IB CHEMISTRY SL/HL
TOPICAL QUESTIONS
1. QUANTITATIVE CHEMISTRY
2. ATOMIC THEORY
3. PERIODICITY
4. BONDING
5. ENERGETICS
6. KINETICS
7. EQUILIBRIUM
8. ACIDS AND BASES
9. REDOX
10. ORGANIC CHEMISTRY
IB Chemistry SL
Topic1 Questions and Answers
1.
What amount of oxygen, O2 , (in moles) contains 1.8×1022 molecules?
A.
0.0030
B.
0.030
C.
0.30
D.
3.0
(Total 1 mark)
2.
Which compound has the empirical formula with the greatest mass?
A.
C2 H6
B.
C4 H10
C.
C5 H10
D.
C6 H6
(Total 1 mark)
3.
__C2 H2 (g) + __O2(g) → __ CO2(g) + __ H2 O(g)
When the equation above is balanced, what is the coefficient for oxygen?
A.
2
B.
3
C.
4
D.
5
(Total 1 mark)
4.
3.0 dm3 of sulfur dioxide is reacted with 2.0 dm3 of oxygen according to the equation below.
2SO2 (g) + O2 (g) → 2SO3 (g)
What volume of sulfur trioxide (in dm3 ) is formed? (Assume the reaction goes to completion
and all gases are measured at the same temperature and pressure.)
A.
5.0
B.
4.0
C.
3.0
D.
2.0
(Total 1 mark)
1
5.
What will happen to the volume of a fixed mass of gas when its pressure and temperature
(in Kelvin) are both doubled?
A.
It will not change.
B.
It will increase.
C.
It will decrease.
D.
The change cannot be predicted.
(Total 1 mark)
6.
What amount (in moles) is present in 2.0 g of sodium hydroxide, NaOH?
A.
0.050
B.
0.10
C.
20
D.
80
(Total 1 mark)
7.
A hydrocarbon contains 90% by mass of carbon. What is its empirical formula?
A.
CH2
B.
C3 H4
C.
C7 H10
D.
C9 H10
(Total 1 mark)
8.
Copper can react with nitric acid as follows.
3Cu +_HNO3 → _Cu(NO3 )2 +_H2O + _NO
What is the coefficient for HNO3 when the equation is balanced?
A.
4
B.
6
C.
8
D.
10
(Total 1 mark)
2
9.
Lithium hydroxide reacts with carbon dioxide as follows.
2LiOH + CO2 → Li2 CO3 + H2 O
What mass (in grams) of lithium hydroxide is needed to react with 11 g of carbon dioxide?
A.
6
B.
12
C.
24
D.
48
(Total 1 mark)
10.
Which change in conditions would increase the volume of a fixed mass of gas?
Pressure /kPa
Temperature /K
A.
Doubled
Doubled
B.
Halved
Halved
C.
Doubled
Halved
D.
Halved
Doubled
(Total 1 mark)
11.
How many hydrogen atoms are contained in one mole of ethanol, C 2 H5 OH?
A.
5
B.
6
C.
1.0×1023
D.
3.6×1024
(Total 1 mark)
12.
The percentage by mass of the elements in a compound is
C = 72%,
H = 12%,
O = 16%.
What is the mole ratio of C:H in the empirical formula of this compound?
A.
1:1
B.
1:2
C.
1:6
D.
6:1
(Total 1 mark)
3
13.
What is the coefficient for O 2 (g) when the equation below is balanced?
__C3 H8 (g) + __O2(g) → __CO2 (g) + __H2 O(g)
A.
2
B.
3
C.
5
D.
7
(Total 1 mark)
14.
What amount of NaCl (in moles) is required to prepare 250 cm3 of a 0.200 mol dm–3 solution?
A.
50.0
B.
1.25
C.
0.800
D.
0.0500
(Total 1 mark)
15.
For which set of conditions does a fixed mass of an ideal gas have the greatest volume?
Temperature
Pressure
A.
low
low
B.
low
high
C.
high
high
D.
high
low
(Total 1 mark)
16.
Which of the following contains the greatest number of molecules?
A.
1 g of CH3 Cl
B.
1 g of CH2 Cl2
C.
1 g of CHCl3
D.
1 g of CCl4
(Total 1 mark)
4
17.
Which of the following compounds has/have the empirical formula CH 2 O?
I.
CH3 COOH
II.
C6 H12 O6
III.
C12 H22 O11
A.
II only
B.
III only
C.
I and II only
D.
II and III only
(Total 1 mark)
18.
Assuming complete reaction, what volume of 0.200 mol dm–3 HCl(aq) is required to neutralize
25.0 cm3 of 0.200 mol dm–3 Ba(OH)2 (aq)?
A.
12.5 cm3
B.
25.0 cm3
C.
50.0 cm3
D.
75.0 cm3
(Total 1 mark)
19.
Under what conditions would one mole of methane gas, CH 4 , occupy the smallest volume?
A.
273 K and 1.01×105 Pa
B.
273 K and 2.02×105 Pa
C.
546 K and 1.01×105 Pa
D.
546 K and 2.02×105 Pa
(Total 1 mark)
20.
The temperature in Kelvin of 2.0 dm3 of an ideal gas is doubled and its pressure is increased by
a factor of four. What is the final volume of the gas?
A.
1.0 dm3
B.
2.0 dm3
C.
3.0 dm3
D.
4.0 dm3
(Total 1 mark)
5
21.
Which is a correct definition of the term empirical formula?
A.
formula showing the numbers of atoms present in a compound
B.
formula showing the numbers of elements present in a compound
C.
formula showing the actual numbers of atoms of each element in a compound
D.
formula showing the simplest ratio of numbers of atoms of each element in a compound
(Total 1 mark)
22.
The reaction of ethanal and oxygen can be represented by the unbalanced equation below.
__ CH3 CHO + __ O2 → __ CO2 + __ H2 O
When the equation is balanced using the smallest possible integers, what is the coefficient
for O2 ?
A.
3
B.
4
C.
5
D.
6
(Total 1 mark)
23.
The equation for the complete combustion of butane is
2C4 H10 + 13O2 → 8CO2 + 10H2 O
What is the amount (in mol) of carbon dioxide formed by the complete combustion of three
moles of butane?
A.
4
B.
8
C.
12
D.
24
(Total 1 mark)
24.
Which solution contains the greatest amount (in mol) of solute?
A.
10.0 cm3 of 0.500 mol dm–3 NaCl
B.
20.0 cm3 of 0.400 mol dm–3 NaCl
C.
30.0 cm3 of 0.300 mol dm–3 NaCl
D.
40.0 cm3 of 0.200 mol dm–3 NaCl
(Total 1 mark)
6
25.
A fixed mass of an ideal gas has a volume of 800 cm3 under certain conditions. The pressure (in
kPa) and temperature (in K) are both doubled. What is the volume of the gas after these changes
with other conditions remaining the same?
A.
200 cm3
B.
800 cm3
C.
1600 cm3
D.
3200 cm3
(Total 1 mark)
26.
The complete oxidation of propane produces carbon dioxide and water as shown below.
C3 H8 + __O2 →__CO2 + __H2O
What is the total of the coefficients for the products in the balanced equation for 1 mole of
propane?
A.
6
B.
7
C.
12
D.
13
(Total 1 mark)
27.
The relative molecular mass (M r) of a compound is 60. Which formulas are possible for this
compound?
I.
CH3 CH2 CH2NH2
II.
CH3 CH2 CH2OH
III.
CH3 CH(OH)CH3
A.
I and II only
B.
I and III only
C.
II and III only
D.
I, II and III
(Total 1 mark)
28.
Which sample has the least number of atoms?
A.
1 mol of H2 SO4
B.
1 mol of CH3 COOH
C.
2 mol of H2 O2
D.
2 mol of NH3
(Total 1 mark)
7
29.
Avogadro’s constant has the same value as the number of
A.
molecules in 1 mol of solid iodine.
B.
atoms in 1 mol of chlorine gas.
C.
ions in 1 mol of solid potassium bromide.
D.
protons in 1 mol of helium gas.
(Total 1 mark)
30.
What is the total of all the coefficients in the balanced equation for the reduction of 1 mol of
MnO4 – ?
__ MnO4 – +__H+ + __ e– →__Mn2+ + __H2 O
A.
5
B.
9
C.
17
D.
19
(Total 1 mark)
31.
Which contains the same number of ions as the value of Avogadro’s constant?
A.
0.5 mol NaCl
B.
0.5 mol MgCl2
C.
1.0 mol Na 2 O
D.
1.0 mol MgO
(Total 1 mark)
32.
A reaction occurring in the extraction of lead from its ore can be represented by this unbalanced
equation:
__ PbS + __O2 → __ PbO + __ SO2
When the equation is balanced using the smallest possible whole numbers, what is the
coefficient for O2 ?
A.
1
B.
2
C.
3
D.
4
(Total 1 mark)
8
33.
The equation for a reaction occurring in the synthesis of methanol is
CO2 + 3H2 → CH3 OH + H2 O
What is the maximum amount of methanol that can be formed from 2 mol of carbon dioxide and
3 mol of hydrogen?
A.
1 mol
B.
2 mol
C.
3 mol
D.
5 mol
(Total 1 mark)
34.
Which solution contains 0.1 mol of sodium hydroxide?
A.
1 cm3 of 0.1 mol dm–3 NaOH
B.
10 cm3 of 0.1 mol dm–3 NaOH
C.
100 cm3 of 1.0 mol dm–3 NaOH
D.
1000 cm3 of 1.0 mol dm–3 NaOH
(Total 1 mark)
35.
A cylinder of gas is at a pressure of 40 kPa. The volume and temperature (in K) are both
doubled. What is the pressure of the gas after these changes?
A.
10 kPa
B.
20 kPa
C.
40 kPa
D.
80 kPa
(Total 1 mark)
36.
Which of the following quantities has units?
A.
Relative atomic mass
B.
Relative molecular mass
C.
Molar mass
D.
Mass number
(Total 1 mark)
9
37.
The empirical formula of a compound is C 2 H4 O. Which molecular formulas are possible for this
compound?
I.
CH3 COOH
II.
CH3 CH2 CH2COOH
III.
CH3 COOCH2CH3
A.
I and II only
B.
I and III only
C.
II and III only
D.
I, II and III
(Total 1 mark)
38.
Calcium carbonate decomposes on heating as shown below.
CaCO3 → CaO + CO2
When 50 g of calcium carbonate are decomposed, 7 g of calcium oxide are formed. What is the
percentage yield of calcium oxide?
A.
7
B.
25
C.
50
D.
75
(Total 1 mark)
39.
Sodium reacts with water as shown below.
__ Na + __ H2 O → __ NaOH + __ H2
What is the total of all the coefficients when the equation is balanced using the smallest possible
whole numbers?
A.
3
B.
4
C.
6
D.
7
(Total 1 mark)
10
40.
What is the total number of ions present in the formula, Al2 (SO4 )3 ?
A.
2
B.
3
C.
5
D.
6
(Total 1 mark)
41.
A 4 g sample of sodium hydroxide, NaOH, is dissolved in water and made up to 500 cm3 of
aqueous solution. What is the concentration of the resulting solution?
A.
0.1 mol dm–3
B.
0.2 mol dm–3
C.
0.5 mol dm–3
D.
1.0 mol dm–3
(Total 1 mark)
42.
Methane, CH4 , burns in oxygen gas to form carbon dioxide and water. How many moles of
carbon dioxide will be formed from 8.0 g of methane?
A.
0.25
B.
0.50
C.
1.0
D.
2.0
(Total 1 mark)
43.
What is the empirical formula of a compound containing 50 by mass of element X (Ar = 20)
and 50 by mass of element Y (Ar = 25)?
A.
XY
B.
X3 Y2
C.
X4 Y5
D.
X5 Y4
(Total 1 mark)
11
44.
Assuming complete reaction, what volume of 0.200 mol dm–3 potassium hydroxide solution
(KOH(aq)), is required to neutralize 25.0 cm3 of 0.200 mol dm–3 aqueous sulfuric acid,
(H2 SO4 (aq))?
A.
12.5 cm3
B.
25.0 cm3
C.
50.0 cm3
D.
75.0 cm3
(Total 1 mark)
45.
Consider the following reaction.
N2 (g) + 3H2 (g)
2NH3 (g)
If the reaction is made to go to completion, what volume of ammonia (in dm3 ) can be prepared
from 25 dm3 of nitrogen and 60 dm3 of hydrogen? All volumes are measured at the same
temperature and pressure.
A.
40
B.
50
C.
85
D.
120
(Total 1 mark)
46.
The temperature in Kelvin of 1.0 dm3 of an ideal gas is doubled and its pressure is tripled. What
is the final volume of the gas in dm3 ?
A.
1
3
B.
2
3
C.
3
2
D.
1
6
(Total 1 mark)
12
47.
On complete combustion, a sample of a hydrocarbon compound produces 1.5 mol of carbon
dioxide and 2.0 mol of water. What is the molecular formula of this hydrocarbon?
A.
C2 H2
B.
C2 H4
C.
C3 H4
D.
C3 H8
(Total 1 mark)
48.
When excess BaCl2 (aq) was added to a sample of Fe(NH 4 )2 (SO4)2 (aq) to determine the amount
in moles of sulfate present, 5.02×10–3 mol of BaSO4 was obtained. How many moles of sulfate
ions and iron ions were in the sample of Fe(NH 4 )2 (SO4 )2?
Amount of sulfate ions / moles
–3
Amount of iron ions / moles
A.
5.02×10
2.51×10–3
B.
10.04×10–3
5.02×10–3
C.
2.51×10–3
5.02×10–3
D.
10.04×10–3
2.51×10–3
(Total 1 mark)
49.
What volume of 0.500 mol dm–3 sulfuric acid solution is required to react completely with 10.0
g of calcium carbonate according to the equation below?
CaCO3 (s) + H2 SO4 (aq) → CaSO4(aq) + H2 O(l) + CO2 (g)
A.
100 cm3
B.
200 cm3
C.
300 cm3
D.
400 cm3
(Total 1 mark)
50.
Which expression gives the amount (in mol) of a substance, if the mass is given in grams?
A.
mass
molar mass
B.
molar mass
mass
C.
1
molar mass
D.
mass × molar mass
(Total 1 mark)
13
51.
What is the total number of atoms in 0.20 mol of propanone, CH 3 COCH3 ?
A.
1.2×1022
B.
6.0×1023
C.
1.2×1024
D.
6.0×1024
(Total 1 mark)
52.
When the equation below is balanced for 1 mol of C 3 H4 , what is the coefficient for O 2 ?
C3 C4 + O2 → CO2 + H2 O
A.
2
B.
3
C.
4
D.
5
(Total 1 mark)
53.
Ethyne, C2 H2 , reacts with oxygen according to the equation below. What volume of oxygen (in
dm3 ) reacts with 0.40 dm3 of C2 H2 ?
2C2 H2 (g) + 5O2 (g) → 4CO2 (g) + 2H2 O(g)
A.
0.40
B.
0.80
C.
1.0
D.
2.0
(Total 1 mark)
54.
Ethyne, C2 H2 , reacts with oxygen according to the equation below. What volume of oxygen (in
dm3 ) reacts with 0.40 dm3 of C2 H2 ?
2C2 H2 (g) + 5O2 (g) → 4CO2 (g) + 2H2 O(g)
A.
0.40
B.
0.80
C.
1.0
D.
2.0
(Total 1 mark)
14
55.
What is the coefficient for H + when the redox equation below is balanced?
Ag(s) + NO3 –(aq) +H+(aq) → Ag+(aq) + NO(g) + H2 O(l)
A.
1
B.
2
C.
3
D.
4
(Total 1 mark)
56.
How many hydrogen atoms are in one mole of ethanol, C 2 H5 OH?
A.
1.00×1023
B.
3.61×1024
C.
5.00
D.
6.00
(Total 1 mark)
57.
What is the coefficient for H 2 SO4 (aq) when the following equation is balanced, using the
smallest possible integers?
__Mg3 N2 (s) + __H2SO4(aq) → __MgSO4 (aq) + __(NH4)2SO4 (aq)
A.
1
B.
3
C.
4
D.
7
(Total 1 mark)
58.
Air bags in cars inflate when sodium azide decomposes to form sodium and nitrogen:
2NaN3 (s) → 2Na(s) + 3N2 (g)
Calculate the amount, in moles, of nitrogen gas produced by the decomposition of 2.52 mol of
NaN3 (s).
A.
1.68
B.
2.52
C.
3.78
D.
7.56
(Total 1 mark)
15
–3
3
59.
What volume, in cm , of 0.200 mol dm
A.
12.5
B.
25.0
C.
50.0
D.
75.0
–3
3
HCl(aq) is required to neutralize 25.0 cm of 0.200 mol dm
Ba(OH)2 (aq)?
(Total 1 mark)
60.
The relative molecular mass of aluminium chloride is 267 and its composition by mass is 20.3%
Al and 79.7% chlorine. Determine the empirical and molecular formulas of aluminium chloride.
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(Total 4 marks)
61.
Sodium reacts with water as follows.
2Na(s) + 2H2 O(l) → 2NaOH(aq) + H2 (g)
1.15 g of sodium is allowed to react completely with water. The resulting solution is diluted to
250 cm3 . Calculate the concentration, in mol dm–3 , of the resulting sodium hydroxide solution.
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(Total 3 marks)
16
62.
(i)
Calcium carbonate is added to separate solutions of hydrochloric acid and ethanoic acid
of the same concentration. State one similarity and one difference in the observations you
could make.
…………………………………………………………………………………………..
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(2)
(ii)
Write an equation for the reaction between hydrochloric acid and calcium carbonate.
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(2)
(iii)
Determine the volume of 1.50 mol dm–3 hydrochloric acid that would react with exactly
1.25 g of calcium carbonate.
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(3)
(iv)
Calculate the volume of carbon dioxide, measured at 273 K and
1.01×105 Pa, which would be produced when 1.25 g of calcium carbonate reacts
completely with the hydrochloric acid.
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(2)
(Total 9 marks)
17
63.
An organic compound, A, containing only the elements carbon, hydrogen and oxygen was
analysed.
(a)
A was found to contain 54.5 C and 9.1 H by mass, the remainder being oxygen.
Determine the empirical formula of the compound.
(3)
(b)
A 0.230 g sample of A, when vaporized, had a volume of 0.0785 dm3 at 95C and 102
kPa. Determine the relative molecular mass of A.
(3)
(c)
Determine the molecular formula of A using your answers from parts (a) and (b).
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(1)
(Total 7 marks)
64.
An organic compound A contains 62.0 by mass of carbon, 24.1 by mass of nitrogen, the
remainder being hydrogen.
(i)
Determine the percentage by mass of hydrogen and the empirical formula of A.
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(3)
(ii)
Define the term relative molecular mass.
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(2)
18
(iii)
The relative molecular mass of A is 116. Determine the molecular formula of A.
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(1)
(Total 6 marks)
65.
An organic compound A contains 62.0 by mass of carbon, 24.1 by mass of nitrogen, the
remainder being hydrogen.
(i)
Determine the percentage by mass of hydrogen and the empirical formula of A.
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(3)
(ii)
Define the term relative molecular mass.
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(2)
(iii)
The relative molecular mass of A is 116. Determine the molecular formula of A.
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(1)
(Total 6 marks)
66.
Propane and oxygen react according to the following equation.
C3 H8 (g) + 5O2 (g) → 3CO2 (g) + 4H2 O(g)
Calculate the volume of carbon dioxide and water vapour produced and the volume of oxygen
remaining, when 20.0 dm3 of propane reacts with 120.0 dm3 of oxygen. All gas volumes are
measured at the same temperature and pressure.
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(Total 3 marks)
19
67.
State and explain what would happen to the pressure of a given mass of gas when its absolute
temperature and volume are both doubled.
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(Total 3 marks)
68.
(i)
Crocetin consists of the elements carbon, hydrogen and oxygen. Determine the empirical
formula of crocetin, if 1.00 g of crocetin forms 2.68 g of carbon dioxide and 0.657 g of
water when it undergoes complete combustion.
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(6)
(ii)
Determine the molecular formula of crocetin given that 0.300 mole of crocetin has a mass
of 98.5 g
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(2)
(Total 8 marks)
20
69.
A solution containing ammonia requires 25.0 cm3 of 0.100 mol dm–3 hydrochloric acid to reach
the equivalence point of a titration.
(i)
Write an equation for the reaction of ammonia with hydrochloric acid
(1)
(ii)
Calculate the amount (in mol) of hydrochloric acid and ammonia that react.
(2)
(iii)
Calculate the mass of ammonia in the solution.
(2)
(Total 5 marks)
70.
A toxic gas, A, consists of 53.8% nitrogen and 46.2% carbon by mass. At 273 K and
1.01×105 Pa, 1.048 g of A occupies 462 cm3 . Determine the empirical formula of A.
Calculate the molar mass of the compound and determine its molecular structure.
(Total 3 marks)
71.
An oxide of copper was reduced in a stream of hydrogen as shown below.
excess hydrogen
burning
hydrogen
gas
oxide of copper
in a dish
HEAT
After heating, the stream of hydrogen gas was maintained until the apparatus had cooled.
The following results were obtained.
Mass of empty dish = 13.80 g
Mass of dish and contents before heating = 21.75 g
Mass of dish and contents after heating and leaving to cool = 20.15 g
(a)
Explain why the stream of hydrogen gas was maintained until the apparatus cooled.
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(1)
(b)
Calculate the empirical formula of the oxide of copper using the data above, assuming
complete reduction of the oxide.
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(3)
21
(c)
Write an equation for the reaction that occurred.
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(1)
(d)
State two changes that would be observed inside the tube as it was heated.
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(2)
(Total 7 marks)
72.
Copper metal may be produced by the reaction of copper(I) oxide and copper(I) sulfide
according to the below equation.
2Cu2 O + Cu2 S → 6Cu + SO2
A mixture of 10.0 kg of copper(I) oxide and 5.00 kg of copper(I) sulfide was heated until no
further reaction occurred.
(a)
Determine the limiting reagent in this reaction, showing your working.
....................................................................................................................................
....................................................................................................................................
....................................................................................................................................
....................................................................................................................................
....................................................................................................................................
....................................................................................................................................
(3)
(b)
Calculate the maximum mass of copper that could be obtained from these masses of
reactants.
....................................................................................................................................
....................................................................................................................................
....................................................................................................................................
....................................................................................................................................
(2)
(Total 5 marks)
22
73.
The reaction below represents the reduction of iron ore to produce iron.
2Fe2 O3 + 3C → 4Fe + 3CO2
A mixture of 30 kg of Fe 2 O3 and 5.0 kg of C was heated until no further reaction occurred.
Calculate the maximum mass of iron that can be obtained from these masses of reactants.
..............................................................................................................................................
..............................................................................................................................................
..............................................................................................................................................
..............................................................................................................................................
..............................................................................................................................................
..............................................................................................................................................
..............................................................................................................................................
..............................................................................................................................................
..............................................................................................................................................
..............................................................................................................................................
(Total 5 marks)
74.
0.502 g of an alkali metal sulfate is dissolved in water and excess barium chloride solution,
BaCl2 (aq) is added to precipitate all the sulfate ions as barium sulfate, BaSO 4 (s). The precipitate
is filtered and dried and weighs 0.672 g.
(a)
Calculate the amount (in mol) of barium sulfate formed.
....................................................................................................................................
....................................................................................................................................
....................................................................................................................................
....................................................................................................................................
(2)
(b)
Determine the amount (in mol) of the alkali metal sulfate present.
....................................................................................................................................
(1)
(c)
Determine the molar mass of the alkali metal sulfate and state its units.
....................................................................................................................................
....................................................................................................................................
....................................................................................................................................
(2)
23
(d)
Deduce the identity of the alkali metal, showing your workings.
....................................................................................................................................
....................................................................................................................................
....................................................................................................................................
....................................................................................................................................
(2)
(e)
Write an equation for the precipitation reaction, including state symbols.
....................................................................................................................................
....................................................................................................................................
(2)
(Total 9 marks)
75.
0.600 mol of aluminium hydroxide is mixed with 0.600 mol of sulfuric acid, and the following
reaction occurs:
2Al(OH)3 (s) + 3H2 SO4 (aq) → Al2 (SO4 )3 (aq) + 6H2 O(l)
(a)
Determine the limiting reactant.
.....................................................................................................................................
.....................................................................................................................................
.....................................................................................................................................
.....................................................................................................................................
(2)
(b)
Calculate the mass of Al2 (SO4 )3 produced.
.....................................................................................................................................
.....................................................................................................................................
.....................................................................................................................................
.....................................................................................................................................
(2)
(c)
Determine the amount (in mol) of excess reactant that remains.
.....................................................................................................................................
.....................................................................................................................................
(1)
24
(d)
Define a Brønsted-Lowry acid and a Lewis base.
Brønsted-Lowry acid
.....................................................................................................................................
Lewis base
.....................................................................................................................................
(2)
(e)
H2 SO4 (aq) is a strong acid. State the name and the formula of any weak acid.
.....................................................................................................................................
.....................................................................................................................................
(1)
(Total 8 marks)
25
1.
B
2.
B
3.
D
4.
C
5.
A
6.
A
7.
B
8.
C
9.
B
10.
11.
D
D
12.
B
13.
C
14.
D
15.
D
16.
A
17.
C
18.
C
19.
B
20.
21.
A
D
22.
C
23.
C
24.
C
25.
B
26.
B
27.
C
28.
A
29.
A
30.
D
31.
A
26
32.
C
33.
A
34.
C
35.
C
36.
C
37.
C
38.
B
39.
D
40.
C
41.
B
42.
B
43.
D
44.
C
45.
A
46.
B
47.
D
48.
A
49.
B
50.
A
51.
C
52.
C
53.
C
54.
C
55.
D
56.
B
57.
C
58.
C
59.
C
60.
Al 20.3 Cl 79.70 or similar working (no penalty for use of 27 or 35.5);
26 .98 35 .45
27
empirical formula AlCl3 ;
molecular formula: n = 267 = 2;
133 .5
Al2 Cl6 ;
Full credit can be obtained if the calculations are carried out
by another valid method. Two correct formulas but no valid
method scores [2 max].
[4]
61.
moles of Na = 1.15 = 0.05;
23
moles of NaOH = 0.05;
Accept “same as moles of Na”
concentration =  0.05  = 0.20 (mol dm–3 )
 0.25 
Allow ECF from moles of NaOH
3
[3]
62.
(i)
bubbling/effervescence/dissolving of CaCO 3 /gas given off
(do not accept CO2 produced);
more vigorous reaction with HCl/OWTTE;
(ii)
2HCl(aq) + CaCO3 (s) → CaCl2 (aq) + CO2 (g) + H2 O(1);
[1] for correct formulas, [1] for balanced, state symbols not
(iii)
amount of CaCO3 = 1.25 (no penalty for use of 100);
100.09
amount of HCl = 2×0.0125 = 0.0250 mol (allow ECF);
volume of HCl = 0.0167 dm3 /16.7 cm3 (allow ECF);
(iv)
2
2
3
1:1 ratio of CaCO3 to CO2 /use 0.0125 moles CO2 (allow ECF);
(0.0125×22.4) = 0.28 dm3 /280 cm3 /2.8×10–4 m3 (allow ECF);
Accept calculation using pV=nRT.
1
[9]
63.
(a)
% of oxygen = 36.4;
C=
54.5
9.1
36.4
,H=
,O=
;
12.01
1.01
16.00
Do not penalize if 12, 1 and 16 are used.
C2 H4 O;
3
If atomic numbers or incorrect Ar values used, only first mark
can be scored.
Award [3] for correct formula without working.
(b)
pV = nRT/pV =
mRT
/correct rearrangement;
Mr
28
Mr =
0.230  8.31  368
;
102 10 3  0.0785 10 −3
Award [1] for 368 even if incorrect expression given.
Mr = 87.8;
3
Accept answer in range 87.8 to 88.
Do not allow ECF.
Award [3] for correct final answer
(c)
C4 H8 O2 ;
1
Answer does not need to show working to receive the mark.
Do not allow ECF.
[7]
64.
(i)
62.0
/ 5.16
12.01
C
N
H
24.1
13.9
/ 1.72
/ 13.8
14.01
1.01
Award [2] for above.
No penalty for use of whole number atomic masses.
If atomic numbers used then only mark for % of H can be
awarded.
If H % and calculation missing, award [1], and last mark
cannot be scored.
If H % calculation incorrect apply ECF.
C3 NH8 ;
3
Correct empirical formula scores [3].
(ii)
the average mass of a molecule;
compared to 1/12 of (the mass of) one atom of 12 C/compared to
C-12 taken as 12;
OR
average mass of a molecule
mass of 1 /12 of one atom of 12 C
Award [2] for the equation above.
(iii)
C6 N2 H16;
2
1
[6]
65.
(i)
C
62.0
/ 5.16
12.01
N
H
24.1
13.9
/ 1.72
/ 13.8
14.01
1.01
Award [2] for above.
No penalty for use of whole number atomic masses.
If atomic numbers used then only mark for % of H can be
awarded.
If H % and calculation missing, award [1], and last mark
29
cannot be scored.
If H % calculation incorrect apply ECF.
C3 NH8 ;
Correct empirical formula scores [3].
(ii)
3
the average mass of a molecule;
compared to 1/12 of (the mass of) one atom of 12 C/compared to
C-12 taken as 12;
OR
average mass of a molecule
mass of 1/12 of one atom of 12 C
Award [2] for the equation above.
(iii)
C6 N2 H16;
2
1
[6]
66.
60.0 dm3 CO2 ;
80.0 dm3 H2 O;
20.0 dm3 O2 ;
3
Apply −1(U).
[3]
67.
overall there will be no change to the pressure;
double absolute temperature and the pressure doubles;
double volume and the pressure halves;
Apply ECF if points 2 and 3 are incorrect.
OR
Use PV = nRT, Since n and R are constant;
V and T are both doubled;
P will remain unchanged;
OR
OWTTE for mathematical interpretation
e.g. T  P, therefore 2P;
V  1/P, therefore ½P;
No change to P, ½P×2P = P;
68.
(i)
3
n(C)(= n(CO2 ) = 2.68 g÷44.01 g mol−1 ) = 0.0609 mol;
n(H)(= 2×n(H2 O) = 0.657 g÷18.02 g mol−1 ) = 0.0729 mol;
m(C) = 0.0609 mol×12.01 g mol−1 = 0.731 g
and m(H) = 0.0729 mol×1.01 g mol−1 = 0.0736 g;
m(O) = (1.00 − 0.731 − 0.0736)g = 0.195g;
n(C)
0.0609
n(H)
0.0730
0.0609
0.0609
0.0730
0.0730
n(O)
0.195
16.00
0.0122
0.0122
30
0.0122
0.0122
0.0122
4.99
5.98
1.00;
empirical formula: C5 H6 O;
For C5 H6 award [4 max].
6
Steps used to arrive at the correct amounts (in moles) are
required for full marks.
(ii)
M(crocetin) = 98.5 g÷0.300 mol = 328 (g mol−1 );
(
328
= 4)
82.11
molecular formula: C20 H24O4;
ECF from (i).
2
[8]
69.
(i)
(ii)
(iii)
NH3 (aq) + HCl(aq) → NH4 Cl(aq) ;
States not required for mark
1
n(HCl) = cV = 0.100 mol dm−3 ×0.0250 dm3 = 0.00250 mol;
n(NH3 ) = n(HCl) = 0.00250 mol;
ECF
2
(M (NH3 ) = 14.01 + 3(1.01) =) 17.04/17.0 (g mol−1 );
m(NH3 ) = 0.00250 mol×17.04g mol −1 = 0.0426g/0.0425g;
ECF
2
[5]
70.
empirical formula = CN;
Working must be shown to get point.
M r = 51.9 (g mol–1 );
:N
C⎯C
N:;
3
[3]
71.
(a)
to prevent (re)oxidation of the copper/OWTTE;
(b)
number of moles of oxygen = 1.60 = 0.10;
16 .00
6.35
number of moles of copper =
= 0.10;
63 .55
empirical formula = Cu (0.10) : O (0.10) = CuO;
Allow ECF.
Award [1] for CuO with no working.
Alternate solution
6.35
= 79.8%
7.95
1.60
= 20.2%
7.95
70 .8
= 1.25
63 .5
20 .2
= 1.29
16
1
3
31
(c)
H2 + CuO → Cu + H2 O;
Allow ECF.
(d)
(black copper oxide) solid turns red/brown;
condensation/water vapour (on sides of test tube);
Accept change colour.
Do not accept reduction of sample size.
1
2
[7]
72.
(a)
n(Cu2 O) = 10.0×103 ÷143.1 = 69.9 mol;
n(Cu2 S) = 5.00×103 ÷159.16 = 31.4 mol;
Penalise failure to convert kg → g once only.
Cu2 S is the limiting reagent;
ECF from above answers.
(b)
n(Cu) = 6×n(Cu2 S) = 6×31.4 = 188 mol;
m(Cu) = 188×63.55 = 11900 − 12000 g/11.9 − 12.0 kg;
If Cu 2 O given in (a), allow 3×n(Cu2 O) and 3×n(Cu2O)×63.55.
3
2
Allow ECF from (a).
[5]
73.
n(Fe2 O3 ) = 30×103 ÷159.7/n(Fe2 O3 ) = 188 mol;
n(C) = 5.0×103 ÷12.01/n(C) = 416 mol;
Fe2 O3 is the limiting reagent or implicit in calculation;
n(Fe) = 2×n(Fe 2 O3 ) = 2×188 = 376 mol;
m(Fe) = 376×55.85 = 21 kg;
Accept 2 sig. fig. or 3 sig. fig., otherwise use − 1(SF).
Correct final answers score [5].
Allow ECF.
[5]
74.
(a)
M(BaSO4 ) (= 137.34 + 32.06 + 4(16.00)) = 233.40 (g mol−1 );
Accept 233.4 but not 233


0.672 g
 = 0.00288 / 2.88×10−3 (mol);
n(BaSO4 )  =
−1 
233
.
40
g
mol


ECF from M value
(b)
(c)
n (alkali metal sulfate) = 0.00288 / 2.88×10−3 (mol);
ECF
m
0.502 g
M =  =
 n 0.00288 mol
ECF
units: g mol−1 ;
2
1

=  174.31 / 174.3 / 174;

2
32
(d)
 (174 − (32 + (4 16 ) 
(2(Ar) + 32 + 4(16) = 174, thus) Ar = 39 / Ar = 
 =39;
2


Accept answer between 38.9 and 39.2
ECF
potassium/K;
ECF from Ar value
2
(e)
K2 SO4 (aq) + BaCl2 (aq) → BaSO4 (s) + 2KCl(aq)
Award [1] for balanced equation and [1] for state symbols
ECF if another alkali metal arrived at in (d)
Accept net ionic equation
If no answer arrived at in (d), but correct equation given
involving any alkali metal, then award [1 max]
2
[9]
75.
(a)
(b)
(c)
0.600 mol Al(OH)3 ≡ (1.5)(0.600) mol H2 SO4 /0.900 mol H2 SO4
needed, but only 0.600 mol used;
H2 SO4 limiting reactant;
Some working must be shown in order to score the second
point.
2
0.200 mol Al2 (SO4 )3 ;
68.4(g);
Penalize incorrect units.
2
0.200 mol;
1
Use ECF from (a).
(d)
(e)
A Brønsted-Lowry acid is a proton/H+ donor;
A Lewis base is an electron-pair donor;
H2 CO3 and carbonic acid/CH3 COOH and ethanoic acid;
Accept any other weak acid and correct formula.
2
1
[8]
33
IB Chemistry – SL
Topic 2 Questions
1.
Consider the composition of the species W, X, Y and Z below. Which species is an anion?
Species
Number of protons
Number of neutrons
Number of electrons
W
9
10
10
X
11
12
11
Y
12
12
12
Z
13
14
10
A.
W
B.
X
C.
Y
D.
Z
(Total 1 mark)
2.
Energy levels for an electron in a hydrogen atom are
A.
evenly spaced.
B.
farther apart near the nucleus.
C.
closer together near the nucleus.
D.
arranged randomly.
(Total 1 mark)
3.
Which is related to the number of electrons in the outer main energy level of the elements from
the alkali metals to the halogens?
I.
Group number
II.
Period number
A.
I only
B.
II only
C.
Both I and II
D.
Neither I nor II
(Total 1 mark)
Page 1 of 19
4.
How do bond length and bond strength change as the number of bonds between two atoms
increases?
Bond length
Bond strength
A.
Increases
increases
B.
Increases
decreases
C.
Decreases
increases
D.
Decreases
decreases
(Total 1 mark)
5.
Which of the following is true for CO 2 ?
C
CO2 molecule
O bond
A.
Polar
non-polar
B.
non-polar
polar
C.
Polar
polar
D.
non-polar
non-polar
(Total 1 mark)
6.
The molar masses of C2 H6 , CH3 OH and CH3 F are very similar. How do their boiling points
compare?
A.
C2 H6 < CH3 OH < CH3 F
B.
CH3 F < CH3 OH < C2 H6
C.
CH3 OH < CH3 F < C2 H6
D.
C2 H6 < CH3 F < CH3 OH
(Total 1 mark)
7.
What is the correct number of each particle in a fluoride ion,
protons
neutrons
electrons
A.
9
10
8
B.
9
10
9
C.
9
10
10
D.
9
19
10
19 –
F?
(Total 1 mark)
8.
Which statement is correct for the emission spectrum of the hydrogen atom?
A.
The lines converge at lower energies.
B.
The lines are produced when electrons move from lower to higher energy levels.
C.
The lines in the visible region involve electron transitions into the energy level closest to
the nucleus.
D.
The line corresponding to the greatest emission of energy is in the ultraviolet region.
Page 2 of 19
9.
Which is the correct description of polarity in F2 and HF molecules?
A.
Both molecules contain a polar bond.
B.
Neither molecule contains a polar bond.
C.
Both molecules are polar.
D.
Only one of the molecules is polar.
(Total 1 mark)
10.
Which types of bonding are present in CH 3 CHO in the liquid state?
I.
Single covalent bonding
II.
Double covalent bonding
III.
Hydrogen bonding
A.
I and II only
B.
I and III only
C.
II and III only
D.
I, II and III
(Total 1 mark)
11.
Which statement(s) is/are generally true about the melting points of substances?
I.
Melting points are higher for compounds containing ions than for compounds
containing molecules.
II.
A compound with a low melting point is less volatile than a compound with a high
melting point.
III.
The melting point of a compound is decreased by the presence of impurities.
A.
I only
B.
I and III only
C.
II and III only
D.
I, II and III
(Total 1 mark)
12.
How many valence electrons are present in an atom of an element with atomic number 16?
A.
2
B.
4
C.
6
D.
8
(Total 1 mark)
Page 3 of 19
13.
A certain sample of element Z contains 60% of 69 Z and 40% of 71 Z. What is the relative atomic
mass of element Z in this sample?
A.
69.2
B.
69.8
C.
70.0
D.
70.2
(Total 1 mark)
14.
What is the difference between two neutral atoms represented by the symbols
A.
The number of neutrons only.
B.
The number of protons and electrons only.
C.
The number of protons and neutrons only.
D.
The number of protons, neutrons and electrons.
59
27
Co and
59
28
Ni?
(Total 1 mark)
15.
How many electrons are there in one
A.
10
B.
12
C.
14
D.
22
24
+2
12 Mg
ion?
(Total 1 mark)
16.
The electron arrangement of sodium is 2.8.1. How many occupied main electron energy levels
are there in an atom of sodium?
A.
1
B.
3
C.
10
D.
11
(Total 1 mark)
Page 4 of 19
17.
Information is given about four different atoms:
Atom
neutrons
protons
W
22
18
X
18
20
Y
22
16
Z
20
18
Which two atoms are isotopes?
A.
W and Y
B.
W and Z
C.
X and Z
D.
X and Y
(Total 1 mark)
18.
Which statement is correct about a line emission spectrum?
A.
Electrons absorb energy as they move from low to high energy levels.
B.
Electrons absorb energy as they move from high to low energy levels.
C.
Electrons release energy as they move from low to high energy levels.
D.
Electrons release energy as they move from high to low energy levels.
(Total 1 mark)
19.
How many neutrons are there in the ion
A.
8
B.
10
C.
16
D.
20
18
O2– ?
(Total 1 mark)
20.
What is the electron arrangement of silicon?
A.
2.4
B.
2.8
C.
2.8.4
D.
2.8.8
(Total 1 mark)
Page 5 of 19
21.
Which statement is correct about the isotopes of an element?
A.
They have the same mass number
B.
They have the same electron arrangement
C.
They have more protons than neutrons
D.
They have the same numbers of protons and neutrons
(Total 1 mark)
22.
What is the difference between two neutral atoms represented by the symbols
A.
The number of neutrons only.
B.
The number of protons and electrons only.
C.
The number of protons and neutrons only.
D.
The number of protons, neutrons and electrons.
210
84
210
85
Po and
At?
(Total 1 mark)
23.
Which statements are correct for the emission spectrum of the hydrogen atom?
I.
II.
III.
The lines converge at lower energies.
Electron transition to n =1 are responsible for lines in the UV region.
Lines are produced when electrons move from higher to lower energy levels.
A.
I and II only
B.
I and III only
C.
II and III only
D.
I, II and III
(Total 1 mark)
24.
What is the symbol for a species that contains 15 protons, 16 neutrons and 18 electrons?
A.
31
16
S
B.
31
16
S 3−
C.
31
15
P−
D.
31
15
P 3−
(Total 1 mark)
Page 6 of 19
25.
What is the electron arrangement of an Al3+ ion?
A.
2, 8
B.
2, 3
C.
2, 8, 3
D.
2, 8, 8
(Total 1 mark)
26.
What will happen to the volume of a fixed mass of gas if the pressure and the Kelvin
temperature are both doubled?
A.
It will remain the same.
B.
It will be double its initial volume.
C.
It will be one-half its initial volume.
D.
It will be four times its initial volume.
(Total 1 mark)
27.
Which species has 54 electrons and 52 protons?
A.
128
2−
52 Te
B.
132
2+
54 Xe
C.
132
2−
54 Xe
D.
128
2+
52 Te
(Total 1 mark)
28.
What is the correct sequence for the processes occurring in a mass spectrometer?
A.
vaporization, ionization, acceleration, deflection
B.
vaporization, acceleration, ionization, deflection
C.
ionization, vaporization, acceleration, deflection
D.
ionization, vaporization, deflection, acceleration
(Total 1 mark)
29.
The percentage composition by mass of a hydrocarbon is C = 85.6% and H = 14.4%.
(a)
Calculate the empirical formula of the hydrocarbon.
.....................................................................................................................................
.....................................................................................................................................
.....................................................................................................................................
.....................................................................................................................................
(2)
Page 7 of 19
(b)
A 1.00 g sample of the hydrocarbon at a temperature of 273 K and a pressure of
1.01×105 Pa (1.00 atm) has a volume of 0.399 dm3 .
(i)
Calculate the molar mass of the hydrocarbon.
..........................................................................................................................
..........................................................................................................................
..........................................................................................................................
..........................................................................................................................
(2)
(ii)
Deduce the molecular formula of the hydrocarbon.
(1)
(Total 5 marks)
30.
State the number of protons, electrons and neutrons in the ion
15
7
N3– .
……………………………………………………………………………………………
……………………………………………………………………………………………
……………………………………………………………………………………………
(Total 2 marks)
31.
A sample of germanium is analysed in a mass spectrometer. The first and last processes in mass
spectrometry are vaporization and detection.
(a)
(i)
State the names of the other three processes in the order in which they occur in a
mass spectrometer.
.........................................................................................................................
.........................................................................................................................
(2)
(ii)
For each of the processes named in (a) (i), outline how the process occurs.
.........................................................................................................................
.........................................................................................................................
.........................................................................................................................
.........................................................................................................................
.........................................................................................................................
.........................................................................................................................
(3)
Page 8 of 19
(b)
The sample of germanium is found to have the following composition:
Isotope
Relative abundance / 
(i)
70
Ge
22.60
72
Ge
25.45
74
Ge
36.73
76
Ge
15.22
Define the term relative atomic mass.
.........................................................................................................................
.........................................................................................................................
.........................................................................................................................
(2)
(ii)
Calculate the relative atomic mass of this sample of germanium, giving your
answer to two decimal places.
.........................................................................................................................
.........................................................................................................................
.........................................................................................................................
.........................................................................................................................
(2)
(Total 9 marks)
32.
Define the following terms.
(i)
atomic number
....................................................................................................................................
....................................................................................................................................
(1)
(ii)
mass number
....................................................................................................................................
....................................................................................................................................
(1)
(Total 2 marks)
33.
State the electron arrangements of the following species:
Si ........................................................................................................................................
P 3– ........................................................................................................................................
(Total 2 marks)
34.
Identify the numbers of protons, neutrons and electrons in the species 33 S2– .
..............................................................................................................................................
..............................................................................................................................................
(Total 1 mark)
Page 9 of 19
35.
State the electron arrangement for atoms of aluminium, nitrogen and fluorine.
(Total 2 marks)
36.
The relative atomic mass of chlorine is 35.45. Calculate the percentage abundance of the two
isotopes of chlorine, 35 Cl and 37 Cl in a sample of chlorine gas.
(Total 2 marks)
37.
(a)
Describe the following stages in the operation of the mass spectrometer.
(i)
ionization
(2)
(ii)
deflection
(2)
(iii)
acceleration
(1)
(b)
(i)
State the meaning of the term isotopes of an element.
(1)
(ii)
Calculate the percentage abundance of the two isotopes of rubidium
85
Rb and 87 Rb.
(2)
(iii)
State two physical properties that would differ for each of the rubidium isotopes.
(1)
(iv)
Determine the full electron configuration of an atom of Si, an Fe 3+ ion and a P 3–
ion.
(3)
(Total 12 marks)
38.
Naturally occurring copper has a relative atomic mass, (Ar), of 63.55 and consists of two
isotopes 63 Cu and 65 Cu.
(i)
Define the term relative atomic mass, Ar.
....................................................................................................................................
....................................................................................................................................
....................................................................................................................................
(1)
Page 10 of 19
(ii)
State and explain which is the more abundant isotope.
....................................................................................................................................
....................................................................................................................................
....................................................................................................................................
(1)
(Total 2 marks)
39.
The element vanadium has two isotopes,
(a)
50
23
V and
51
23V,
and a relative atomic mass of 50.94.
Define the term isotope.
……………………………………………………………………………………….
……………………………………………………………………………………….
(1)
(b)
State the number of protons, electrons and neutrons in
50
23V.
……………………………………………………………………………………….
……………………………………………………………………………………….
(2)
(c)
State and explain which is the more abundant isotope.
……………………………………………………………………………………….
……………………………………………………………………………………….
(1)
(d)
State the name and the mass number of the isotope relative to which all atomic masses
are measured.
……………………………………………………………………………………….
(1)
(Total 5 marks)
40.
(a)
State a physical property that is different for isotopes of an element.
.....................................................................................................................................
(1)
(b)
Chlorine exists as two isotopes, 35 Cl and 37 Cl. The relative atomic mass of chlorine is
35.45. Calculate the percentage abundance of each isotope.
.....................................................................................................................................
.....................................................................................................................................
.....................................................................................................................................
.....................................................................................................................................
(2)
(Total 3 marks)
Page 11 of 19
41.
(a)
Define the term isotope.
.....................................................................................................................................
.....................................................................................................................................
.....................................................................................................................................
(2)
(b)
A sample of argon exists as a mixture of three isotopes.
mass number 36, relative abundance 0.337%
mass number 38, relative abundance 0.0630%
mass number 40, relative abundance 99.6%
Calculate the relative atomic mass of argon.
................................................................................................................................... ..
.....................................................................................................................................
.....................................................................................................................................
.....................................................................................................................................
(2)
(c)
State the number of electrons, protons and neutrons in the ion
56
Fe3+.
electrons: ............................. protons: ............................. neutrons: ...........................
(2)
(Total 6 marks)
42.
The element bromine exists as the isotopes
79.90.
(a)
79
Br and 81 Br, and has a relative atomic mass of
Complete the following table to show the numbers of sub-atomic particles in the species
shown.
an atom of 79 Br
an ion of 81 Br–
protons
neutrons
electrons
(3)
(b)
State and explain which of the two isotopes
bromine.
79
Br and 81 Br is more common in the element
...................................................................................................................................
...................................................................................................................................
(1)
Page 12 of 19
(c)
The element calcium is in the same period of the Periodic Table as bromine.
(i)
Write the electron arrangement for an atom of calcium.
.........................................................................................................................
(1)
(ii)
Deduce the formula of the compound calcium bromide.
.........................................................................................................................
(1)
(Total 6 marks)
43.
Some vaporized magnesium is introduced into a mass spectrometer. One of the ions that reaches
the detector is 25 Mg+.
(a)
Identify the number of protons, neutrons and electrons in the
25
Mg+ ion.
....................................................................................................................................
....................................................................................................................................
(1)
(b)
State how this ion is accelerated in the mass spectrometer.
....................................................................................................................................
....................................................................................................................................
(1)
(c)
The 25 Mg2+ ion is also detected in this mass spectrometer by changing the magnetic field.
Deduce and explain, by reference to the m/z values of these two ions of magnesium,
which of the ions 25 Mg2+ and 25 Mg+ is detected using a stronger magnetic field.
....................................................................................................................................
....................................................................................................................................
....................................................................................................................................
....................................................................................................................................
(2)
(Total 4 marks)
44.
(a)
List the following types of electromagnetic radiation in order of increasing wavelength
(shortest first).
I.
Yellow light
II.
Red light
III.
Infrared radiation
IV.
Ultraviolet radiation
.....................................................................................................................................
(1)
Page 13 of 19
(b)
Distinguish between a continuous spectrum and a line spectrum.
.....................................................................................................................................
.....................................................................................................................................
.....................................................................................................................................
(1)
(c)
The thinning of the ozone layer increases the amount of UV-B radiation that reaches the
Earth’s surface.
Type of Radiation
Wavelength / nm
UV-A
320–380
UV-B
290–320
Based on the information in the table above explain why UV-B rays are more dangerous
than UV-A.
.....................................................................................................................................
.....................................................................................................................................
.....................................................................................................................................
.....................................................................................................................................
.....................................................................................................................................
(3)
(Total 5 marks)
Page 14 of 19
IB Chemistry – SL
Topic 2 Answers
1.
A
[1]
2.
B
[1]
3.
A
[1]
4.
C
[1]
5.
A
[1]
6.
D
[1]
7.
C
[1]
8.
D
[1]
9.
D
[1]
10.
A
[1]
11.
B
[1]
12.
C
[1]
13.
B
[1]
14.
D
[1]
15.
A
[1]
16.
B
[1]
17.
B
[1]
18.
D
[1]
19.
B
[1]
20.
C
[1]
21.
B
[1]
Page 15 of 19
22.
D
[1]
23.
C
[1]
24.
D
[1]
25.
A
[1]
26.
A
[1]
27.
A
[1]
28.
A
[1]
29.
(a)
mole ratio C:H = 85 .6 : 14 .4 = 7.13:4.3;
12 .01 1.01
No penalty for using integer atomic masses.
empirical formula is CH 2 ;
(b)
(i)
2
1.01  10 2 kPa (0.399 dm 3 );
mass
;
number of moles of gas n = PV =
;
RT molar mass
8.314 J (273 K)
mol K
1.00 g
= 56.3 (g mol–1 )
0.017 mol
2
OR
molar mass is the mass of the molar3 volu me at STP;
22.4 dm
= 1.00  22.4 = 56.1 (g mol–1 )
0.399
Accept answers in range 56.0 to 56.3.
Accept two, three or four significant figures.
(ii)
C4 H8 ;
1
No ECF.
[5]
30.
7 protons, 8 neutrons, 10 electrons;
Award [2] for three correct and [1] for two correct.
2
[2]
31.
(a)
(i)
ionization, acceleration, deflection/separation;
Award [1] for all three names and [1] for correct order.
Award [1] for two names in correct order.
(ii)
ionization: sample bombarded with high-energy or high-speed
electrons/OWTTE;
acceleration: electric field/oppositely charged plates;
deflection: (electro)magnet/magnetic field;
2
3
Page 16 of 19
(b)
(i)
(ii)
average or (weighted) mean of masses of all isotopes of an element;
relative to (one atom of) 12 C;
Both marks available from a suitable expression.
2
Ar = (700.2260) + (720.2545) + (740.3673) + (760.1522);
= 72.89;
No other final answer acceptable.
Award [2] for correct final answer.
2
[9]
32.
(i)
number of protons in the nucleus/atom;
Do not accept protons and electrons.
1
(ii)
number of protons and neutrons in the nucleus/atom;
1
[2]
33.
Si
2.8.4/2,8,4;
3−
P
2.8.8/2,8,8;
2
[2]
34.
16 protons and 17 neutrons and 18 electrons;
1
[1]
35.
Al − 2,8,3;
N − 2,5;
F − 2,7;
2
Award [2] for three correct, [1] for two or one correct.
Accept correct configuration using s,p,d notation.
[2]
36.
Ar(Cl) = 35.45 =
35
35x + 37(100 − x)
;
100
Cl = 77.5% and 37 Cl = 22.5%;
2
[2]
37.
(a)
(i)
(ii)
(b)
to produce positively charged ions;
by the bombardment of fast moving electrons;
magnetic field at right angles to path of ions/accept suitably
labelled diagram;
moves ions in curve path/deflects ions;
dependent on mass/charge ratio;
Award [1] each for any 2 points.
2
2
(iii)
acceleration of the ions by electric field/towards negative plate/cathode;
1
(i)
atoms with the same number of protons/positive charges/atomic
number but different number of neutrons/mass number;
1
Page 17 of 19
(ii)
85x + 87(100 − x)
;
100
Accept other valid mathematical alternatives.
Ar(Rb) = 85.47 =
85
(iii)
(iv)
Rb = 76.5 and 87 Rb = 23.5%;
mass; density;
boiling point;
melting point;
rate of diffusion in the gas phase;
enthalpy of vaporization;
enthalpy of fusion;
rate of reaction in the gas/liquid phase;
Any two for one mark
2
1
Si: 1s2 2s2 2p6 3s2 3p2 ;
Fe3+: 1s2 2s2 2p6 3s2 3p6 3d5 ;
P 3− : 1s2 2s2 2p6 3s2 3p6 ;
Allow [1 max] for 3 correct abbreviated structures using noble
gas symbols.
38.
(i)
(ii)
1
the mass of C-12 isotope/
12
average mass of an atom on a scale where one atom of C-12 has a mass of 12/sum
of the weighted average mass of isotopes of an element compared to
C-12/OWTTE;
Award no mark if ‘element’ is used in place of ‘atom’
3
ratio of average mass of an atom to
63
Cu (more abundant) since Ar (Cu) is closer in mass to 63;
Explanation needed for mark
1
1
[2]
39.
(a)
(b)
(c)
(d)
atom of same element/same number of protons but with different
mass number/number of neutrons;
protons
electrons
neutrons
51
23
1
23
23
27
Three correct [2], two correct [1].
2
V /51 nearer to Ar value of 50.94;
1
carbon, 12/12 C;
1
[5]
40.
(a)
(b)
mass/density/for gases: rate of effusion or diffusion/melting point/
boiling point
Do not accept mass number.
1
if 35 Cl = x, then (x = 35.00) + (1 – x) 37.00 = 35.45
Award [1] for set up.
Page 18 of 19
therefore, x = 0.775;
35
Cl = 77.5% and 37 Cl = 22.5%;
(need both for mark);
2
[3]
41.
(a)
(b)
(c)
atoms of the same element/same number of protons/same atomic number;
having different numbers of neutrons/different (mass number);
Award only [1] max if reference made to elements but
not atoms.
relative atomic mass = 36  0.337 + 38  0.0630 + 40  99 .6 ;
100
23 electrons;
26 protons;
30 neutrons;
Award [2] for three correct, [1] for two correct.
2
2
2
[6]
42.
(a)
an atom of 79 Br
an ion of 81 Br–
protons
35
35
;
neutrons
44
46
;
electrons
35
36
;
3
(b)
79
(c)
(i)
2,8,8,2/2.8.8.2;
(ii)
CaBr2 ;
Br because Ar is closer to 79/OWTTE;
1
1
1
[6]
43.
(a)
12 protons and 13 neutrons and 11 electrons;
1
(b)
electric field/oppositely charged plates/potential difference/OWTTE;
1
(c)
25
Mg+;
greater m/z value/less highly charged ions need stronger fields to
deflect them/OWTTE;
Do not accept greater mass with no reference to charge, or
greater mass and smaller charge.
2
[4]
44.
(a)
(b)
(c)
IV < I < II < III/
ultra violet radiation < yellow light < red light < infrared radiation;
1
A continuous spectrum has all colours/wavelengths/frequencies whereas
a line spectrum has only (lines of) sharp/discrete/specific colours/
wavelengths/frequencies;
1
UV-B radiation has shorter wavelength;
hence, has higher energy;
increases risk of damage to skin cells/OWTTE/causes cancer;
3
[5]
Page 19 of 19
IB Chemistry – HL
Topic 2 Questions
1.
What is the electron configuration for an atom with Z = 22?
A.
1s2 2s2 2p6 3s2 3p6 3d4
B.
1s2 2s2 2p6 3s2 3p6 4s2 4p2
C.
1s2 2s2 2p6 3s2 3p6 3d2 4p2
D.
1s2 2s2 2p6 3s2 3p6 4s2 3d2
(Total 1 mark)
2.
What is the total number of p orbitals containing one or more electrons in germanium (atomic
number 32)?
A.
2
B.
3
C.
5
D.
8
(Total 1 mark)
3.
How many electrons are there in all the d orbitals in an atom of xenon?
A.
10
B.
18
C.
20
D.
36
(Total 1 mark)
4.
Which is correct about the element tin (Sn) (Z = 50)?
Number of electrons in
highest main energy level
Number of main energy levels
containing electrons
4
4
A.
14
4
B.
4
5
C.
14
5
D.
(Total 1 mark)
Page 1 of 5
5.
What is the total number of electrons in p orbitals in an atom of iodine?
A.
5
B.
7
C.
17
D.
23
(Total 1 mark)
6.
A transition metal ion X2+ has the electronic configuration [Ar]3d9 . What is the atomic number
of the element?
A.
27
B.
28
C.
29
D.
30
(Total 1 mark)
7.
How many orbitals are there in the n = 3 level of an atom?
A.
3
B.
5
C.
7
D.
9
(Total 1 mark)
8.
What is the electron configuration for the copper(I) ion, (Z = 29)?
A.
[Ar]4s2 3d9
B.
[Ar]4s1 3d10
C.
[Ar]4s1 3d9
D.
[Ar]3d10
(Total 1 mark)
9.
(i)
State the full electron configuration for argon.
……………………………………………………………………………………………
(1)
(ii)
Give the formulas of two oppositely charged ions which have the same electron
configuration as argon.
……………………………………………………………………………………………
(2)
(Total 3 marks)
Page 2 of 5
10.
(a)
Use the Aufbau principle to write the electron configuration of an atom of germanium.
...................................................................................................................................
(1)
(b)
The successive ionization energies of germanium are shown in the following table:
5th
4th
3rd
2nd
1st
8950
4390
3300
1540
760
(i)
Ionization energy /
kJ mol–1
Identify the sub-level from which the electron is removed when the first ionization
energy of germanium is measured.
.........................................................................................................................
(1)
(ii)
Write an equation, including state symbols, for the process occurring when
measuring the second ionization energy of germanium.
.........................................................................................................................
(1)
(iii)
Explain why the difference between the 4th and 5th ionization energies is much
greater than the difference between any two other successive values.
.........................................................................................................................
.........................................................................................................................
.........................................................................................................................
.........................................................................................................................
(2)
(Total 5 marks)
11.
(i)
Explain why successive ionization energies of an element increase.
...................................................................................................................................
...................................................................................................................................
(1)
(ii)
Explain how successive ionization energies account for the existence of three main
energy levels in the sodium atom.
...................................................................................................................................
...................................................................................................................................
...................................................................................................................................
...................................................................................................................................
...................................................................................................................................
...................................................................................................................................
(3)
(Total 4 marks)
Page 3 of 5
IB Chemistry – HL
Topic 2 Answers
1.
D
[1]
2.
D
[1]
3.
C
[1]
4.
C
[1]
5.
D
[1]
6.
C
[1]
7.
D
[1]
8.
D
[1]
9.
(i)
1s2 2s2 2p6 3s2 3p6 ;
Do not accept [Ne] 3s2 3p6 or 2, 8, 8.
(ii)
K+ /Ca2+ /Sc3+/Ti4+;
Cl– /S2– /P 3– ; 2
Accept other suitable pairs of ions.
1
[3]
10.
(a)
1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p2 /[Ar] 4s2 3d10 4p2 ;
Do not penalize for interchanging 4s2 and 3d10 .
(b)
(i)
(ii)
(iii)
(4)p;
Ge+(g) → Ge2+(g) + e−;
Do not penalize for e −(g).
Accept loss of electron on LHS.
5th electron removed from energy level closer to nucleus/5th electron
removed from 3rd energy level and 4th electron from 4th energy
level/OWTTE;
attraction by nucleus or protons greater (for electrons closer to
nucleus)/OWTTE;
1
1
1
2
[5]
11.
(i)
(ii)
same nuclear charge, fewer electrons (thus more energy required to
remove successive electrons)/harder to remove an electron from
an ion with increasing positive charge/nucleus has greater effect
on smaller number of electrons/OWTTE;
1
large increases in IE when 2nd and 10th electron removed; thus,
1st electron further from nucleus than 2nd electron; and 9th electron
Page 4 of 5
further from nucleus than 10th electron; large increases indicate
changes in main energy levels/OWTTE;
OR
outermost/3p electron has low IE because it is far/furthest from the nucleus;
electron(s) in second shell/2p electrons are much closer (to nucleus) and
need much more energy to remove/IE much higher/very high/there is a
big jump in IE;
electron(s) in first/innermost shell/1s electrons are even closer (to nucleus)
and need much more energy to remove (than those in second shell/2s or 2p
electrons);
3
[4]
Page 5 of 5
IB Chemistry – HL
Topic 3 Questions
1.
Which of the following salts form coloured solutions when dissolved in water?
I.
ScCL3
II.
FeCl3
III.
NiCl2
IV.
ZnCl2
A.
I and II only
B.
II and III only
C.
III and IV only
D.
I, II, III and IV
(Total 1 mark)
2.
Which is an essential feature of a ligand?
A.
a negative charge
B.
an odd number of electrons
C.
the presence of two or more atoms
D.
the presence of a non-bonding pair of electrons
(Total 1 mark)
3.
Which equation represents the third ionization energy of an element M?
A.
M+(g) → M4+(g) + 3e–
B.
M2+(g) → M3+(g) + e–
C.
M(g) → M3+(g) + 3e–
D.
M3+(g) → M4+(g) + e–
(Total 1 mark)
4.
Which electrons are lost by an atom of iron when it forms the Fe 3+ ion?
A.
One s orbital electron and two d orbital electrons
B.
Two s orbital electrons and one d orbital electron
C.
Three s orbital electrons
D.
Three d orbital electrons
(Total 1 mark)
Page 1 of 9
5.
Which properties are typical of d-block elements?
I.
complex ion formation
II.
catalytic behaviour
III.
colourless compounds
A.
I and II only
B.
I and III only
C.
II and III only
D.
I, II and III
(Total 1 mark)
6.
Which combination of ion charge and ion size produces the greatest lattice enthalpy?
A.
High charge, large size
B.
High charge, small size
C.
Low charge, small size
D.
Low charge, large size
(Total 1 mark)
7.
Which salts form coloured solutions when dissolved in water?
I.
II.
III.
FeCl3
NiCl2
ZnCl2
A.
I and II only
B.
I and III only
C.
II and III only
D.
I, II and III
(Total 1 mark)
8.
Which combination is correct for the complex ion in [Co(NH 3 )4 (H2O)Cl]Br?
Oxidation state of cobalt
Shape of the complex ion
Overall charge of the
complex ion
A.
+2
Octahedral
+2
B.
+3
Octahedral
–1
C.
+2
Octahedral
+1
D.
+2
Tetrahedral
+1
(Total 1 mark)
Page 2 of 9
9.
Define the term ligand. Cu2+(aq) reacts with ammonia to form the complex ion
[Cu(NH3 )4 ]2+. Explain this reaction in terms of an acid-base theory, and outline the bonding in
the complex ion formed between Cu2+ and NH3 .
................................................................................................................................................
................................................................................................................................................
................................................................................................................................................
................................................................................................................................................
................................................................................................................................................
..................................................................................................................................... ...........
................................................................................................................................................
................................................................................................................................................
................................................................................................................................................
................................................................................................................................................
(Total 4 marks)
10.
By reference to the structure and bonding in the compounds NaCl and SiCl4
(i)
state and explain the differences in conductivity in the liquid state.
.....................................................................................................................................
.....................................................................................................................................
.....................................................................................................................................
.....................................................................................................................................
(3)
(ii)
predict an approximate pH value for a solution formed by adding each compound
separately to water.
.....................................................................................................................................
.....................................................................................................................................
.....................................................................................................................................
.....................................................................................................................................
.....................................................................................................................................
.....................................................................................................................................
(4)
(Total 7 marks)
Page 3 of 9
11.
Two characteristics of the d-block (transition) elements are that they exhibit variable oxidation
states and form coloured compounds.
(i)
State two possible oxidation states for iron and explain these in terms of electron
arrangements.
.....................................................................................................................................
.....................................................................................................................................
.....................................................................................................................................
.....................................................................................................................................
.....................................................................................................................................
.....................................................................................................................................
(2)
(ii)
Explain why many compounds of d-block (transition) elements are coloured.
.....................................................................................................................................
.....................................................................................................................................
.....................................................................................................................................
.....................................................................................................................................
.....................................................................................................................................
.....................................................................................................................................
(3)
(Total 5 marks)
12.
Silicon tetrachloride, SiCl4 , reacts with water to form an acidic solution.
(i)
Explain why silicon tetrachloride has a low melting point.
...................................................................................................................................
...................................................................................................................................
(2)
(ii)
Write an equation for the reaction of silicon tetrachloride with water.
...................................................................................................................................
(1)
(Total 3 marks)
13.
Magnesium chloride and silicon(IV) chloride have very different properties.
(i)
Give the formula and physical state at room temperature of each chloride.
(2)
(ii)
State the conditions under which, if at all, each chloride conducts electricity.
(2)
Page 4 of 9
(iii)
Each chloride is added to water in separate experiments. Suggest an approximate pH
value for the solution formed, and write an equation for any reaction that occurs.
(3)
(Total 7 marks)
14.
The elements in the d-block in the periodic table have several characteristics in common.
(i)
Give the electronic configuration of Ni2+.
(1)
(ii)
Explain what is meant by a ligand, and describe the type of bond formed between a ligand
and a d-block element.
(2)
(iii)
Determine the oxidation numbers of copper in the species
[Cu(NH3 )4 ]2+ and [CuCl4 ]2–
(2)
(iv)
Explain why the species in (iii) are coloured.
(3)
(v)
Identify the d-block element used as a catalyst in the Haber process and write an equation
for the reaction occurring.
(2)
(Total 10 marks)
15.
(i)
Explain why complexes of Zn2+ are colourless whereas complexes containing Cu2+ are
coloured.
(3)
(ii)
Give the formula and describe the shape of the complex ion formed between Fe 3+ and the
ligand CN– .
(2)
(Total 5 marks)
16.
Consider the transition metal complex, K 3 [Fe(CN)6 ].
(i)
Define the term ligand, and identify the ligand in this complex.
(1)
(ii)
Write the full electron configuration and draw the orbital box diagram of iron in its
oxidation state in this complex, and hence, determine the number of unpaired electrons in
this state.
(3)
(iii)
Explain why many transition metal d-block complexes are coloured.
(3)
(Total 7 marks)
17.
By reference to the structure and bonding in NaCl and SiCl4 :
(i)
State and explain the differences in electrical conductivity in the liquid state.
(3)
(ii)
Predict an approximate pH value for the solutions formed by adding each compound
separately to water. Explain your answer.
(4)
Page 5 of 9
(Total 7 marks)
18.
Elements with atomic number 21 to 30 are d-block elements.
(a)
Identify which of these elements are not considered to be typical transition elements.
.....................................................................................................................................
.....................................................................................................................................
(1)
(b)
Complex ions consist of a central metal ion surrounded by ligands. Define the term
ligand.
.....................................................................................................................................
.....................................................................................................................................
.....................................................................................................................................
(2)
(c)
Complete the table below to show the oxidation state of the transition element.
(3)
ion
Cr2 O7 2–
[CuCl4 ]2–
[Fe(H2 O)6 ]3+
oxidation state
(d)
Identify two transition elements used as catalysts in industrial processes, stating the
process in each case.
.....................................................................................................................................
.....................................................................................................................................
.....................................................................................................................................
.....................................................................................................................................
(2)
(e)
Apart from the formation of complex ions and apart from their use as catalysts, state two
other properties of transition elements.
.....................................................................................................................................
.....................................................................................................................................
(2)
(Total 10 marks)
Page 6 of 9
IB Chemistry – HL
Topic 3 Answers
1.
B
[1]
2.
D
3.
B
[1]
[1]
4.
B
[1]
5.
A
[1]
6.
B
[1]
7.
A
[1]
8.
C
[1]
9.
ligand: a molecule or ion that can bond to a (central) metal ion
(to form a complex);
NH3 : Lewis base and Cu2+: Lewis acid (need both for mark);
each NH3 /ligand donates an electron pair (to Cu2+);
forming coordinate covalent/dative covalent bond;
4
[4]
10.
(i)
(ii)
NaCl conducts and SiCl4 does not;
NaCl ionic and SiCl4 covalent;
ions can move in liquid (in NaCl)/OWTTE;
3
NaCl pH = 7;
salt of strong acid and strong base/Na + and Cl– not hydrolysed;
SiCl4
pH = 0 to 3;
HCl is formed/strong acid formed;
4
[7]
11.
(i)
(ii)
+2 and +3/Fe 2+ and Fe3+;
both s electrons are lost giving Fe 2+ and one more d electron is
also lost to form Fe 3+;
presence of unpaired electrons;
the d orbitals are split into two energy levels;
electrons move between these energy levels;
electrons can absorb energy from light of visible wavelength
/OWTTE;
Award [1] each for any three.
2
3
[5]
Page 7 of 9
12.
(i)
(ii)
van der Waals’ forces (between molecules);
Accept London or dispersion forces or temporary dipole-dipole
attractions.
(these forces are) weak/easily overcome;
2
SiCl4 + 4H2 O → Si(OH)4 + 4HCl;
Ignore state symbols, accept SiO 2.2H2O or H4 SiO4 as product.
1
[3]
13.
(i)
(ii)
(iii)
MgCl2 and SiCl4 ;
MgCl2 solid and SiCl4 liquid;
2
MgCl2 (conducts electricity) when molten/dissolved in water;
SiCl4 does not conduct (under any conditions);
2
MgCl2 pH value in range 5.0 to 6.9/just under 7;
SiCl4 pH value in range 0 to 3;
SiCl4 + 4H2 O → Si(OH)4 + 4HCl/SiO2 .2H2 O + 4HCl;
Do not accept SiCl4 + 2H2O → SiO2 + 4HCl.
3
[7]
14.
(i)
Ni2+ 1s2 2s2 2p6 3s2 3p6 3d8 / [Ar]3d8 ;
1
(ii)
species with lone pair of electrons used to bond with the ion;
co-ordinate bond/dative (covalent) bond;
2
+2;
+2;
2
(iii)
Accept 2+ but not 2 or II.
(iv)
(v)
d orbitals/sub-levels (in complexes) split (into two sets at different
energy levels);
energy difference corresponds to frequency/wavelength of (part of)
visible light;
part of visible spectrum absorbed by electrons;
when they move between energy levels;
OWTTE for all of the above.
Award [1] each for any two of the last three.
iron;
N2 + 3H2
2NH3 ;
No penalty for →.
3
2
[10]
15.
(i)
(ii)
Zn2+ has full d sub-shell / Zn2+ does not have partially filled d sub-shell/
Cu2+ has partially filled d sub-shell/orbitals;
d orbitals are split (into two sets of different energy levels);
colour due to electron transition between (split) d orbitals;
[Fe(CN)6 ]3−;
octahedral/suitable diagram;
Accept square bipyramidal
3
2
[5]
Page 8 of 9
16.
(i)
(ii)
an ion or molecule, with a lone pair of electrons that coordinates to a metal
atom or to a metal ion to form a complex/(OWTTE) and cyanide/CN−;
Fe3+ = 1s2 , 2s2 , 2p 6 , 3s2 , 3p 6 , 3d 5 ;
;
3d





1
[Ar]
5
5 unpaired electrons;
(iii)
3
presence of unpaired electrons;
the d orbitals are split into two energy levels;
electrons move between these energy levels;
absorb energy from light of visible wavelength/OWTTE;
max
Award [1] each for any three.
3
[7]
17.
(i)
(ii)
NaCl conducts and SiCl4 does not;
NaCl ionic and SiCl4 covalent;
ions can move in liquid (in NaCl);
3
NaCl pH = 7;
salt of strong acid and strong base/Na + and Cl− not hydrolysed;
SiCl4 pH = 0 to 3;
HCl is formed/strong acid formed;
4
[7]
18.
(a)
scandium and zinc/Sc and Zn;
Both needed for the mark.
Accept copper/Cu if given in addition to Sc and Zn i.e. all three
needed for the mark.
(b)
species/neutral molecules/anions which contain a non-bonding pair
of electrons; able to form coordinate/dative covalent bonds;
1
2
(c)
[Fe(H2 O)6 ]3+
+3
2[CuCl4 ]2–
+2
Cr2 O7 2–
+6
ion
oxidation state
Accept 6+, 2+, 3+. If given as 6, 2, 3 or (VI), (II), (III),
Award [2] only.
(d)
(e)
V/V2 O5 in the contact process;
Fe in the Haber process;
Ni in the conversion of alkenes to alkanes/hydrogenation reactions;
Award [1] each for any two.
Accept any other suitable examples.
variable oxidation states; coloured compounds;
Accept any other suitable examples.
3
2 max
2
[10]
Page 9 of 9
IB Chemistry – SL
Topic 3 Questions
1.
Which pair of elements reacts most readily?
A.
Li + Br2
B.
Li + Cl2
C.
K + Br2
D.
K + Cl2
(Total 1 mark)
2.
Which of the following properties of the halogens increase from F to I?
I.
Atomic radius
II.
Melting point
III.
Electronegativity
A.
I only
B.
I and II only
C.
I and III only
D.
I, II and III
(Total 1 mark)
3.
Which pair would react together most vigorously?
A.
Li and Cl2
B.
Li and Br2
C.
K and Cl2
D.
K and Br2
(Total 1 mark)
4.
For which element are the group number and the period number the same?
A.
Li
B.
Be
C.
B
D.
Mg
(Total 1 mark)
Page 1 of 20
5.
Which of the physical properties below decrease with increasing atomic number for both the
alkali metals and the halogens?
I.
Atomic radius
II.
Ionization energy
III.
Melting point
A.
I only
B.
II only
C.
III only
D.
I and III only
(Total 1 mark)
6.
Rubidium is an element in the same group of the periodic table as lithium and sodium.
It is likely to be a metal which has a
A.
high melting point and reacts slowly with water.
B.
high melting point and reacts vigorously with water.
C.
low melting point and reacts vigorously with water.
D.
low melting point and reacts slowly with water.
(Total 1 mark)
7.
When the following species are arranged in order of increasing radius, what is the correct
order?
A.
Cl– , Ar, K+
B.
K+, Ar , Cl–
C.
Cl– , K+, Ar
D.
Ar, Cl– , K+
(Total 1 mark)
8.
What increases in equal steps of one from left to right in the periodic table for the elements
lithium to neon?
A.
the number of occupied electron energy levels
B.
the number of neutrons in the most common isotope
C.
the number of electrons in the atom
D.
the atomic mass
(Total 1 mark)
Page 2 of 20
9.
Which property decreases down group 7 in the periodic table?
A.
atomic radius
B.
electronegativity
C.
ionic radius
D.
melting point
(Total 1 mark)
10.
Which properties are typical of most non-metals in period 3 (Na to Ar)?
I.
They form ions by gaining one or more electrons.
II.
They are poor conductors of heat and electricity.
III.
They have high melting points.
A.
I and II only
B.
I and III only
C.
II and III only
D.
I, II and III
(Total 1 mark)
11.
A potassium atom has a larger atomic radius than a sodium atom. Which statement about
potassium correctly explains this difference?
A.
It has a larger nuclear charge.
B.
It has a lower electronegativity.
C.
It has more energy levels occupied by electrons.
D.
It has a lower ionization energy.
(Total 1 mark)
12.
Which factors lead to an element having a low value of first ionization energy?
I.
large atomic radius
II.
high number of occupied energy levels
III.
high nuclear charge
A.
I and II only
B.
I and III only
C.
II and III only
D.
I, II and III
(Total 1 mark)
Page 3 of 20
13.
Which statement about electronegativity is correct?
A.
Electronegativity decreases across a period.
B.
Electronegativity increases down a group.
C.
Metals generally have lower electronegativity values than non-metals.
D.
Noble gases have the highest electronegativity values.
(Total 1 mark)
14.
Which statement is correct for a periodic trend?
A.
Ionization energy increases from Li to Cs.
B.
Melting point increases from Li to Cs.
C.
Ionization energy increases from F to I.
D.
Melting point increases from F to I.
(Total 1 mark)
15.
Which compound of an element in period 3 reacts with water to form a solution with a pH
greater than 7?
A.
SiO2
B.
SiCl4
C.
NaCl
D.
Na2 O
(Total 1 mark)
16.
Which equation represents the first ionization energy of fluorine?
A.
F(g) + e– → F– (g)
B.
F– (g) → F(g) + e–
C.
F+(g) → F(g) + e–
D.
F(g) → F+(g) + e–
(Total 1 mark)
17.
Which statement is correct for the halogen group?
A.
Halide ions are all reducing agents, with iodide ions being the weakest.
B.
Halogens are all oxidizing agents, with chlorine being the strongest.
C.
Chloride ions can be oxidized to chlorine by bromine.
D.
Iodide ions can be oxidized to iodine by chlorine.
(Total 1 mark)
Page 4 of 20
18.
Which of the following statements are correct?
I.
II.
III.
The melting points decrease from Li → Cs for the alkali metals.
The melting points increase from F → I for the halogens.
The melting points decrease from Na → Ar for the period 3 elements.
A.
I and II only
B.
I and III only
C.
II and III only
D.
I, II and III
(Total 1 mark)
19.
Which element is a transition metal?
A.
Ca
B.
Cr
C.
Ge
D.
Se
(Total 1 mark)
20.
When Na, K, and Mg are arranged in increasing order of atomic radius (smallest first), which
order is correct?
A.
Na, K, Mg
B.
Na, Mg, K
C.
K, Mg, Na
D.
Mg, Na, K
(Total 1 mark)
21.
Which oxides produce an acidic solution when added to water?
I.
SiO2
II.
P 4 O6
III.
SO2
A.
I and II only
B.
I and III only
C.
II and III only
D.
I, II and III
(Total 1 mark)
Page 5 of 20
22.
Which series is arranged in order of increasing radius?
A.
Ca2+ < Cl– < K+
B.
K+ < Ca2+ < Cl–
C.
Ca2+ < K+ < Cl–
D.
Cl– < K+ < Ca2+
(Total 1 mark)
23.
Describe the acid-base character of the oxides of the period 3 elements Na to Ar. For sodium
oxide and sulfur trioxide, write balanced equations to illustrate their acid-base character.
........................................................................................................................................................
........................................................................................................................................................
........................................................................................................................................................
........................................................................................................................................................
........................................................................................................................................................
(Total 4 marks)
24.
Table 6 of the Data Booklet lists melting points of the elements. Explain the trend in the melting
points of the alkali metals, halogens and period 3 elements.
........................................................................................................................................................
........................................................................................................................................................
........................................................................................................................................................
........................................................................................................................................................
........................................................................................................................................................
........................................................................................................................................................
........................................................................................................................................................
........................................................................................................................................................
........................................................................................................................................................
........................................................................................................................................................
........................................................................................................................................................
........................................................................................................................................................
........................................................................................................................................... .............
........................................................................................................................................................
........................................................................................................................................................
(Total 8 marks)
Page 6 of 20
25.
(i)
Explain how the first ionization energy of K compares with that of Na and Ar.
…………………………………………………………………………………………..
…………………………………………………………………………………………..
…………………………………………………………………………………………..
…………………………………………………………………………………………..
…………………………………………………………………………………………..
…………………………………………………………………………………………..
(3)
(ii)
Explain the difference between the first ionization energies of Na and Mg.
…………………………………………………………………………………………..
…………………………………………………………………………………………..
…………………………………………………………………………………………..
…………………………………………………………………………………………..
…………………………………………………………………………………………..
…………………………………………………………………………………………..
(4)
(iii)
Suggest why much more energy is needed to remove an electron from Na + than from
Mg+.
…………………………………………………………………………………………..
…………………………………………………………………………………………..
…………………………………………………………………………………………..
(1)
(Total 8 marks)
26.
Nitrogen is found in period 2 and group 5 of the periodic table.
(i)
Distinguish between the terms period and group.
……………………………………………………………………………………………
……………………………………………………………………………………………
(1)
(ii)
State the electron arrangement of nitrogen and explain why it is found in period 2 and
group 5 of the periodic table.
……………………………………………………………………………………………
……………………………………………………………………………………………
(3)
(Total 4 marks)
Page 7 of 20
27.
Table 8 of the Data Booklet gives the atomic and ionic radii of elements. State and explain the
difference between
(i)
the atomic radius of nitrogen and oxygen.
……………………………………………………………………………………………
……………………………………………………………………………………………
……………………………………………………………………………………………
……………………………………………………………………………………………
(2)
(ii)
the atomic radius of nitrogen and phosphorus.
……………………………………………………………………………………………
……………………………………………………………………………………………
……………………………………………………………………………………………
……………………………………………………………………………………………
(1)
(iii)
the atomic and ionic radius of nitrogen.
……………………………………………………………………………………………
……………………………………………………………………………………………
……………………………………………………………………………………………
……………………………………………………………………………………………
(2)
(Total 5 marks)
28.
State and explain the trends in the atomic radius and the ionization energy
(i)
for the alkali metals Li to Cs.
……………………………………………………………………………………………
……………………………………………………………………………………………
……………………………………………………………………………………………
……………………………………………………………………………………………
……………………………………………………………………………………………
……………………………………………………………………………………………
(4)
Page 8 of 20
(ii)
for the period 3 elements Na to Cl.
……………………………………………………………………………………………
……………………………………………………………………………………………
……………………………………………………………………………………………
……………………………………………………………………………………………
……………………………………………………………………………………………
……………………………………………………………………………………………
(4)
(Total 8 marks)
29.
(i)
Describe three similarities and one difference in the reactions of lithium and potassium
with water.
……………………………………………………………………………………………
……………………………………………………………………………………………
……………………………………………………………………………………………
……………………………………………………………………………………………
……………………………………………………………………………………………
……………………………………………………………………………………………
……………………………………………………………………………………………
……………………………………………………………………………………………
(4)
(ii)
Give an equation for one of these reactions. Suggest a pH value for the resulting solution,
and give a reason for your answer.
……………………………………………………………………………………………
……………………………………………………………………………………………
……………………………………………………………………………………………
(3)
(Total 7 marks)
30.
(a)
Classify each of the following oxides as acidic, basic or amphoteric.
(i)
aluminium oxide
……………………………………………………………………………………
(1)
(ii)
sodium oxide
……………………………………………………………………………………
(1)
Page 9 of 20
(iii)
sulfur dioxide
……………………………………………………………………………………
(1)
(b)
Write an equation for each reaction between water and
(i)
sodium oxide
……………………………………………………………………………………
……………………………………………………………………………………
(1)
(ii)
sulfur dioxide.
……………………………………………………………………………………
……………………………………………………………………………………
(1)
(Total 5 marks)
31.
This question is about Period 3 elements and their compounds.
(a)
Explain, in terms of their structure and bonding, why the element sulfur is a nonconductor of electricity and aluminium is a good conductor of electricity.
...................................................................................................................................
...................................................................................................................................
...................................................................................................................................
...................................................................................................................................
...................................................................................................................................
...................................................................................................................................
(4)
(b)
Explain, in terms of its structure and bonding, why silicon dioxide, SiO 2 , has a high
melting point.
...................................................................................................................................
...................................................................................................................................
...................................................................................................................................
(2)
(Total 6 marks)
32.
Explain why
(i)
the first ionization energy of magnesium is lower than that of fluorine.
(2)
(ii)
magnesium has a higher melting point than sodium.
(3)
(Total 5 marks)
Page 10 of 20
33.
Discuss the acid-base nature of the period 3 oxides. Write an equation to illustrate the reaction
of one of these oxides to produce an acid, and another equation of another of these oxides to
produce a hydroxide.
(Total 5 marks)
34.
Information about the halogens appears in the Data Booklet.
(i)
Explain why the ionic radius of chlorine is less than that of sulfur.
(2)
(ii)
Explain what is meant by the term electronegativity and explain why the electronegativity
of chlorine is greater than that of bromine.
(3)
(Total 5 marks)
35.
(a)
(i)
State the meaning of the term electronegativity and explain why the noble gases are
not assigned electronegativity values.
(2)
(ii)
State and explain the trend in electronegativity across period 3 from Na to Cl.
(2)
(iii)
Explain why Cl2 rather than Br2 would react more vigorously with a solution of I– .
(2)
(b)
State the acid-base properties of the following period 3 oxides.
MgO
Al2 O3
P 4 O6
Write equations to demonstrate the acid-base properties of each compound.
(7)
(Total 13 marks)
36.
(i)
Define the term ionization energy.
(1)
(ii)
Write an equation for the reaction of lithium with water.
(1)
(iii)
State and explain the trend in the ionization energy of alkali metals down the group.
(3)
(iv)
Explain why the electronegativity of phosphorus is greater than that of aluminium.
(2)
(v)
Table 8 in the Data Booklet contains two values for the ionic radius of silicon. Explain,
by reference to atomic structure and electron arrangements, why the two values are very
different.
(4)
(Total 11 marks)
37.
Explain why sulfur has a lower first ionization energy than oxygen, and also a lower first
ionization energy than phosphorus.
(Total 4 marks)
38.
With reference to the types of bonding present in period 3 elements:
(i)
explain why Mg has a higher melting point than Na.
(2)
Page 11 of 20
(ii)
explain why Si has a very high melting point.
(2)
(iii)
explain why the other non-metal elements of period 3 have low melting points.
(2)
(Total 6 marks)
39.
Describe the acid-base character of the oxides of the period 3 elements Na to Ar. For sodium
oxide and sulfur trioxide, write balanced equations to illustrate their acid-base character.
(Total 3 marks)
40.
Explain the following statements.
(a)
The first ionization energy of sodium is
(i)
less than that of magnesium.
…………………………………………………………………………………
…………………………………………………………………………………
…………………………………………………………………………………
…………………………………………………………………………………
(2)
(ii)
greater than that of potassium.
…………………………………………………………………………………
…………………………………………………………………………………
(1)
(b)
The electronegativity of chlorine is higher than that of sulfur.
……………………………………………………………………………………….
……………………………………………………………………………………….
……………………………………………………………………………………….
……………………………………………………………………………………….
(2)
(Total 5 marks)
41.
(a)
(i)
Define the term ionization energy.
.........................................................................................................................
.........................................................................................................................
(2)
(ii)
Write an equation, including state symbols, for the process occurring when
measuring the first ionization energy of aluminium.
.........................................................................................................................
(1)
Page 12 of 20
(b)
The first ionization energies of the elements are shown in Table 7 of the Data Booklet.
Explain why the first ionization energy of magnesium is greater than that of sodium.
...................................................................................................................................
...................................................................................................................................
...................................................................................................................................
...................................................................................................................................
(2)
(c)
Lithium reacts with water. Write an equation for the reaction and state two observations
that could be made during the reaction.
...................................................................................................................................
...................................................................................................................................
...................................................................................................................................
(3)
(Total 8 marks)
42.
(a)
State the meaning of the term electronegativity.
....................................................................................................................................
....................................................................................................................................
(1)
(b)
State and explain the trend in electronegativity across period 3 from Na to Cl.
....................................................................................................................................
....................................................................................................................................
....................................................................................................................................
....................................................................................................................................
(2)
(c)
Explain why Cl2 rather than Br2 would react more vigorously with a solution of I – .
....................................................................................................................................
....................................................................................................................................
....................................................................................................................................
....................................................................................................................................
(2)
(Total 5 marks)
Page 13 of 20
IB Chemistry – SL
Topic 3 Answers
1.
D
[1]
2.
B
[1]
3.
C
[1]
4.
B
[1]
5.
B
[1]
6.
C
[1]
7.
B
[1]
8.
C
[1]
9.
B
[1]
10.
A
[1]
11.
C
[1]
12.
A
[1]
13.
C
[1]
14.
D
[1]
15.
D
[1]
16.
D
[1]
17.
D
[1]
18.
A
[1]
19.
B
[1]
20.
D
[1]
21.
C
Page 14 of 20
[1]
22.
C
[1]
23.
oxides of:
Na, Mg: basic;
Al: amphoteric;
Si to Cl: acidic;
Ar: no oxide;
All four correct [2], two or three correct [1].
Na2 O + H2 O → 2NaOH;
SO3 + H2 O → H2 SO4 ;
Must be balanced for marks.
Award marks for alternative correct equations such as SO3 with
NaOH.
4
[4]
24.
alkali metals:
metallic bonding/a bed of cations in a sea of electrons;
as radius increases down the group, valence electrons are further away from
nucleus (and strength of metallic bonding decreases);
halogens:
non-polar/van der Waals’ forces between molecules;
as size increases van der Waals’ forces increase (and melting point increases);
period 3 elements:
increase in melting points of metals (Na, Mg, Al) due to increase in number of
valence electrons and decrease in size/the way atoms are packed as solids;
Award mark just for “increased number of delocalized or valence electrons”.
silicon:
network covalent solid (with very high melting point);
Award mark also for “many or strong covalent bonds”.
P → Ar:
simple molecular (atomic in case of Ar) substances with weak van
der Waals’ forces (and lower melting points);
trend in P 4 , S8 , Cl2 , Ar due to size/mass of particles;
Award mark for “decreasing mass or size”.
Molecular formulae not necessary.
8
[8]
25.
(i) and (ii) marked together.
K less than Na because
electron removed (from K) is from higher energy level/further from
nucleus/in n = 4 compared to n = 3;
this is more important than the extra 8 protons in K/OWTTE;
increase repulsion by extra shell of electrons/greater shielding effect;
so less strongly attracted by nucleus;
K less than Ar because
electron removed (from K) is from higher energy level/further from nucleus/
in n = 4 compare to n = 3;
and has only one more proton;
Page 15 of 20
increase repulsion by extra shell of electrons/greater shielding effect;
so less strongly attracted by nucleus;
Mg greater than Na because
(Mg has) greater nuclear charge/one more proton/12 protons compare to 11;
electron removed is in same (main) higher energy level/shell;
smaller (atomic) radius;
so more strongly attracted by nucleus;
Accept opposite worded arguments, i.e. why Na is
greater than K.
Award [7] for any seven correct but accept less/more
strongly attracted to nucleus once only.
(iii)
second electron in Na removed from n = 2, whereas second electron in Mg
removed from n = 3
7
1
[8]
26.
(i)
(ii)
period is a horizontal row in the periodic table and a group is a
vertical column/OWTTE;
1
2,5;
electrons in two energy levels/shells;
five outer/valence electrons;
3
[4]
27.
(i)
(ii)
(iii)
atomic radius of N > O because O has greater nuclear charge;
greater attraction for the outer electrons/OWTTE;
2
atomic radius of P > N because P has outer electrons in an energy
level further from the nucleus/OWTTE;
1
N3– >N/ionic radius > atomic radius because N 3– has more electrons
than protons; so the electrons are held less tightly/OWTTE;
2
3–
Award [1] for greater repulsion in N due to more
electrons (no reference to protons).
[5]
28.
(i)
(ii)
Li to Cs
atomic radius increases;
because more full energy levels are used or occupied/outer electrons
further from nucleus/outer electrons in a higher shell;
ionization energy decreases;
because the electron removed is further from the nucleus/increased
repulsion by inner-shell electrons;
Accept increased shielding effect.
Na to Cl
atomic radius decreases;
because nuclear charge increases and electrons are added to same
main (outer) energy level;
ionization energy increases;
because nuclear charge increases and the electron removed is closer to the
nucleus/is in the same energy level;
Accept “core charge” for “nuclear charge”.
In (i) and (ii) explanation mark dependent on correct trend.
4
4
[8]
Page 16 of 20
29.
(i)
similarities [3 max]
the metal floats/moves on the surface;
fizzing/effervescence/bubbles; (accept sound is produced)
solution gets hot;
solution becomes alkaline/basic;
they react to form the metal hydroxide;
hydrogen is evolved;
differences [1 max]
flame/hydrogen burns with potassium (and not with lithium)
/reaction faster/more vigorous with potassium/slower or
less vigorous with lithium;
max
(ii)
4
2Li + 2H2 O → 2Li+ + 2OH– + H2 / 2K + 2H2 O → 2K+ + 2OH– + H2 ;
Accept LiOH/KOH.
pH  11;
LiOH/KOH is a strong base/strong alkali/high concentration of OH– ;
3
[7]
30.
(a)
(b)
(i)
aluminium oxide
amphoteric;
(ii)
sodium oxide
basic;
(iii)
sulfur dioxide
acidic;
(i)
Na2 O + H2 O → 2Na+ + 2OH– ;
(ii)
SO2 + H2 O → H2 SO3 ;
Accept NaOH and H+ + HSO3–/2H+ + SO32–.
3
2
[5]
31.
(a)
(b)
sulfur is (simple) molecular;
(contains) covalent bonds/no delocalized electrons/all (outer) electrons used
in bonding;
aluminium contains positive ions and delocalized electrons;
(delocalized) electrons move (when voltage applied or current flows);
silicon dioxide is macromolecular/giant covalent;
many/strong covalent bonds must be broken;
Award max [1] if no mention of covalent.
Do not accept weakened instead of broken.
4
2
[6]
32.
(i)
electron removed from higher energy level/further from nucleus/
greater atomic radius;
increased repulsion by extra inner shell electrons/increased shielding
effect;
(ii)
2
Mg has twice as many/more delocalized electrons (compared to Na);
the ionic charge is twice as big/greater in Mg (than Na);
Page 17 of 20
(electrostatic) attraction between ions and electrons is much greater;
3
[5]
33.
oxides of
Na, Mg are basic
Al is amphoteric
Si, P, S and Cl are acidic
Award 7 correct [3], 6/5 correct [2] and 4/3 correct [1].
SO2 + H2 O → H2 SO3 /SO3 + H2 O → H2 SO4/
P 4 O10 + 6H2 O → 4H3 PO4/P 4O6 + 6H2 O → 4H3 PO3 ;
Na2 O + H2 O → 2NaOH/MgO + H2 O → Mg(OH)2 ;
Accept equation using P2 O3 or P2 O5.
5
[5]
34.
(i)
(ii)
(chlorine has) an extra proton/more protons/greater nuclear charge/
17+ compared to 16+;
outer electrons attracted more strongly;
2
ability of atom to attract bonding pair of electrons/electrons in a
covalent bond;
chlorine has a smaller radius/(electrons) closer to nucleus/in lower
energy level;
repelled by fewer inner electrons/decreased shielding effect;
3
[5]
35.
(a)
(i)
(ii)
(iii)
(b)
the ability of an atom to attract a bonding pair of electrons;
inert/do not react/do not attract electrons/stable electron
configuration/full outer electron shell/do not form bonds;
electronegativity increases (along period 3 from Na to Cl);
number of protons increases/nuclear charge increase/core charge
increase/size of atom decreases;
Do not accept “greater nuclear attraction”.
Cl2 stronger oxidising agent;
Cl2 has greater attraction for electrons/has a higher electron affinity;
Accept converse statements for Br 2 .
2
2
2
MgO − basic oxide/alkali;
MgO + 2HCl → MgCl2 + H2 O/MgO + H2 O → Mg(OH)2 ;
Al2 O3 − amphoteric oxide/acidic and basic oxide;
Al2 O3 + 6HCl → 2AlCl3 + 3H2 O;
Al2 O3 + 2OH− + 3H2 O → 2Al(OH)4 – /Al2 O3 + 2OH− → 2AlO2 – + H2 O;
P 4 O6 − acidic oxide;
P 4 O6 + 6H2 O → 4H3 PO3;
All equations must be balanced.
7
[13]
36.
(i)
minimum energy required to remove one (mole of) electron(s) from
(one mole of) (a) gaseous atom(s)/OWTTE;
1
Page 18 of 20
(ii)
(iii)
(iv)
(v)
2Li(s) + 2H2 O(1) → 2LiOH(aq) + H2 (g)/Li(s) + H2 O(1) → LiOH(aq)
+ 1/2H2 (g);
State symbols not required
1
(ionization energy) decreases;
radius increases/valence electrons further away from nucleus/
electron removed from higher shell;
(nuclear charge increases but) shielding/screening effect increases/
more electrons between nucleus and valence electron/lower effective
nuclear charge/Zeff;
3
phosphorus has a higher (effective) nuclear charge/Zeff;
radius of P is smaller;
electron pair/bonding electrons attracted more strongly;
2
both have same number of protons/14 protons/nuclear charge/core charge;
Si4+ formed by electron loss, Si4− formed by electron gain;
Si4+ : 2.8 arrangement/2 (complete) energy levels/electrons in n = 2;
Si4− : 2.8.8 arrangement/3 (complete) energy levels/electrons in n = 3;
explanation of proton : electron ratio;
higher effective nuclear charge/Zeff in Si4+;
4
[11]
37.
IES < IEO:
valence electron in S in n = 3, in O in n = 2/e − further away/S has another
electron shell/atomic radius of S greater than that of O;
less attracted to nucleus/experiences greater screening from inner electrons;
IES < IEP :
electron removed from S is paired;
greater repulsion due to two electrons in the same (p) orbital/paired
electrons in S;
4
[4]
38.
(i)
(ii)
(iii)
Mg has greater nuclear charge/greater charge on cation/more
valence e−/greater number of delocalized electrons/Na has lesser
nuclear charge/lesser charge on cation/less valence e −/lesser number of delocalized
electrons; stronger attraction between cation and delocalized/
free/valence electrons;
If neither mark scored, accept stronger metallic bonding in Mg
for [1 max].
giant/network/lattice/macromolecular structure;
many/strong covalent bonds (need to be broken);
(simple) molecular substances;
weak van der Waals’/dispersion/London forces between molecules;
“Weak intermolecular forces” not sufficient for second mark
2
2
2
[6]
39.
Oxides of: Na and Mg are basic;
Al is amphoteric;
Si to Cl are acidic;
Page 19 of 20
Ar has no oxide;
All four correct award [2], two or three correct award [1].
Na2 O + H2 O → 2NaOH and SO3 + H2 O → H2 SO4 ;
Must be balanced for mark.
Award marks for alternative correct equations such as SO3 with
NaOH.
3
[3]
40.
(a)
(i)
(ii)
(b)
Na has lower nuclear charge/number of protons; electrons
being removed are from same energy level/shell;
or Na has larger radius/electron further from nucleus;
max
Award this mark if both electron arrangements are given.
2
Na electron closer to nucleus/in lower energy level/Na has less
shielding effect;
Allow counter arguments for Mg in (i) and K in (ii).
1
chlorine has a higher nuclear charge;
attracts the electron pair/electrons in bond more strongly;
2
[5]
41.
(a)
(i)
(ii)
(b)
(c)
the (minimum) energy required/needed for the removal of one electron;
from a gaseous/isolated atom;
Al(g) → Al+(g) + e;
Do not penalize the answer if (g) is after e.
greater nuclear charge/greater number of protons/atom radius g is smaller;
stronger attraction (for electron);
2
1
2
2Li + 2H2 O → 2LiOH + H2 ;
Ignore state symbols.
effervescence/fizzing/bubbles/OWTTE;
lithium moves around/decrease in size of piece;
Accept dissolves or disappears.
heat produced;
Award [1] each for any two of last three observations.
3
[8]
42.
(a)
the ability of an element/atom/nucleus to attract a bonding pair of electrons;
(b)
electronegativity increases (along period 3 from Na to Cl);
number of protons increases/nuclear charge increases/core charge increases
/size of atoms decreases;
Do not accept greater nuclear attraction.
(c)
Cl2 is a stronger oxidizing agent/Chlorine’s outer shell closer to nucleus;
Cl2 has greater attraction for electrons/has a higher electron affinity;
Accept converse argument for Br 2 .
1
2
2
[5]
Page 20 of 20
IB Chemistry – SL
Topic 4 Questions
1.
What is the formula for the compound formed by calcium and nitrogen?
A.
CaN
B.
Ca2 N
C.
Ca2 N3
D.
Ca3 N2
(Total 1 mark)
2.
Element X is in group 2, and element Y in group 7, of the periodic table. Which ions will be
present in the compound formed when X and Y react together?
A.
X+ and Y–
B.
X 2+ and Y–
C.
X+ and Y2–
D.
X2– and Y+
(Total 1 mark)
3.
Based on electronegativity values, which bond is the most polar?
A.
B―C
B.
C―O
C.
N―O
D.
O―F
(Total 1 mark)
4.
What is the Lewis (electron dot) structure for sulfur dioxide?
A.
O S O
B.
O S O
C.
O
S
D.
O
S O
O
(Total 1 mark)
1
5.
Which substance is most soluble in water (in mol dm–3 ) at 298 K?
A.
CH3 CH3
B.
CH3 OCH3
C.
CH3 CH2 OH
D.
CH3 CH2 CH2CH2OH
(Total 1 mark)
6.
According to VSEPR theory, repulsion between electron pairs in a valence shell decreases in
the order
A.
lone pair-lone pair > lone pair-bond pair > bond pair-bond pair.
B.
bond pair-bond pair > lone pair-bond pair > lone pair-lone pair.
C.
lone pair-lone pair > bond pair-bond pair > bond pair-lone pair.
D.
bond pair-bond pair > lone pair-lone pair > lone pair-bond pair.
(Total 1 mark)
7.
Which molecule is linear?
A.
SO2
B.
CO2
C.
H2 S
D.
Cl2 O
(Total 1 mark)
8.
Why is the boiling point of PH3 lower than that of NH3 ?
A.
PH3 is non-polar whereas NH3 is polar.
B.
PH3 is not hydrogen bonded whereas NH 3 is hydrogen bonded.
C.
Van der Waals’ forces are weaker in PH 3 than in NH3 .
D.
The molar mass of PH3 is greater than that of NH 3 .
(Total 1 mark)
9.
Which molecule is non-polar?
A.
H2 CO
B.
SO3
C.
NF3
D.
CHCl3
(Total 1 mark)
2
10.
What happens when sodium and oxygen combine together?
A.
Each sodium atom gains one electron.
B.
Each sodium atom loses one electron.
C.
Each oxygen atom gains one electron.
D.
Each oxygen atom loses one electron.
(Total 1 mark)
11.
Which statement is correct about two elements whose atoms form a covalent bond with
each other?
A.
The elements are metals.
B.
The elements are non-metals.
C.
The elements have very low electronegativity values.
D.
The elements have very different electronegativity values.
(Total 1 mark)
12.
Which substance has the lowest electrical conductivity?
A.
Cu(s)
B.
Hg(l)
C.
H2 (g)
D.
LiOH(aq)
(Total 1 mark)
13.
When the following bond types are listed in decreasing order of strength (strongest first), what
is the correct order?
A.
covalent  hydrogen  van der Waals’
B.
covalent  van der Waals’  hydrogen
C.
hydrogen  covalent  van der Waals’
D.
van der Waals’  hydrogen  covalent
(Total 1 mark)
14.
Which statement is true for most ionic compounds?
A.
They contain elements of similar electronegativity.
B.
They conduct electricity in the solid state.
C.
They are coloured.
D.
They have high melting and boiling points.
(Total 1 mark)
3
15.
What is the valence shell electron pair repulsion (VSEPR) theory used to predict?
A.
The energy levels in an atom
B.
The shapes of molecules and ions
C.
The electronegativities of elements
D.
The type of bonding in compounds
(Total 1 mark)
16.
Which fluoride is the most ionic?
A.
NaF
B.
CsF
C.
MgF2
D.
BaF2
(Total 1 mark)
17.
Which substance is most similar in shape to NH 3 ?
A.
GaI3
B.
BF3
C.
FeCl3
D.
PBr3
(Total 1 mark)
18.
Which statement is a correct description of electron loss in this reaction?
2Al + 3S → Al2 S3
A.
Each aluminium atom loses two electrons.
B.
Each aluminium atom loses three electrons.
C.
Each sulfur atom loses two electrons.
D.
Each sulfur atom loses three electrons.
(Total 1 mark)
19.
Which molecule has the smallest bond angle?
A.
CO2
B.
NH3
C.
CH4
D.
C2 H4
(Total 1 mark)
4
20.
In which substance is hydrogen bonding present?
A.
CH4
B.
CH2 F2
C.
CH3 CHO
D.
CH3 OH
(Total 1 mark)
21.
Which is a correct description of metallic bonding?
A.
Positively charged metal ions are attracted to negatively charged ions.
B.
Negatively charged metal ions are attracted to positively charged metal ions.
C.
Positively charged metal ions are attracted to delocalized electrons.
D.
Negatively charged metal ions are attracted to delocalized electrons.
(Total 1 mark)
22.
What intermolecular forces are present in gaseous hydrogen?
A.
Hydrogen bonds
B.
Covalent bonds
C.
Dipole-dipole attractions
D.
Van der Waals’ forces
(Total 1 mark)
23.
Which molecule is polar?
A.
CO2
B.
PF3
C.
CH4
D.
BF3
(Total 1 mark)
24.
What are responsible for the high electrical conductivity of metals?
A.
Delocalized positive ions
B.
Delocalized valence electrons
C.
Delocalized atoms
D.
Delocalized negative ions
(Total 1 mark)
5
25.
Which compound has the least covalent character?
A.
SiO2
B.
Na2 O
C.
MgCl2
D.
CsF
(Total 1 mark)
26.
When C2 H4 , C2 H2 and C2 H6 are arranged in order of increasing C–C bond length, what is the
correct order?
A.
C2 H6 , C2 H2 , C2 H4
B.
C2 H4 , C2 H2 , C2 H6
C.
C2 H2 , C2 H4 , C2 H6
D.
C2 H4 , C2 H6 , C2 H2
(Total 1 mark)
27.
Which compound contains both ionic and covalent bonds?
A.
MgCl2
B.
HCl
C.
H2 CO
D.
NH4 Cl
(Total 1 mark)
28.
When the species BF2 +, BF3 and BF4 – are arranged in order of increasing F−B−F bond angle,
what is the correct order?
A.
BF3 , BF4 – , BF2 +
B.
BF4 – , BF3 , BF2 +
C.
BF2 +, BF4 – , BF3
D.
BF2 +, BF3 , BF4 –
(Total 1 mark)
29.
Which species has a trigonal planar shape?
A.
CO3 2–
B.
SO3 2–
C.
NF3
D.
PCl3
6
30.
When C2 H4 , C2 H2 and C2 H6 are arranged in order of increasing C–C bond length, what is the
correct order?
A.
C2 H6 , C2 H2 , C2 H4
B.
C2 H4 , C2 H2 , C2 H6
C.
C2 H2 , C2 H4 , C2 H6
D.
C2 H4 , C2 H6 , C2 H2
(Total 1 mark)
31.
What is the formula for an ionic compound formed between an element, X, from group 2 and an
element, Y, from group 6?
A.
XY
B.
X2 Y
C.
XY2
D.
X2 Y6
(Total 1 mark)
32.
In the molecules N2 H4 , N2 H2 , and N2 , the nitrogen atoms are linked by single, double and triple
bonds, respectively. When these molecules are arranged in increasing order of the lengths of
their nitrogen to nitrogen bonds (shortest bond first) which order is correct?
A.
N2 H 4 , N2 , N 2 H2
B.
N2 H 4 , N2 H2 , N2
C.
N2 H 2 , N2 , N 2 H4
D.
N2 , N 2 H 2 , N 2 H4
(Total 1 mark)
33.
The compounds listed have very similar molar masses. Which has the strongest intermolecular
forces?
A.
CH3 CHO
B.
CH3 CH2 OH
C.
CH3 CH2 F
D.
CH3 CH2 CH3
(Total 1 mark)
7
34.
What is the shape of the CO 3 2– ion and the approximate O–C–O bond angle?
A.
Linear, 180
B.
Trigonal planar, 90
C.
Trigonal planar, 120
D.
Pyramidal, 109
(Total 1 mark)
35.
Which combination of Hvaporization and boiling point is the result of strong intermolecular
forces?
Hvaporization
Boiling Point
A.
large
high
B.
large
low
C.
small
low
D.
small
high
(Total 1 mark)
36.
What is the formula of the compound formed when aluminium reacts with oxygen?
A.
Al3 O2
B.
Al2 O3
C.
AlO2
D.
AlO3
(Total 1 mark)
37.
Which statement is true for compounds containing only covalent bonds?
A.
They are held together by electrostatic forces of attraction between oppositely charged
ions.
B.
They are made up of metal elements only.
C.
They are made up of a metal from the far left of the periodic table and a non-metal from
the far right of the periodic table.
D.
They are made up of non-metal elements only.
(Total 1 mark)
38.
How many electrons are used in the carbon-carbon bond in C2 H2 ?
A.
4
B.
6
C.
10
D.
12
(Total 1 mark)
8
39.
Which compound has the highest boiling point?
A.
CH3 CH2 CH3
B.
CH3 CH2 OH
C.
CH3 OCH3
D.
CH3 CHO
(Total 1 mark)
40.
What type of solid materials are typically hard, have high melting points and poor electrical
conductivities?
I.
II.
III.
Ionic
Metallic
Covalent-network
A.
I and II only
B.
I and III only
C.
II and III only
D.
I, II and III
(Total 1 mark)
41.
The boiling points of the hydrides of the group 6 elements are shown below.
400
300
Boiling point / K
200
100
0
H2 O
(i)
H2 S
H 2 Se
H 2 Te
Explain the trend in boiling points from H 2 S to H2 Te.
……………………………………………………………………………………….
……………………………………………………………………………………….
……………………………………………………………………………………….
(2)
9
(ii)
Explain why the boiling point of water is higher than would be expected from the group
trend.
……………………………………………………………………………………….
……………………………………………………………………………………….
……………………………………………………………………………………….
……………………………………………………………………………………….
(2)
(Total 4 marks)
42.
(i)
State the shape of the electron distribution around the oxygen atom in the water molecule
and state the shape of the molecule.
………………………………………………………………………………………….
………………………………………………………………………………………….
(2)
(ii)
State and explain the value of the HOH bond angle.
…………………………………………………………………………………………..
…………………………………………………………………………………………..
…………………………………………………………………………………………..
…………………………………………………………………………………………..
(2)
(Total 4 marks)
43.
Explain why the bonds in silicon tetrachloride, SiCl4 , are polar, but the molecule is not.
……………………………………………………………………………………………
……………………………………………………………………………………………
……………………………………………………………………………………………
……………………………………………………………………………………………
(Total 2 marks)
44.
The diagrams below represent the structures of iodine, sodium and sodium iodide.
A
B
C
10
(a)
(i)
Identify which of the structures (A, B and C) correspond to iodine, sodium and
sodium iodide.
……………………………………………………………………………………….
……………………………………………………………………………………….
……………………………………………………………………………………….
(1)
(ii)
State the type of bonding in each structure.
……………………………………………………………………………………….
……………………………………………………………………………………….
……………………………………………………………………………………….
(3)
(b)
(i)
Sodium and sodium iodide can both conduct electricity when molten, but only
sodium can conduct electricity when solid. Explain this difference in conductivity
in terms of the structures of sodium and sodium iodide.
……………………………………………………………………………………….
……………………………………………………………………………………….
……………………………………………………………………………………….
……………………………………………………………………………………….
……………………………………………………………………………………….
……………………………………………………………………………………….
(4)
(ii)
Explain the high volatility of iodine compared to sodium and sodium iodide.
……………………………………………………………………………………….
……………………………………………………………………………………….
……………………………………………………………………………………….
……………………………………………………………………………………….
(2)
(Total 10 marks)
45.
(i)
Draw Lewis (electron dot) structures for CO 2 and H2 S showing all valence electrons.
(2)
11
(ii)
State the shape of each molecule and explain your answer in terms of VSEPR theory.
CO2 .............................................................................................................................
.....................................................................................................................................
.....................................................................................................................................
H2 S .............................................................................................................................
.....................................................................................................................................
.....................................................................................................................................
(4)
(iii)
State and explain whether each molecule is polar or non-polar.
.....................................................................................................................................
.....................................................................................................................................
.....................................................................................................................................
.....................................................................................................................................
(2)
(Total 8 marks)
46.
Identify the strongest type of intermolecular force in each of the following compounds.
CH3 Cl ...................................................................................................................................
CH4 .......................................................................................................................................
CH3 OH .......................................................................................................................... .......
(Total 3 marks)
47.
(a)
An important compound of nitrogen is ammonia, NH 3 . The chemistry of ammonia is
influenced by its polarity and its ability to form hydrogen bonds. Polarity can be
explained in terms of electronegativity.
(i)
Explain the term electronegativity.
……………………………………………………………………………………
……………………………………………………………………………………
(2)
(ii)
Draw a diagram to show hydrogen bonding between two molecules of NH 3 .
The diagram should include any dipoles and/or lone pairs of electrons
……………………………………………………………………………………
……………………………………………………………………………………
……………………………………………………………………………………
……………………………………………………………………………………
(3)
12
(iii)
State the H–N–H bond angle in an ammonia molecule.
………………………………………………………………………………………
(1)
(iv)
Explain why the ammonia molecule is polar.
……………………………………………………………………………………
……………………………………………………………………………………
……………………………………………………………………………………
(1)
(b)
Ammonia reacts with hydrogen ions forming ammonium ions, NH 4 +.
(i)
State the H–N–H bond angle in an ammonium ion.
……………………………………………………………………………………
(1)
(ii)
Explain why the H–N–H bond angle of NH3 is different from the H–N–H bond
angle of NH4 +; referring to both species in your answer.
……………………………………………………………………………………
……………………………………………………………………………………
……………………………………………………………………………………
……………………………………………………………………………………
……………………………………………………………………………………
……………………………………………………………………………………
(3)
(Total 11 marks)
48.
State the type of bonding in the compound SiCl4 . Draw the Lewis structure for this compound.
……………………………………………………………………………………………
……………………………………………………………………………………………
……………………………………………………………………………………………
……………………………………………………………………………………………
……………………………………………………………………………………………
……………………………………………………………………………………………
(Total 3 marks)
13
49.
Outline the principles of the valence shell electron pair repulsion (VSEPR) theory.
……………………………………………………………………………………………
……………………………………………………………………………………………
……………………………………………………………………………………………
……………………………………………………………………………………………
……………………………………………………………………………………………
……………………………………………………………………………………………
……………………………………………………………………………………………
……………………………………………………………………………………………
(Total 3 marks)
50.
(i)
Use the VSEPR theory to predict and explain the shape and the bond angle of each of the
molecules SCl2 and C2 Cl2
……………………………………………………………………………………………
……………………………………………………………………………………………
……………………………………………………………………………………………
……………………………………………………………………………………………
……………………………………………………………………………………………
……………………………………………………………………………………………
……………………………………………………………………………………………
……………………………………………………………………………………………
……………………………………………………………………………………………
……………………………………………………………………………………………
(6)
(ii)
Deduce whether or not each molecule is polar, giving a reason for your answer.
……………………………………………………………………………………………
……………………………………………………………………………………………
……………………………………………………………………………………………
……………………………………………………………………………………………
……………………………………………………………………………………………
……………………………………………………………………………………………
(3)
(Total 9 marks)
14
51.
Draw a Lewis structure of a water molecule, name the shape of the molecule and state and
explain why the bond angle is less than the bond angle in a tetrahedral molecule such as
methane.
(Total 4 marks)
52.
Predict and explain the order of the melting point for propanol, butane and propanone with
reference to their intermolecular forces.
(Total 4 marks)
53.
The elements sodium, aluminium, silicon, phosphorus and sulfur are in period 3 of the periodic
table.
Describe the metallic bonding present in aluminium and explain why aluminium has a higher
melting point than sodium.
..............................................................................................................................................
..............................................................................................................................................
..............................................................................................................................................
............................................................................................................................................ ..
..............................................................................................................................................
..............................................................................................................................................
(Total 3 marks)
54.
Draw the Lewis structure of NCl3 . Predict, giving a reason, the Cl – N – Cl bond angle in NCl3 .
(Total 3 marks)
55.
Arrange the following in decreasing order of bond angle (largest one first), and explain your
reasoning.
NH2 – , NH3 , NH4 +
(Total 5 marks)
56.
(i)
Outline the principles of the valence shell electron pair repulsion (VSEPR) theory.
(3)
(ii)
Use the VSEPR theory to deduce the shape of H 3 O+ and C2 H4 . For each species, draw the
Lewis structure, name the shape, and state the value of the bond angle(s).
(6)
(iii)
Predict and explain whether each species is polar.
(2)
(iv)
Using Table 7 of the Data Booklet, predict and explain which of the bonds O-H, O-N or
N-H would be most polar.
(2)
(Total 13 marks)
57.
Predict and explain which of the following compounds consist of molecules:
NaCl, BF3 , CaCl2 , N2 O, P 4 O6 , FeS and CBr4 .
(Total 2 marks)
15
58.
Diamond, graphite and C60 fullerene are three allotropes of carbon.
(i)
Describe the structure of each allotrope.
(3)
(ii)
Compare the bonding in diamond and graphite.
(2)
(Total 5 marks)
59.
State two physical properties associated with metals and explain them at the atomic level.
(Total 4 marks)
60.
(a)
Draw the Lewis structure of methanoic acid, HCOOH.
(1)
(b)
In methanoic acid, predict the bond angle around the
(2)
(c)
(i)
carbon atom. .....................................................................................................
(ii)
oxygen atom bonded to the hydrogen atom. ...................................................
State and explain the relationship between the length and strength of the bonds between
the carbon atom and the two oxygen atoms in methanoic acid.
......................................................................................................................................
......................................................................................................................................
......................................................................................................................................
......................................................................................................................................
......................................................................................................................................
......................................................................................................................................
(3)
(Total 6 marks)
16
IB Chemistry – SL
Topic 4 Answers
1.
D
[1]
2.
B
[1]
3.
B
[1]
4.
D
[1]
5.
C
[1]
6.
A
[1]
7.
B
[1]
8.
B
[1]
9.
B
10.
B
11.
B
12.
C
13.
A
14.
D
15.
B
16.
B
17.
D
[1]
[1]
[1]
[1]
[1]
[1]
[1]
[1]
[1]
18.
B
[1]
19.
B
[1]
20.
D
17
[1]
21.
C
[1]
22.
D
[1]
23.
B
[1]
24.
B
[1]
25.
D
[1]
26.
C
[1]
27.
D
[1]
28.
B
[1]
29.
A
[1]
30.
C
[1]
31.
A
[1]
32.
D
[1]
33.
B
[1]
34.
C
[1]
35.
A
[1]
36.
B
[1]
37.
D
[1]
38.
A
[1]
39.
B
[1]
40.
B
18
[1]
41.
(i)
(ii)
as molecules become larger/heavier/have higher M r values/
number of electrons increases; van der Waals’/London/
dispersion forces increase;
2
hydrogen bonding between molecules in H2 O; this bonding is stronger
(than van der Waals’ forces);
Must be an implied comparison with (i)
2
[4]
42.
(i)
(ii)
tetrahedral (accept correct 3-D diagram);
bent/V-shape/angular (accept suitable diagram);
2
105° (accept 103 – 106°);
lone pairs repel each other more than bonding pairs;
Do not accept repulsion of atoms.
2
[4]
43.
bonds are polar as Cl more electronegative than Si;
Allow “electronegativities are different”
molecule is symmetrical, hence polar effects cancel out/OWTTE;
2
[2]
44.
(a)
(b)
(i)
A – sodium iodide, B – sodium, C – iodine (three correct [1]);
Accept correct formulas.
(ii)
A – ionic bonding;
B – metallic bonding;
C – van der Waals’ forces (and covalent bonding);
(i)
(for Na)
(for NaI)
(ii)
1
3
(lattice of) positive ions/atoms;
delocalized/free electrons/sea of electrons;
oppositely charged ions/positive and negative ions;
free to move (only) in molten state;
4
forces between I2 molecules are weak;
ionic/metallic bonding strong(er);
2
[10]
45.
(i)
O
C
O ;
2
H S H;
Accept dots, crosses, a combination of dots and crosses or a
line to represent a pair of electrons.
(ii)
(iii)
CO2 is linear;
two charge centres or bonds and no lone pairs (around C);
H2 S is bent/v-shaped/angular;
two bond pairs, two lone pairs (around S);
4
CO2 is non-polar, H2 S is polar;
bond polarities cancel CO2 but not in H2 S;
2
[8]
46. CH3 Cl – dipole-dipole attractions;
CH4 – van der Waals’/dispersion/London forces;
19
CH3OH – hydrogen bond;
3
[3]
47.
(a)
(i)
(relative) measure of an atom’s attraction for electrons; in a bond;
2
(ii)
–
+
xx
hydrogen bonding
H
N
+
+H
H
xx
–
+
H
N
+
+H
(iii)
H
Suitable diagram indicating
dipoles;
lone pairs of electrons;
hydrogen bonding;
3
107°;
1
Accept answer in range 107 to 109° .
(b)
(iv)
molecule is asymmetrical/OWTTE;
(i)
109.5°;
(ii)
NH4 + has four bonding pairs
(around central atom so is a regular tetrahedron);
NH3 has three bonding pairs (of electrons) and one non-bonding pair;
non-bonding pairs (of electrons) exert a greater repulsive force;
Accept suitable diagrams.
1
1
3
[11]
48.
Si—Cl bonds are covalent;
3
Cl
Cl
Si
Cl
Cl
Accept lines for electron pairs.
Award [1] for covalent bonds and [1] for lone pairs.
[3]
49. find number of electron pairs/charge centres in (valence shell of) central atom;
electron pairs/charge centres (in valence shell) of central atom repel each other;
to positions of minimum energy/repulsion/maximum stability;
pairs forming a double or triple bond act as a single bond;
non-bonding pairs repel more than bonding pairs/OWTTE;
3 max
Do not accept repulsion between bonds or atoms.
Award [1] each for any three points.
[3]
20
50.
(i)
SCl2 two bonding pairs, two non-bonding pairs;
angular/bent/non-linear/V-shaped;
Both these marks can be scored from a diagram.
90° < angle < 107°;
C2 Cl2 two charge centres around each C;
linear;
Both these marks can be scored from a diagram.
angle = 180°;
(ii)
6
SCl2 is polar;
C2 Cl2 is non-polar;
No net dipole movement for C2 Cl2 but angular SCl2 has a
resultant dipole / OWTTE;
Mark can be scored from a diagram.
Allow ECF based on the answers given to (i).
3
[9]
51.
H
H
H
O
H
O
Allow a combination of dots, crosses or lines.
bent/V shaped/angular
104.5;
Accept answers in range 104 to 106.
repulsion of the two non-bonding pairs of electrons forces bond angle
to be smaller/non-bonding pairs repel more than bonding pairs;
4
[4]
52.
butane < propanone < propanol;
butane has van der Waals’ forces;
Accept vdW, dispersion or London forces or attractions
between temporary dipoles.
propanone has dipole-dipole attractions;
propanol has (the stronger) H-bonding;
4
[4]
53.
delocalized electrons;
(attracted) to positive ions;
more delocalized/mobile/outer shell electrons/higher ionic charge;
3
[3]
21
54.
Cl
N
Cl
Cl
All electrons must be shown.
Accept molecular structures using lines to represent bonding
and lone electron pairs.
bond angle: 107−109
greater repulsion between lone pair and bonding pairs/OWTTE;
NOT between electron pairs and atoms.
Award [1 max] if lone pair missed on nitrogen, ECF for bond
angle of 120.
3
[3]
55.
NH4 + > NH3 > NH2 – ;
NH4 + has four bonded electron pairs (and no lone electron pairs);
NH3 has three bonded electron pairs and one electron lone pair;
NH2 – has two bonded electron pairs and two electron lone pairs;
Accept correct Lewis structures with lone electron pairs clearly
shown.
lone pair-lone pair > lone pair-bonded pair > bonded pair-bonded pair/
lone pairs of electrons repel more than bonding pairs of electrons/OWTTE;
Do not accept repulsion between atoms.
5
[5]
56.
(i)
Find number of electron pairs/charge centres in (valence shell of)
central atom;
electron pairs/charge centres (in valence shell) of central atom repel
each other;
Any one of the following:
to positions of minimum energy/repulsion/maximum stability;
pairs forming a double or triple bond act as a single bond;
non-bonding pairs repel more than bonding pairs/OWTTE;
max
Do not accept repulsion between bonds or atoms.
3
(ii)
6
Species
H 3 O+
C2 H 4
Lewis (electron-dot)
structure
+
–
O
H ;
H H
H
H
H
H
;
Shape
Bond angle(s)
Trigonal/triangular
pyramidal;
Allow values in the
range 106° to 109.5°;
Trigonal/triangular
planar;
Allow values of
approximately 120°;
22
Accept crosses and dots for electrons in Lewis structures also.
As the Lewis structures were asked for, and not 3D
representations, do not penalize incorrectly drawn geometries.
Do not accept structure of hydronium cation without lone pair
on oxygen.
No penalty for missing charge.
(iii)
H3 O+: is polar and explanation either using a diagram or in words,
involving the net dipole moment;
e.g. the three individual O-H bond dipole moments add as
vectors to give a net dipole moment.
C2 H4 : is non-polar and explanation either using a diagram or in words,
involving no net dipole moment;
e.g. the vector sum of the individual bond dipole moments is
zero.
For simple answers such as bond polarities do not cancel for
H3 O+ and do cancel for C2H4, Award [1], only for the last two
marking points.
(iv)
O-H is most polar;
O-H has greatest difference between electronegativities/calculation
showing values of 1.4, 0.5 and 0.9 respectively;
2
2
[13]
57.
BF3 , N2 O, P 4 O6 and CBr4 ;
Non-metals only/small difference in electronegativity values of the elements;
2
[2]
58.
(i)
3
Allotrope
(ii)
Structure
Diamond
3D array/network involving tetrahedral
carbons/each carbon atom joined to four others;
Graphite
layer structure involving trigonal (triangular) planar
carbons/with each carbon atom joined to three
others/with hexagonal (six-membered) rings of
carbon atoms;
C60 fullerene
truncated icosahedrons;
Accept carbon atoms form a ‘ball’ with 32 faces, of
which 12 are pentagons and 20 are hexagons,
exactly like a soccer ball. Do not accept soccer ball
alone.
Diamond: covalent bonds (only);
Graphite: covalent bonds and the separated layers held together by
(weak) London/van der Waals’/dispersion forces;
2
[5]
59.
Electrical conductivity:
Bonding electrons are delocalised;
Current flow occurs without displacement of atoms within the metal/
able to flow within the metal;
23
Malleability:
Can be hammered into thin sheets;
atoms capable of slipping with respect to one another;
4
[4]
60.
(a)
O
H
C
O
H
1
No mark without lone electron pairs.
Correct shape not necessary.
Do not award mark if dots/crosses and bond lines are shown.
Accept lone pairs represented as straight lines.
(b)
(c)
O − C − O = 120°/H − C − O = 120°;
C − O − H = 109°/<109°;
No mark for 109.5°
Accept answer in range 100–109°
2
length: C = O < C − O;
strength: C = O > C – O;
greater number of electrons between nuclei pull atoms together and require greater
energy to break;
Or
double bonds are shorter/single bonds are longer;
double bonds are stronger/single bonds are weaker;
Accept stronger attraction between nuclei and (bonding)
electrons.
3
[6]
24
IB Chemistry HL
Topic4 Questions
1.
2.
What is the best description of the carbon-oxygen bond lengths in CO3 2– ?
A.
One short and two long bonds
B.
One long and two short bonds
C.
Three bonds of the same length
D.
Three bonds of different lengths
What is the number of sigma () and pi () bonds and the hybridization of the carbon atom in
O
H
3.
4.
C
O
H
Sigma
Pi
Hybridization
A.
4
1
sp2
B.
4
1
sp3
C.
3
2
sp3
D.
3
1
sp2
Which of the following contain a bond angle of 90°?
I.
PC14 +
II.
PCl5
III.
PCl6 –
A.
I and II only
B.
I and III only
C.
II and III only
D.
I, II and III
Which allotropes contain carbon atoms with sp2 hybridization?
I.
Diamond
II.
Graphite
III.
C60 fullerene
A.
I and II only
B.
I and III only
C.
II and III only
D.
I, II and III
1
5.
6.
7.
What is the molecular shape and the hybridization of the nitrogen atom in NH 3 ?
Molecular shape
Hybridization
A.
tetrahedral
sp3
B.
trigonal planar
sp2
C.
trigonal pyramidal
sp2
D.
trigonal pyramidal
sp3
Which statement about sigma and pi bonds is correct?
A.
Sigma bonds are formed only by s orbitals and pi bonds are formed only by p orbitals.
B.
Sigma bonds are formed only by p orbitals and pi bonds are formed only by s orbitals.
C.
Sigma bonds are formed by either s or p orbitals, pi bonds are formed only by p orbitals.
D.
Sigma and pi bonds are formed by either s or p orbitals.
Consider the following statements.
I.
All carbon-oxygen bond lengths are equal in CO 3 2– .
II.
All carbon-oxygen bond lengths are equal in CH 3 COOH.
III.
All carbon-oxygen bond lengths are equal in CH 3 COO– .
Which statements are correct?
8.
A.
I and II only
B.
I and III only
C.
II and III only
D.
I, II and III
Which statement is correct about multiple bonding between carbon atoms?
A.
Double bonds are formed by two π bonds.
B.
Double bonds are weaker than single bonds.
C.
π bonds are formed by overlap between s orbitals.
D.
π bonds are weaker than sigma bonds.
2
9.
10.
11.
12.
Which particles can act as ligands in complex ion formation?
I.
C1–
II.
NH3
III.
H2 O
A.
I and II only
B.
I and III only
C.
II and III only
D.
I, II and III
Which statements correctly describe the NO 2 – ion?
I.
It can be represented by resonance structures.
II.
It has two lone pairs of electrons on the N atom.
III.
The N atom is sp2 hybridized.
A.
I and II only
B.
I and III only
C.
II and III only
D.
I, II and III
Which is the smallest bond angle in the PF5 molecule?
A.
90
B.
109.5
C.
120
D.
180
Which types of hybridization are shown by the carbon atoms in the compound CH 2 = CH−CH3 ?
I.
sp
II.
sp2
III.
sp3
A.
I and II only
B.
I and III only
C.
II and III only
D.
I, II and III
3
13.
Identify the types of hybridization shown by the carbon atoms in the molecule
CH3 CH2 CH2COOH
14.
15.
I.
sp
II.
sp2
III.
sp3
A.
I and II only
B.
I and III only
C.
II and III only
D.
I, II and III
Which molecule is square planar in shape?
A.
XeO4
B.
XeF4
C.
SF4
D.
SiF4
What is the hybridization of nitrogen atoms I, II, III and IV in the following molecules?
H 2N N H 2
I
16.
II
HNN H
IV
III
I
II
III
IV
A.
sp2
sp2
sp3
sp3
B.
sp3
sp3
sp2
sp2
C.
sp2
sp2
sp
sp
D.
sp3
sp3
sp
sp
What is the molecular geometry and the Cl–I–Cl bond angle in the ICl4 – ion?
A.
Square planar 90
B.
Square pyramidal 90
C.
Tetrahedral 109
D.
Trigonal pyramidal 107
4
17.
18.
19.
20.
What is the geometry of the bonds around an atom with sp2 hybridization?
A.
2 bonds at 180
B.
3 bonds at 120
C.
2 bonds at 90, 1 bond at 180
D.
4 bonds at 109
How many sigma (σ) and pi (π) bonds are present in the structure of HCN?
σ
π
A.
1
3
B.
2
3
C.
2
2
D.
3
1
How many lone pairs and bonding pairs of electrons surround xenon in the XeF 4 molecule?
Lone pairs
Bonding pairs
A.
4
8
B.
0
8
C.
0
4
D.
2
4
(a)
Explain the meaning of the term hybridization.
....................................................................................................................................
....................................................................................................................................
(1)
(b)
State the type of hybridization shown by the carbon atom in the H–C≡N molecule, and
the number of  and  bonds present in the C≡N bond.
....................................................................................................................................
....................................................................................................................................
(2)
(c)
Describe how  and  bonds form.
....................................................................................................................................
....................................................................................................................................
....................................................................................................................................
....................................................................................................................................
(4)
(Total 7 marks)
5
21.
(i)
Draw the Lewis structures for carbon monoxide, carbon dioxide and the carbonate ion.
................................................................................................................................................
................................................................................................................................................
................................................................................................................................................
................................................................................................................................................
................................................................................................................................................
................................................................................................................................................
................................................................................................................................................
................................................................................................................................................
(3)
(ii)
Identify the species with the longest carbon-oxygen bond and explain your answer.
................................................................................................................................................
................................................................................................................................................
................................................................................................................................................
................................................................................................................................................
................................................................................................................................................
................................................................................................................................................
(3)
(Total 6 marks)
22.
In 1954 Linus Pauling was awarded the Chemistry Nobel Prize for his work on the nature of the
chemical bond. Covalent bonds are one example of intramolecular bonding.
Explain the formation of the following.
(i)
σ bonding
……………………………………………………………………………………………
……………………………………………………………………………………………
……………………………………………………………………………………………
(2)
(ii)
π bonding
……………………………………………………………………………………………
……………………………………………………………………………………………
……………………………………………………………………………………………
(2)
6
(iii)
double bonds
……………………………………………………………………………………………
……………………………………………………………………………………………
……………………………………………………………………………………………
(1)
(iv)
triple bonds
……………………………………………………………………………………………
……………………………………………………………………………………………
……………………………………………………………………………………………
(1)
(Total 6 marks)
23.
Atomic orbitals can mix by hybridization to form new orbitals for bonding.
Identify the type of hybridization present in each of the three following molecules.
Deduce and explain their shapes.
(i)
OF2
……………………………………………………………………………………………
……………………………………………………………………………………………
……………………………………………………………………………………………
……………………………………………………………………………………………
(3)
(ii)
H2 CO
……………………………………………………………………………………………
……………………………………………………………………………………………
……………………………………………………………………………………………
……………………………………………………………………………………………
(3)
(iii)
C2 H2
……………………………………………………………………………………………
……………………………………………………………………………………………
……………………………………………………………………………………………
……………………………………………………………………………………………
(3)
(Total 9 marks)
7
24.
For the following compounds
PCl3 , PCl5 , POCl3
(i)
Draw a Lewis structure for each molecule in the gas phase.
(Show all non-bonding electron pairs.)
.....................................................................................................................................
.....................................................................................................................................
.....................................................................................................................................
.....................................................................................................................................
.....................................................................................................................................
.....................................................................................................................................
.....................................................................................................................................
.....................................................................................................................................
.....................................................................................................................................
.....................................................................................................................................
(3)
(ii)
State the shape of each molecule and predict the bond angles.
.....................................................................................................................................
.....................................................................................................................................
.....................................................................................................................................
.....................................................................................................................................
.....................................................................................................................................
.....................................................................................................................................
(6)
(iii)
Deduce whether or not each molecule is polar, giving a reason for your answer.
.....................................................................................................................................
.....................................................................................................................................
.....................................................................................................................................
.....................................................................................................................................
.....................................................................................................................................
.....................................................................................................................................
(3)
(Total 12 marks)
8
25.
(i)
Explain the meaning of the term hybridization.
.....................................................................................................................................
.....................................................................................................................................
.....................................................................................................................................
(1)
(ii)
Discuss the bonding in the molecule CH 3 CHCH2 with reference to
•
the formation of σ and π bonds
•
the length and strength of the carbon-carbon bonds
•
the types of hybridization shown by the carbon atoms
.....................................................................................................................................
.....................................................................................................................................
.....................................................................................................................................
.....................................................................................................................................
.....................................................................................................................................
.....................................................................................................................................
.....................................................................................................................................
.....................................................................................................................................
(6)
(Total 7 marks)
26.
(a)
Draw the Lewis structures for the compounds XeF4 , PF5 and BF4 – .
(3)
(b)
Use the valance shell electron pair repulsion (VSEPR) theory to predict the shapes of the
three compounds in (a). State and explain the bond angles in each of the three
compounds.
(3)
(Total 6 marks)
27.
(a)
State the meaning of the term hybridization. State the type of hybridization shown by the
nitrogen atoms in N2 , N2 H2 and N2 H4 .
(4)
(b)
By referring to the N2 H2 molecule describe how sigma () and pi () bonds form and
describe how single and double bonds differ.
(4)
(Total 8 marks)
28.
(i)
Explain why the first ionization energy of magnesium is lower than that of fluorine.
(2)
(ii)
Write an equation to represent the third ionization energy of magnesium. Explain why the
third ionization energy of magnesium is higher than that of fluorine.
(3)
(Total 5 marks)
9
29.
Draw the Lewis structures, state the shapes and predict the bond angles for the following
species.
(i)
PCl5
(3)
(ii)
SCl2
(3)
(iii)
ICl4 –
(3)
(Total 9 marks)
30.
(a)
(i)
State the meaning of the term hybridization.
(1)
(ii)
State the type of hybridization around the carbon atoms in C 60 fullerene, diamond
and graphite.
(3)
(iii)
Explain why graphite and C60 fullerene can conduct electricity.
(2)
(b)
(i)
Compare how atomic orbitals overlap in the formation of sigma () and pi ()
bonds.
(2)
(ii)
State the number of sigma bonds and pi bonds in H 2 CC(CH3 )CHCH2.
(2)
(Total 10 marks)
31.
(i)
Apply the VSEPR theory to deduce the shape of NO 2− , ICl5 and SF4 . For each species,
draw the Lewis (electron dot) structure, name the shape, and state the value of the bond
angle(s).
(9)
(ii)
Discuss the bond angle(s) in SF4 .
(1)
(iii)
Explain the hybridization involved in the C2 H4 molecule.
(4)
(iv)
State the hybridization involved in the NO 2− ion and comment on the nitrogen-oxygen
bond distances.
(2)
(v)
Using Table 7 of the Data Booklet, predict and explain which of the bonds O-H, O-N or
N-H would be most polar.
(2)
(Total 18 marks)
10
IB Chemistry HL
Topic4 Answers
1.
C
2.
A
3.
C
4.
5.
C
D
6.
C
7.
B
8.
9.
D
D
10.
B
11.
A
12.
C
13.
C
14.
B
15.
B
16.
A
17.
B
18.
C
19.
D
20.
(a)
(b)
mixing/joining together/combining/merging of atomic orbitals to
form molecular orbitals/new orbitals/orbitals of equal energy;
Accept specific example such as mixing of s and p orbitals.
sp;
Do not award mark if sp 2 or sp 3 is also stated.
one sigma and two pi (bonds);
(c)
1
( bond formed by) end-on/axial overlap;
electrons/electron density between the two (carbon) atoms/OWTTE;
(π bond formed by) sideways/parallel overlap;
electrons/electron density above and below  bond/OWTTE;
Marks can be scored from a suitable diagram.
2
4
Do not award 2 nd and 4 th marks if electrons are not mentioned.
[7]
11
21.
(i)
2–
O
C
O
O
C
C
O
O
O
OTTWE
3
CO32–
Award [1] each. Need charge on
for [1].
Penalize missing lone electron pairs only once.
(ii)
CO3 2– ;
bond order 1 13 / 1 13 bonds each compared to double bonds in CO 2 and
triple bond in CO;
the fewer the number of bonding electrons, the less tightly nuclei
are held together, the longer the bond;
3
[6]
22.
(i)
(ii)
“head on” overlap of (2) orbitals;
along axial symmetry/along a line drawn through the
2 nuclei/OWTTE;
Accept suitable diagram for 2nd mark.
2
parallel p orbitals overlap sideways on;
above and below the line drawn through the 2 nuclei/OWTTE;
Accept suitable diagram for 2nd mark.
2
(iii)
1 σ and 1 π/σ and π;
1
(iv)
1 σ and 2 π/σ and π;
1
[6]
23.
(i)
(ii)
(iii)
OF2
sp3 ;
V-shaped/bent/angular;
2 bonding + 2 non-bonding (electron pairs);
H2 CO
sp2 ;
trigonal planar;
2 areas of electron density/negative charge centres;
C2 H2
sp;
linear;
2 areas of electron density/negative charge centres;
Accept suitable diagrams for shapes.
Allow [2] for ECF if correct explanation given for
incorrect formula, e.g. C2 H4.
3
3
3
[9]
12
24.
(i)
Award [1] for each correct Lewis structure.
PCl 3
Cl
P
Cl
Cl
PCl 5
Cl
Cl
Cl
P
Cl
Cl
POCl 3
O
Cl
P
Cl
Cl
3
Accept use of dots or crosses to represent electron pairs.
Subtract [1] if non-bonding pair on P in PCl3 is missing.
Subtract [1] if non-bonding pair(s) on Cl or O are missing.
Accept legitimate alternatives for POCl3, e.g. see below.
O
Cl
P
Cl
Cl
Cl
P
O
Cl
Cl
(ii)
PCl3
trigonal pyramid;
Accept answers in range
100° to 108°;
PCl5
trigonal bipyramid;
90° and 120°;
POCl3
tetrahedral;
Accept answers in range
100° to 112°;
Allow ECF if based on legitimate chemical structure.
6
13
(iii)
PCl3
polar, polarities do not
cancel/OWTTE;
PCl5
non-polar, polarities
cancel/OWTTE;
POCl3
polar, polarities do not
cancel/OWTTE;
3
Award [2] for three polarities correct, [1] for two polarities
correct, and [1] for correct reason(s).
Accept argument based on dipole moments.
Allow ECF if based on legitimate chemical structure.
[12]
25.
(i)
combining of atomic orbitals to form new orbitals/OWTTE;
(ii)
σ : overlap of orbitals between nuclei/end-on overlap;
 : overlap above and below line joining nuclei/sideways overlap;
Award [1] if candidate counts bonds (8 σ, 1 π), or
describes all three types of bonds
(i.e. C—H is σ, C—C is σ, C=C is σ and π).
1
single bonds longer than double;
double bonds stronger than single;
C of CH3 is sp3 ;
other two C are sp2 ;
Accept suitable diagrams.
6
[7]
26.
(a)
3
xx
x
xF
xx
xx
x
xF
xx
xx
xx
Xe
xx
; lone pairs on Xe required for the mark.
xx
x x
xFx
xx
x
xF
xx
xx
x
xF
xx
P
xx
F xx
xx
x x
x Fx
xx
x
x
xx
x
xF
xx
F xx
xx
xx
F xx
xx
xx
F xx
xx
B
F xx ; square brackets and charge required for the mark.
xx
x
x
x xFx x
14
Accept any combination of dots, crosses and lines.
Penalise missing fluorine lone pairs once only.
(b)
XeF4
Square planar and 90;
PF5
trigonal bipyramid and 90 and 120;
BF4 –
Tetrahedral and 109.5/109;
Allow clear suitable diagrams instead of name.
No ECF from (a).
3
[6]
27.
(a)
hybridization: mixing/merging of atomic orbitals;
N2 −sp;
N2 H2 −sp2 ;
N2 H4 −sp3 ;
(b)
4
 bonds (result from the) overlapping of orbitals end to end/along inter-nuclear
axis;
 bonds (result from the) overlapping of parallel/sideways p orbitals;
(single bonds)  bonds only;
(double bonds) have a σ bond and a  bond;
Suitable clear and labelled diagrams acceptable for all marks.
4
[8]
28.
(i)
electron removed from higher energy level/further from nucleus/greater
atomic radius;
increased repulsion by extra inner shell electrons/increased shielding
effect;
(ii)
2
Mg2+(g) → Mg3+(g) + e;
(even though) valence electrons in the same shell/main energy level/
Mg2+ has noble gas configuration;
Mg has greater nuclear/core charge/more protons;
3
[5]
29.
(i)
Cl
Cl
x
x
Cl
x
P
x
x
Cl
;
Cl
trigonal bipyramidal;
15
90;
120;
180;
3
Award [1] for 2 correct bond angles.
(ii)
xx
Cl x S
S xx
;
x
Cl
Bent/angular/V-shaped;
100−107;
3
(iii)
Cl
Cl
;
I
Cl
Cl
square planar;
90;
No ECF allowed.
Penalize once only [1] mark for missing lone pairs.
Accept structures using lines to represent bonding and lone
electron pairs.
3
[9]
30.
(a)
(i)
mixing/combining of atomic orbitals/OWTTE;
(ii)
C60 fullerene: sp2 ;
1
graphite : sp2 ;
(iii)
(b)
(i)
diamond: sp3 ;
3
each carbon atom is bound to 3 other carbon atoms/ bonding;
leading to delocalized electrons;
2
sigma/ bonds are formed by orbitals overlapping end to end/
along the internuclear axis/along line directly between nuclei;
Accept suitable diagram.
pi/ bonds are formed by p orbitals overlapping sideways;
Accept suitable diagram.
(ii)
12 sigma bonds;
2 pi bonds;
2
2
[10]
16
31.
(i)
Species
NO2 –
Lewis (electron-dot)
structure
N
;
O
O
Cl
Cl
ICl5
Bond angle(s)
Bent/Vshaped/angular;
109.5° < θ < 120°;
Square pyramidal;
|Inplane Cl-I-out-of-plane Cl| < 90°;
Allow corresponding correct
statement for other correctly identified
bond angles.
See-saw;
|Equatorial F-S-Equatorial F| < 120°;
Allow corresponding correct
statement for axial-equatorial and
axial-axial F-S-F angles.
Cl
;
I
Cl
SF4
Shape
Cl
F
F
S
F
F
;
9
Accept crosses and dots for electrons in the Lewis structures
also.
If all ideal bond angles are given, penalize once only.
As the Lewis structures were asked for, and not 3D
representations, do not penalize incorrectly drawn geometries.
(ii)
(iii)
(iv)
(v)
(equatorial F-S-equatorial F) less than 120° since non-bonding electron pairs
(exert greater repulsive forces and thus) compress the bond angles/OWTTE;
1
orbital diagram representation of carbon ground-state going to carbon
excited-state electron configuration;
mixing of orbitals to give three new entirely equivalent hybrid orbitals, sp 2 ,
on each carbon;
sp 2 orbitals trigonal (triangular) planar in shape;
unhybridized orbitals overlap to give π-bond;
4
sp 2 ;
both N-O bond lengths equal, (intermediate between double and single
bonds) due to resonance/delocalisation;
2
O-H is most polar;
O-H has greatest difference between electronegativities/calculation
showing values of 1.4, 0.5 and 0.9 respectively;
2
[18]
17
IB Chemistry SL
Topic 5 Questions
1.
2.
3.
What energy changes occur when chemical bonds are formed and broken?
A.
Energy is absorbed when bonds are formed and when they are broken.
B.
Energy is released when bonds are formed and when they are broken.
C.
Energy is absorbed when bonds are formed and released when they are broken.
D.
Energy is released when bonds are formed and absorbed when they are broken.
The temperature of a 2.0 g sample of aluminium increases from 25°C to 30°C.
How many joules of heat energy were added? (Specific heat of Al = 0.90 J g–1 K–1 )
A.
0.36
B.
2.3
C.
9.0
D.
11
Using the equations below:
C(s) + O2 (g) → CO2 (g)
Mn(s) + O2 (g) → MnO2 (s)
∆H = –390 kJ
∆H = –520 kJ
what is ∆H (in kJ) for the following reaction?
MnO2 (s) + C(s) → Mn(s) + CO2 (g)
4.
A.
910
B.
130
C.
–130
D.
–910
Which statements about exothermic reactions are correct?
I.
They have negative H values.
II.
The products have a lower enthalpy than the reactants.
III.
The products are more energetically stable than the reactants.
A.
I and II only
B.
I and III only
C.
II and III only
D.
I, II and III
1
5.
6.
A sample of a metal is heated. Which of the following are needed to calculate the heat absorbed
by the sample?
I.
The mass of the sample
II.
The density of the sample
III.
The specific heat capacity of the sample
A.
I and II only
B.
I and III only
C.
II and III only
D.
I, II and III
The average bond enthalpies for O—O and O==O are 146 and 496 kJ mol–1 respectively.
What is the enthalpy change, in kJ, for the reaction below?
H—O—O—H(g) → H—O—H(g) + ½O==O(g)
7.
A.
– 102
B.
+ 102
C.
+ 350
D.
+ 394
When the solids Ba(OH)2 and NH4 SCN are mixed, a solution is produced and the
temperature drops.
Ba(OH)2 (s) + 2NH4 SCN(s) → Ba(SCN)2 (aq) + 2NH3 (g) + 2H2 O(l)
Which statement about the energetics of this reaction is correct?
8.
A.
The reaction is endothermic and H is negative.
B.
The reaction is endothermic and H is positive.
C.
The reaction is exothermic and H is negative.
D.
The reaction is exothermic and H is positive.
Using the equations below
Cu(s) +
2Cu(s) +
1
2
O2 (g) → CuO(s)∆Hο = –156 kJ
1
2
O2 (g) → Cu2 O(s)∆Hο = –170 kJ
what is the value of ∆Hο (in kJ) for the following reaction?
2CuO(s) → Cu2 O(s) +
1
2
O2 (g)
A.
142
B.
15
C.
–15
D.
–142
2
9.
Which of the quantities in the enthalpy level diagram below is (are) affected by the use of a
catalyst?
I
Enthalpy
II
III
Time
10.
A.
I only
B.
III only
C.
I and II only
D.
II and III only
Consider the following equations.
Mg(s) +
H2 (g) +
1
2
1
2
O2 (g) → MgO(s)
O2 (g) → H2 O(g)
∆Hο = –602 kJ
∆Hο = –242 kJ
What is the ∆H° value (in kJ) for the following reaction?
MgO(s) + H2 (g) → Mg(s) + H2 O(g)
11.
12.
A.
–844
B.
–360
C.
+360
D.
+844
For which of the following is the sign of the enthalpy change different from the other three?
A.
CaCO3 (s) → CaO(s) + CO2 (g)
B.
Na(g) → Na+(g) + e–
C.
CO2 (s) → CO2 (g)
D.
2Cl(g) → Cl2 (g)
Separate solutions of HCl(aq) and H 2 SO4 (aq) of the same concentration and same volume
were completely neutralized by NaOH(aq). X kJ and Y kJ of heat were evolved respectively.
Which statement is correct?
A.
X=Y
B.
Y = 2X
C.
X = 2Y
D.
Y = 3X
3
13.
14.
15.
Which statements are correct for an endothermic reaction?
I.
The system absorbs heat.
II.
The enthalpy change is positive.
III.
The bond enthalpy total for the reactants is greater than for the products.
A.
I and II only
B.
I and III only
C.
II and III only
D.
I, II and III
The mass m (in g) of a substance of specific heat capacity c (in J g–1 K–1 ) increases by t°C.
What is the heat change in J?
A.
mct
B.
mc(t + 273)
C.
mct
1000
D.
mc (t + 273 )
1000
The average bond enthalpy for the C―H bond is 412 kJ mol–1 . Which process has an enthalpy
change closest to this value?
A.
CH4 (g) → C(s) + 2H2 (g)
B.
CH4 (g) → C(g) + 2H2 (g)
C.
CH4 (g) → C(s) + 4H(g)
D.
CH4 (g) → CH3 (g) + H(g)
(
16.
The following equation shows the formation of magnesium oxide from magnesium metal.
2Mg(s) + O2 (g)→2MgO(s)
HӨ = –1204kJ
Which statement is correct for this reaction?
A.
1204 kJ of energy are released for every mol of magnesium reacted.
B.
602 kJ of energy are absorbed for every mol of magnesium oxide formed.
C.
602 kJ of energy are released for every mol of oxygen gas reacted.
D.
1204 kJ of energy are released for every two mol of magnesium oxide formed.
4
17.
The following equations show the oxidation of carbon and carbon monoxide to carbon dioxide.
C(s) +O2 (g) → CO2 (g)
CO(g) +
1
2
O2 (g)→ CO2 (g)
HӨ = –x kJ mol–1
HӨ = –y kJ mol–l
What is the enthalpy change, in kJ mol–1 , for the oxidation of carbon to carbon monoxide?
C(s) +
18.
19.
1
2
O2 (g)→ CO(g)
A.
x+y
B.
–x–y
C.
y–x
D.
x–y
A simple calorimeter was used to determine the enthalpy of combustion of ethanol. The
experimental value obtained was –920 kJ mol–1 . The Data Booklet value is –1371 kJ mol–1 .
Which of the following best explains the difference between the two values?
A.
incomplete combustion of the fuel
B.
heat loss to the surroundings
C.
poor ventilation in the laboratory
D.
inaccurate temperature measurements
For the reaction
2H2 (g) + O2 (g) → 2H2 O(g)
the bond enthalpies (in kJ mol–1 ) are
H–H
x
O=O
y
O–H
z
Which calculation will give the value, in kJ mol–1 , of HӨ for the reaction?
20.
A.
2x + y –2z
B.
4z – 2x – y
C.
2x + y – 4z
D.
2z –2x – y
Which statement about bond enthalpies is correct?
A.
Bond enthalpies have positive values for strong bonds and negative values for weak
bonds.
B.
Bond enthalpy values are greater for ionic bonds than for covalent bonds.
C.
Bond breaking is endothermic and bond making is exothermic.
D.
The carbon–carbon bond enthalpy values are the same in ethane and ethene.
5
21.
An equation for a reaction in which hydrogen is formed is
CH4 + H2 O → 3H2 + CO
HӨ = +210 kJ
Which energy change occurs when 1 mol of hydrogen is formed in this reaction?
22.
A.
70 kJ of energy are absorbed from the surroundings.
B.
70 kJ of energy are released to the surroundings.
C.
210 kJ of energy are absorbed from the surroundings.
D.
210 kJ of energy are released to the surroundings.
The equations and enthalpy changes for two reactions used in the manufacture of sulfuric acid
are:
S(s) O2 (g) → SO2 (g)
HӨ = –300 kJ
2SO2 (g) + O2 (g) → 2SO3 (g)
HӨ = –200 kJ
What is the enthalpy change, in kJ, for the reaction below?
2S(s) + 3O2 (g) → 2SO3 (g)
23.
A.
–100
B.
–400
C.
–500
D.
–800
Approximate values of the average bond enthalpies, in kJ mol–1 , of three substances are:
H–H
430
F–F
155
H–F
565
What is the enthalpy change, in kJ, for this reaction?
2HF → H2 + F2
A.
+545
B.
+20
C.
–20
D.
–545
6
24.
The standard enthalpy change of formation values of two oxides of phosphorus are:
P 4 (s) + 3O2 (g) → P 4 O6 (s)
HӨf= –1600 kJ mol–1
P 4 (s) + 5O2 (g) → P 4 O10 (s)
HӨf= –3000 kJ mol–1
What is the enthalpy change, in kJ mol–1 , for the reaction below?
P 4 O6 (s) + 2O2 (g) → P 4 O10(s)
25.
26.
A.
+4600
B.
+1400
C.
–1400
D.
–4600
Which statement is correct for an endothermic reaction?
A.
The products are more stable than the reactants and H is positive.
B.
The products are less stable than the reactants and H is negative.
C.
The reactants are more stable than the products and H is positive.
D.
The reactants are less stable than the products and H is negative.
Which statement is correct about the reaction shown?
2SO2 (g) + O2 (g) → 2SO3 (g)
27.
H = –196 kJ
A.
196 kJ of energy are released for every mole of SO 2 (g) reacted.
B.
196 kJ of energy are absorbed for every mole of SO 2 (g) reacted.
C.
98 kJ of energy are released for every mole of SO 2 (g) reacted.
D.
98 kJ of energy are absorbed for every mole of SO 2 (g) reacted.
Which statements are correct for all exothermic reactions?
I.
The enthalpy of the products is less than the enthalpy of the reactants.
II.
The sign of H is negative.
III.
The reaction is rapid at room temperature.
A.
I and II only
B.
I and III only
C.
II and III only
D.
I, II and III
7
28.
Consider the specific heat capacity of the following metals.
Metal
Specific heat capacity / J kg–1 K–1
Cu
385
Ag
234
Au
130
Pt
134
Which metal will show the greatest temperature increase if 50 J of heat is supplied to a 0.001 kg
sample of each metal at the same initial temperature?
29.
A.
Cu
B.
Ag
C.
Au
D.
Pt
Consider the following reactions.
S(s) + 1 12 O2 (g) → SO3 (g)
SO2 (s) +
1
2
O2 (g) → SO3 (g)
HӨ = −395 kJ mol−1
HӨ = −98 kJ mol−1
What is the HӨ value (in kJ mol–1 ) for the following reaction?
S(s) + O2 (g) → SO2 (g)
30.
A.
–297
B.
+297
C.
– 493
D.
+493
Which statement is correct for an endothermic reaction?
A.
Bonds in the products are stronger than the bonds in the reactants.
B.
Bonds in the reactants are stronger than the bonds in the products.
C.
The enthalpy of the products is less than that of the reactants.
D.
The reaction is spontaneous at low temperatures but becomes non-spontaneous at high
temperatures.
8
31.
According to the enthalpy level diagram below, what is the sign for H and what term is used to
refer to the reaction?
H
reactants
products
reaction progress
32.
H
reaction
A.
positive
endothermic
B.
negative
exothermic
C.
positive
exothermic
D.
negative
endothermic
When 40 joules of heat are added to a sample of solid H 2 O at –16.0°C the temperature increases
to –8.0°C. What is the mass of the solid H 2 O sample?
[Specific heat capacity of H 2 O(s) = 2.0 J g–1 K–1 ]
33.
A.
2.5 g
B.
5.0 g
C.
10 g
D.
160 g
The HӨ values for the formation of two oxides of nitrogen are given below.
1
2
HӨ = –57 kJ mol–1
N2 (g) + O2 (g) → NO2 (g)
N2 (g) + 2O2 (g) → N2 O4 (g)
HӨ = +9 kJ mol–1
Use these values to calculate HӨ for the following reaction (in kJ):
2NO2 (g) → N2 O4 (g)
A.
–105
B.
– 48
C.
+66
D.
+123
9
34.
35.
36.
How much energy, in joules, is required to increase the temperature of 2.0 g of aluminium from
25 to 30°C? (Specific heat of Al = 0.90 J g–1 K–1 ).
A.
0.36
B.
4.5
C.
9.0
D.
54
Which combination is correct for a chemical reaction that absorbs heat from the surroundings?
Type of reaction
ΔH at constant pressure
A.
Exothermic
Positive
B.
Exothermic
Negative
C.
Endothermic
Positive
D.
Endothermic
Negative
Using the equations below:
C(s) + O2 (g) → CO2 (g)
∆Hο = –394 kJ mol–1
Mn(s) + O2 (g) → MnO2 (s)
∆Hο = –520 kJ mol–1
What is ∆H, in kJ, for the following reaction?
MnO2 (s) + C(s) → Mn(s) + CO2 (g)
A.
914
B.
126
C.
–126
D.
–914
10
37.
Number of
molecules
Ea
Energy
The diagram shows the distribution of energy for the molecules in a sample of gas at a given
temperature, T1 .
(a)
In the diagram Ea represents the activation energy for a reaction. Define this term.
……………………………………………………………………………………….
……………………………………………………………………………………….
(1)
(b)
On the diagram above draw another curve to show the energy distribution for the same
gas at a higher temperature. Label the curve T2 .
(2)
(c)
With reference to your diagram, state and explain what happens to the rate of a reaction
when the temperature is increased.
……………………………………………………………………………………….
……………………………………………………………………………………….
(2)
(Total 5 marks)
38.
(a)
Define the term average bond enthalpy, illustrating your answer with an equation for
methane, CH4 .
.....................................................................................................................................
(3)
(b)
The equation for the reaction between methane and chlorine is
CH4 (g) + Cl2 (g) → CH3 Cl(g) + HCl(g)
Use the values from Table 10 of the Data Booklet to calculate the enthalpy change for
this reaction.
.....................................................................................................................................
.....................................................................................................................................
.....................................................................................................................................
(3)
11
(c)
Explain why no reaction takes place between methane and chlorine at room
temperature unless the reactants are sparked, exposed to UV light or heated.
.....................................................................................................................................
.....................................................................................................................................
.....................................................................................................................................
.....................................................................................................................................
(2)
(d)
Draw an enthalpy level diagram for this reaction.
(2)
(Total 10 marks)
39.
In aqueous solution, potassium hydroxide and hydrochloric acid react as follows.
KOH(aq) + HCl(aq) → KCl(aq)+ H2 O(l)
The data below is from an experiment to determine the enthalpy change of this reaction.
50.0 cm3 of a 0.500 mol dm–3 solution of KOH was mixed rapidly in a glass beaker with
50.0 cm3 of a 0.500 mol dm–3 solution of HCl.
Initial temperature of each solution = 19.6°C
Final temperature of the mixture = 23.1°C
(a)
State, with a reason, whether the reaction is exothermic or endothermic.
.....................................................................................................................................
(1)
(b)
Explain why the solutions were mixed rapidly.
.....................................................................................................................................
(1)
(c)
Calculate the enthalpy change of this reaction in kJ mol–1 . Assume that the specific heat
capacity of the solution is the same as that of water.
.....................................................................................................................................
.....................................................................................................................................
.....................................................................................................................................
(4)
12
(d)
Identify the major source of error in the experimental procedure described above.
Explain how it could be minimized.
.....................................................................................................................................
.....................................................................................................................................
(2)
(e)
The experiment was repeated but with an HCl concentration of 0.510 mol dm–3 instead
of 0.500 mol dm–3 . State and explain what the temperature change would be.
.....................................................................................................................................
.....................................................................................................................................
(2)
(Total 10 marks)
40.
The data below is from an experiment used to measure the enthalpy change for the combustion
of 1 mole of sucrose (common table sugar), C 12 H22O11 (s). The time-temperature data was taken
from a data-logging software programme.
Mass of sample of sucrose, m = 0.4385 g
Heat capacity of the system, Csystem = 10.114 kJ K–1
(a)
Calculate ΔT, for the water, surrounding the chamber in the calorimeter.
.....................................................................................................................................
(1)
13
(b)
Determine the amount, in moles, of sucrose.
.....................................................................................................................................
.....................................................................................................................................
.....................................................................................................................................
.....................................................................................................................................
(1)
(c)
(i)
Calculate the enthalpy change for the combustion of 1 mole of sucrose.
...........................................................................................................................
...........................................................................................................................
(1)
(ii)
Using Table 12 of the Data Booklet, calculate the percentage experimental error
based on the data used in this experiment.
...........................................................................................................................
...........................................................................................................................
(1)
(d)
A hypothesis is suggested that TNT, 2-methyl-1,3,5-trinitrobenzene, is a powerful
explosive because it has:
• a large enthalpy of combustion
• a high reaction rate
• a large volume of gas generated upon combustion
Use your answer in part (c)(i) and the following data to evaluate this hypothesis:
Equation for combustion
Relative
rate of
combustion
Sucrose
C12 H22 O11(s) + 12O2 (g) → 12CO2 (g) + 11H2 O(g)
Low
TNT
2C7 H5 N3 O6(s) → 7CO(g) + 7C(s) + 5H2 O(g) + 3N2 (g)
High
Enthalpy of
combustion
/ kJ mol–1
3406
.....................................................................................................................................
.....................................................................................................................................
.....................................................................................................................................
.....................................................................................................................................
.....................................................................................................................................
.....................................................................................................................................
(3)
(Total 7 marks)
14
41.
(a)
Define the term average bond enthalpy.
.....................................................................................................................................
.....................................................................................................................................
.....................................................................................................................................
.....................................................................................................................................
(2)
(b)
Use the information from Table 10 of the Data Booklet to calculate the enthalpy change
for the complete combustion of but-1-ene, according to the following equation.
C4 H8 (g) + 6O2 (g) → 4CO2 (g) + 4H2 O(g)
.....................................................................................................................................
.....................................................................................................................................
.....................................................................................................................................
.....................................................................................................................................
(3)
(Total 5 marks)
42.
Given the following data:
C(s) + F2 (g) → CF4 (g); ∆H1 = –680 kJ mol–1
F2 (g) → 2F(g); ∆H2 = +158 kJ mol–1
C(s) → C(g); ∆H3 = +715 kJ mol–1
calculate the average bond enthalpy (in kJ mol–1 ) for the C––F bond.
.....................................................................................................................................
.....................................................................................................................................
.....................................................................................................................................
.....................................................................................................................................
.....................................................................................................................................
.....................................................................................................................................
.....................................................................................................................................
.....................................................................................................................................
.....................................................................................................................................
.....................................................................................................................................
(Total 4 marks)
15
43.
Two reactions occurring in the manufacture of sulfuric acid are shown below:
reaction I
reaction II SO2 (g) +
(i)
HӨ = –297 kJ
S(s) +O2 (g) → SO2 (g)
1
2
O2 (g)
SO3 (g)
HӨ = –92 kJ
State the name of the term HӨ. State, with a reason, whether reaction I would be
accompanied by a decrease or increase in temperature.
(3)
(ii)
At room temperature sulfur trioxide, SO 3 , is a solid. Deduce, with a reason, whether the
HӨ value would be more negative or less negative if SO 3 (s) instead of SO3 (g) were
formed in reaction II.
(2)
(iii)
Deduce the HӨ value of this reaction:
S(s) + 1 12 O2 (g) → SO3 (g)
(1)
(Total 6 marks)
44.
(i)
Define the term average bond enthalpy.
(3)
(ii)
Explain why Br2 is not suitable as an example to illustrate the term average bond
enthalpy.
(1)
(iii)
Using values from Table 10 of the Data Booklet, calculate the enthalpy change for the
following reaction:
CH4 (g) + Br2 (g) → CH3 Br(g) + HBr(g)
(3)
(iv)
Sketch an enthalpy level diagram for the reaction in part (iii).
(2)
(v)
Without carrying out a calculation, suggest, with a reason, how the enthalpy change for
the following reaction compares with that of the reaction in part (iii):
CH3 Br(g) + Br2 (g) → CH2 Br2 (g) + HBr(g)
(2)
(Total 11 marks)
16
45.
But–1–ene gas, burns in oxygen to produce carbon dioxide and water vapour according to the
following equation.
C4 H8 + 6O2 → 4CO2 + 4H2 O
(a)
Use the data below to calculate the value of HӨ for the combustion of but-1-ene.
Bond
C−C
C=C
C−H
O=O
C=O
O–H
348
612
412
496
743
463
Average bond
enthalpy / kJ
mol–1
....................................................................................................................................
....................................................................................................................................
....................................................................................................................................
....................................................................................................................................
....................................................................................................................................
....................................................................................................................................
(3)
(b)
State and explain whether the reaction above is endothermic or exothermic.
....................................................................................................................................
....................................................................................................................................
(1)
(Total 4 marks)
46.
Calculate the enthalpy change, H4 for the reaction
C + 2H2 +
1
2
O2 → CH3 OH
H4
using Hess’s Law and the following information.
CH3 OH + 1 12 O2 → CO2 + 2H2 O
H1 = −676 kJ mol−1
C + O2 → CO2
H2 = −394 kJ mol−1
H2 +
1
2
O2 → H2 O
H3 = −242 kJ mol−1
..............................................................................................................................................
..............................................................................................................................................
..............................................................................................................................................
..............................................................................................................................................
..............................................................................................................................................
(Total 4 marks)
17
47.
Methylamine can be manufactured by the following reaction.
CH3 OH(g) + NH3 (g) → CH3 NH2(g) + H2 O(g)
(a)
Define the term average bond enthalpy.
....................................................................................................................................
....................................................................................................................................
....................................................................................................................................
....................................................................................................................................
(2)
(b)
Use information from Table 10 of the Data Booklet to calculate the enthalpy change for
this reaction.
....................................................................................................................................
....................................................................................................................................
....................................................................................................................................
....................................................................................................................................
....................................................................................................................................
....................................................................................................................................
(4)
(Total 6 marks)
48.
(a)
Define the term average bond enthalpy.
....................................................................................................................................
....................................................................................................................................
....................................................................................................................................
....................................................................................................................................
(2)
(b)
Use the information from Table 10 in the Data Booklet to calculate the enthalpy change
for the complete combustion of but-1-ene according to the following equation
C4 H8 (g) → 4CO2 (g) + 4H2 O(g)
....................................................................................................................................
....................................................................................................................................
....................................................................................................................................
....................................................................................................................................
....................................................................................................................................
(3)
18
(c)
Predict, giving a reason, how the enthalpy change for the complete combustion of
but-2-ene would compare with that of but-1-ene based on average bond enthalpies.
....................................................................................................................................
....................................................................................................................................
(1)
(d)
The enthalpy level diagram for a certain reaction is shown below.
Enthalpy
HR
enthalpy of
reactants
HP
enthalpy of
products
Time
State and explain the relative stabilities of the reactants and products.
....................................................................................................................................
....................................................................................................................................
(2)
(Total 8 marks)
49.
The reaction between ethene and hydrogen gas is exothermic.
(i)
Write an equation for this reaction.
(1)
(ii)
Deduce the relative stabilities and energies of the reactants and products.
(2)
(iii)
Explain, by referring to the bonds in the molecules, why the reaction is exothermic.
(2)
(Total 5 marks)
50.
(i)
Define the term average bond enthalpy.
(2)
(ii)
The equation for the reaction of ethyne and hydrogen is:
C2 H2 (g) + 2H2 (g) → C2 H6 (g)
Use information from Table 10 of the Data Booklet to calculate the change in enthalpy
for the reaction.
(2)
(iii)
State and explain the trend in the bond enthalpies of the C–Cl, C–Br and C–I bonds.
(2)
(Total 6 marks)
19
IB Chemistry SL
Topic 5 Answers
1.
D
2.
C
3.
B
4.
D
5.
B
6.
7.
A
B
8.
A
9.
C
10.
C
11.
D
12.
13.
B
A
14.
A
15.
D
16.
D
17.
C
18.
19.
B
C
20.
C
21.
A
22.
D
23.
A
24.
C
25.
C
26.
C
27.
A
28.
C
29.
A
20
30.
B
31.
B
32.
A
33.
D
34.
C
35.
C
36.
B
37.
(a)
activation energy = minimum energy required for a reaction to occur;
(b)
curve moved to the right;
peak lower,
Deduct [1] if shaded area smaller at T2 or if T2 line touches the
x-axis
(c)
1
2
rate increased;
as more molecules with energy  Ea;
2
[5]
38.
(a)
(b)
(c)
energy for the conversion of a gaseous molecule into (gaseous) atoms;
(average values) obtained from a number of similar bonds/compounds/OWTTE;
CH4 (g) → C(g) + 4H(g);
State symbols needed.
(bond breaking) = 1890/654;
(bond formation) = 2005/769;
enthalpy = –115(kJ mol–1 )
Award [3] for correct final answer.
Penalize [1] for correct answer with wrong sign.
molecules have insufficient energy to react (at room temperature)/
wrong collision geometry/unsuccessful collisions;
extra energy needed to overcome the activation energy/Ea for the reaction;
3
3
2
(d)
Ea
energy
reactants
products
reaction path
exothermic shown;
activation energy/Ea shown;
2
[10]
21
39.
(a)
exothermic because temperature rises/heat is released;
(b)
to make any heat loss as small as possible/so that all the heat will be
given out very quickly;
Do not accept “to produce a faster reaction”.
1
heat released = mass×specific heat capacity×temp increase/q = mc∆T =/
100×4.18×3.5;
= 1463 J/1.463 kJ; (allow 1.47 kJ if specific heat = 4.2)
amount of KOH/HCl used = 0.500×0.050 = 0.025 mol;
∆H = (1.463÷0.025) = –58.5 (kJ mol–1 ); (minus sign needed for mark)
Use ECF for values of q and amount used.
Award [4] for correct final answer.
Final answer of 58.5 or +58.5 scores [3].
Accept 2,3 or 4 significant figures.
4
(c)
(d)
(e)
1
heat loss (to the surroundings);
insulate the reaction vessel/use a lid/draw a temperature versus time graph;
2
3.5°C/temperature change would be the same;
amount of base reacted would be the same/excess acid would not react/
KOH is the limiting reagent;
2
[10]
40.
(a)
ΔT = 23.70 – 23.03 = 0.67 (°C/K);
1
(b)
 0.4385 g 
 = 1.281×10–3 ;
n = 
−1 
 342.34 g mol 
1
(c)
(i)
ΔHc = (C ΔT)/n =
− [(10 .114 kJ K −1 )(0.67 K)]
(1.281  10
−3
mol)
= –5.3×103 kJ mol–1 ;
1
Use ECF for values of T and n.
(ii)
(d)
 (−5.3 10 3 ) + (5.6 10 3 ) 
Percentage experimental error = 
 100 = 5.4%;
(−5.6 10 3 )


Use ECF for values of ΔHc.
enthalpy change of combustion of sucrose > TNT, and therefore not important;
rate of reaction for TNT is greater than that of sucrose, so this is valid;
amount of gas generated (in mol) for sucrose > than that of TNT
(according to the given equation), so this is not important;
1
3
[7]
41.
(a)
(b)
The amount of energy needed to break 1 mole of (covalent) bonds;
in the gaseous state;
average calculated from a range of compounds;
max
Award [1] each for any two points above.
Bonds broken
(612) + (2×348) + (8×412) + (6×496)/7580 (kJ mol–1 );
Bonds made
(8×743) + (8×463) / 9648 (kJ mol–1 );
H = –2068 (kJ mol–1 );
2
3
22
Award [3] for the correct answer.
Allow full ECF.
Allow kJ but no other incorrect units.
Even if the first two marks are lost, the candidate can score [1]
for a clear correct subtraction for H.
[5]
42.
C(s) + 2F2 (g) → CF4 (g)
∆H1 = –680 kJ;
4F(g) → 2F2 (g)
∆H2 = 2(–158) kJ;
C(g) → C(s)
∆H3 = –715 kJ;
Accept reverse equations with +∆H values.
C(g) + 4F(g) → CF4 (g)
∆H = –1711 kJ,
so average bond enthalpy = – 1711
4
= –428 kJ mol–1 ;
4
Accept + or – sign.
Lots of ways to do this! The correct answer is very different
from the value in the Data Booklet, so award [4] for final
answer with/without sign units not needed, but deduct [1] if
incorrect units. Accept answer in range of 427 to 428 without
penalty for sig figs.
If final answer is not correct use following;
Award [1] for evidence of cycle or enthalpy diagram or adding
of equations.
Award [1] for 2F2 (g) → 4F(g) 2×158 seen.
Award [1] for dividing 1711 or other value by 4.
[4]
43.
(a)
(i)
(ii)
(iii)
standard enthalpy (change) of reaction;
(temperature) increase;
reaction is exothermic/sign of H is negative;
3
more (negative);
heat given out when gas changes to solid/solid has less enthalpy than
gas/OWTTE;
2
–389 kJ;
1
[6]
44.
(i)
the energy needed to break one bond;
(in a molecule in the) gaseous state;
value averaged using those from similar compounds;
3
(ii)
it is an element/no other species with just a Br-Br bond/OWTTE;
1
(iii)
(sum bonds broken =) 412 + 193 = 605;
(sum bonds formed =) 276 + 366 = 642;
(H =) –37 kJ;
Award [3] for correct final answer.
Award [2] for “+ 37”.
Accept answer based on breaking and making extra C-H bonds.
3
23
(iv)
CH4 + Br2
Enthalpy
;
CH3 Br + HBr
2
Award [1] for enthalpy label and two horizontal lines, [1] for
reactants higher than products.
ECF from sign in (iii), ignore any higher energy level involving
atoms.
(v)
(about) the same/similar;
same (number and type of) bonds being broken and formed;
2
[11]
45.
(a)
(Amount of energy required to break bonds of reactants)
8×412 + 2×348 + 612 + 6×496/7580 (kJ mol−1 );
(Amount of energy released during bond formation)
4×2×743 + 4×2×463/9648 (kJ mol−1 );
H = −2068 (kJ or kJ mol−1 );
3
ECF from above answers.
Correct answer scores [3].
Award [2] for (+)2068.
If any other units apply − 1(U), but only once per paper.
(b)
exothermic and HӨ is negative/energy is released;
Apply ECF to sign of answer in part (a).
Do not mark if no answer to (a).
1
[4]
46.
−1×H1 /676;
1×H2 /–394;
2×H3 /– 484;
H4 = −202 (kJ mol−1 );
Accept alternative methods.
Correct answers score [4].
4
Award [3] for (+)202 or (+)40 (kJ/kJ mol−1 ).
−1(U) if units incorrect (ignore if absent).
[4]
47.
(a)
(b)
energy needed to break (1 mol of) a bond in a gaseous molecule;
averaged over similar compounds;
2
bonds broken identified as C−O and N−H;
bonds formed identified as C−N and O−H;
H = 748 − 768 (kJ);
= − 20 kJ/kJ mol−1 (units needed for this mark);
If wrong bonds identified apply ECF to 3rd and 4th marks.
4
24
Accept answer based on breaking and making all bonds.
Award [4] for correct final answer.
Award max [3] if only one bond missed.
Answer of 20 or +20 kJ (mol−1 ) scores [3].
[6]
48.
(a)
(b)
amount of energy needed to break one mole of (covalent) bonds;
in the gaseous state;
average calculated from a range of compounds;
Award [1] each for any two points above.
2
bonds broken: 161 + 2×348 + 8×412 + 6×496/7580 kJ mol−1 ;
bonds made: 8×743 + 8×463/9648 kJ mol−1 ;
(bonds broken − bonds made =) H = −2068(kJ mol−1 );
Award [3] for the correct answer.
Allow full ECF − 1 mistake equals 1 penalty.
Allow kJ but not other wrong units.
3
(c)
same/equal, because the same bonds are being broken and formed;
1
(d)
products more stable than reactants;
bonds are stronger in products than reactants/HP < HR/enthalpy/stored
energy of products less than reactants;
2
[8]
49.
(a)
(i)
C2 H4 (g) + H2 (g) → C2H6 (g);
State symbols not required for mark
(ii)
products more stable than reactants/reactants less stable than products;
products lower in energy/reactants higher in energy;
(iii)
1
2
(overall) bonds in reactants weaker/(overall) bonds in product stronger
/all bonds in product are  bonds/weaker  bond broken and a
(stronger)  bond formed;
less energy needed to break weaker bonds/more energy produced
to make stronger bonds (thus reaction is exothermic)/OWTTE;
OR
bond breaking is endothermic/requires energy and bond making is
exothermic/releases energy;
stronger bonds in product mean process is exothermic overall;
2
[5]
50.
(i)
energy required to break (a mole of) bonds in the gaseous state
/energy given out when (a mole of) bonds are made in the
gaseous state;
average value from a number of similar compounds;
2
25
(ii)
(HӨreaction = (∑BEbreak − BEmake))
= [(837) + 2(436)] − [(348 + 4(412)];
= − 287(kJ/kJ mol−1 );
Award [1 max] for 287 or + 287.
(iii)
(BE): C−Cl > C−Br > C−I/C−X bond becomes weaker;
halogen size/radius increases/bonding electrons further away from
the nucleus/bonds become longer;
2
2
[6]
26
IB Chemistry HL
Topic5 Questions
1.
2.
Which combination of ionic charge and ionic radius give the largest lattice enthalpy for an ionic
compound?
Ionic charge
Ionic radius
A.
high
large
B.
high
small
C.
low
small
D.
low
large
The lattice enthalpy values for lithium fluoride and calcium fluoride are shown below.
LiF(s)
CaF2 (s)
∆Hο = +1022 kJ mol–1
∆Hο = +2602 kJ mol–1
Which of the following statements help(s) to explain why the value for lithium fluoride is less
than that for calcium fluoride?
3.
4.
I.
The ionic radius of lithium is less than that of calcium.
II.
The ionic charge of lithium is less than that of calcium.
A.
I only
B.
II only
C.
I and II
D.
Neither I nor II
Which reaction has the most negative ∆Hο value?
A.
LiF(s) → Li+(g) + F– (g)
B.
Li+(g) + F– (g) → LiF(s)
C.
NaCl(s) → Na+(g) + Cl– (g)
D.
Na+(g) + Cl– (g) → NaCl(s)
Which type of reaction is referred to in the definition of standard enthalpy change of formation?
A.
the formation of a compound from its elements
B.
the formation of a crystal from its ions
C.
the formation of a molecule from its atoms
D.
the formation of a compound from other compounds
1
5.
6.
7.
8.
9.
What is the correct order of decreasing entropy for a pure substance?
A.
gas  liquid  solid
B.
solid  liquid  gas
C.
solid  gas  liquid
D.
liquid  solid  gas
Which reaction has the largest positive value of S Ө?
A.
CO2 (g) + 3H2 (g) → CH3 OH(g) + H2 O(g)
B.
2Al(s) + 3S(s) → Al2 S3 (s)
C.
CH4 (g) + H2 O(g) → 3H2 (g) + CO(g)
D.
2S(s) + 3O2 (g) → 2SO3 (g)
Which is a correct equation to represent the lattice enthalpy of magnesium sulfide?
A.
MgS(s) → Mg(s) + S(s)
B.
MgS(s) → Mg(g) + S(g)
C.
MgS(s) → Mg+(g) + S– (g)
D.
MgS(s) → Mg2+(g) + S2– (g)
Which equation represents a change with a negative value for S?
A.
2H2 (g) + O2 (g) → 2H2 O(g)
B.
H2 O(s) → H2 O(g)
C.
H2 (g) + Cl2 (g) → 2HCl(g)
D.
2NH3 (g) → N2 (g) + 3H2 (g)
The expression for the standard free energy change of a reaction is given by
GӨ = HӨ – TS Ө
What are the signs for HӨ and S Ө for a reaction that is spontaneous at all temperatures?
HӨ
S Ө
A.
+
–
B.
–
+
C.
+
+
D.
–
–
2
10.
11.
Which are characteristics of ions in an ionic compound with a large lattice enthalpy value?
A.
Large ionic radius and high ionic charge
B.
Small ionic radius and low ionic charge
C.
Large ionic radius and low ionic charge
D.
Small ionic radius and high ionic charge
The following reaction is spontaneous only at temperatures above 850C.
CaCO3 (s) → CaO(s) + CO2 (g)
Which combination is correct for this reaction at 1000C?
12.
G
H
S
A.
–
–
–
B.
+
+
+
C.
–
+
+
D.
+
–
–
Consider the following information.
Compound
C6 H6 (l)
CO2 (g)
H2 O(l)
HfӨ / kJ mol–1
+49
+394
–286
C6 H6 (l) + 7 12 O2 (g) → 6CO2 (g) + H2 O(l)
Which expression gives the correct value of the standard enthalpy change of combustion for
benzene (l), in kJ mol–1 ?
13.
A.
12(−394) + (−286) −2(49)
B.
12(394) + 6(286) −2(−49)
C.
6(−394) + 3(−286) − (−49)
D.
6(394) + 3(286) − (−49)
Which equation represents the lattice enthalpy of magnesium oxide?
1
2
O2 (g) → MgO(s)
A.
Mg(s) +
B.
Mg2+(g) + O2– (g) → MgO(g)
C.
Mg2+(g) +
D.
Mg2+(g) + O2−(g) → MgO(s)
1
2
O2 (g) → MgO(s)
3
14.
15.
16.
17.
18.
The HӨ and S Ө values for a reaction are both negative. What will happen to the spontaneity
of this reaction as the temperature is increased?
A.
The reaction will become more spontaneous as the temperature is increased.
B.
The reaction will become less spontaneous as the temperature is increased.
C.
The reaction will remain spontaneous at all temperatures.
D.
The reaction will remain non-spontaneous at any temperature.
Which reaction has the most negative ∆Hο value?
A.
LiF(s) → Li+(g) + F– (g)
B.
Li+(g) + F– (g) → LiF(s)
C.
NaCl(s) → Na+(g) + Cl– (g)
D.
Na+(g) + Cl– (g) → NaCl(s)
Which equation represents the electron affinity of calcium?
A.
Ca(g) →Ca+(g) + e–
B.
Ca(g) →Ca– (g) + e–
C.
Ca(g) + e– → Ca– (g)
D.
Ca+(g) + e– → Ca(g)
Which reaction causes a decrease in the entropy of the system?
A.
CaCO3 (s) → CaO(s) + CO2 (g)
B.
2H2 (g) + O2 (g) → 2H2 O(l)
C.
2C(s) + O2 (g) → 2CO(g)
D.
2SO3 (g) → 2SO2 (g) + O2 (g)
What are the signs of ∆Hο and ∆S ο for a reaction that is non-spontaneous at low temperature but
spontaneous at high temperature?
Hο
Sο
A.
–
–
B.
+
–
C.
–
+
D.
+
+
4
19.
The equation for the decomposition of calcium carbonate is given below.
CaCO3 (s) → CaO(s) + CO2 (g)
At 500 K, ∆H for this reaction is +177 kJ mol–1 and ∆S is 161 J K–1 mol–1 .
(a)
Explain why ∆H for the reaction above cannot be described as ∆Hfο .
.....................................................................................................................................
.....................................................................................................................................
(2)
(b)
State the meaning of the term ∆S.
.....................................................................................................................................
.....................................................................................................................................
(1)
(c)
Calculate the value of ∆G at 500 K and determine, giving a reason, whether or not the
reaction will be spontaneous.
.....................................................................................................................................
.....................................................................................................................................
.....................................................................................................................................
(3)
(Total 6 marks)
20.
The standard enthalpy change for the combustion of phenol, C 6 H5 OH(s), is –3050 kJ mol–1
at 298 K.
(a)
Write an equation for the complete combustion of phenol.
......................................................................................................................................
......................................................................................................................................
(1)
(b)
The standard enthalpy changes of formation of carbon dioxide, CO 2 (g), and of water,
H2 O(l), are –394 kJ mol–1 and –286 kJ mol–1 respectively.
Calculate the standard enthalpy change of formation of phenol, C 6 H5 OH(s).
......................................................................................................................................
......................................................................................................................................
......................................................................................................................................
......................................................................................................................................
......................................................................................................................................
(3)
5
(c)
The standard entropy change of formation, ∆S ο , of phenol, C6 H5 OH(s) at
298 K is –385 J K–1 mol –1 . Calculate the standard free energy change of formation,
∆Gο , of phenol at 298 K.
......................................................................................................................................
......................................................................................................................................
......................................................................................................................................
......................................................................................................................................
......................................................................................................................................
......................................................................................................................................
......................................................................................................................................
(3)
(d)
Determine whether the reaction is spontaneous at 298 K, and give a reason.
......................................................................................................................................
......................................................................................................................................
......................................................................................................................................
......................................................................................................................................
(2)
(e)
Predict the effect, if any, of an increase in temperature on the spontaneity of this reaction.
......................................................................................................................................
......................................................................................................................................
......................................................................................................................................
......................................................................................................................................
(2)
(Total 11 marks)
21.
Explain in terms of Gο , why a reaction for which both Hο and S ο are positive is sometimes
spontaneous and sometimes not.
……………………………………………………………………………………………………
……………………………………………………………………………………………………
……………………………………………………………………………………………………
……………………………………………………………………………………………………
……………………………………………………………………………………………………
……………………………………………………………………………………………………
(Total 4 marks)
6
22.
Consider the following reaction.
N2 (g) + 3H2 (g) → 2NH3 (g)
(i)
Using the average bond enthalpy values in Table 10 of the Data Booklet, calculate the
standard enthalpy change for this reaction.
……………………………………………………………………………………………
……………………………………………………………………………………………
……………………………………………………………………………………………
……………………………………………………………………………………………
……………………………………………………………………………………………
……………………………………………………………………………………………
(4)
(ii)
The absolute entropy values, S, at 300 K for N2 (g), H3 (g) and NH2 (g) are 193, 131 and
192 JK–1 mol–1 respectively. Calculate S ο for the reaction and explain the sign of S ο .
……………………………………………………………………………………………
……………………………………………………………………………………………
……………………………………………………………………………………………
……………………………………………………………………………………………
……………………………………………………………………………………………
……………………………………………………………………………………………
(3)
(iii)
Calculate Gο for the reaction at 300 K.
……………………………………………………………………………………………
……………………………………………………………………………………………
……………………………………………………………………………………………
(1)
(iv)
If the ammonia was produced as a liquid and not as a gas, state and explain the effect this
would have on the value of Hο for the reaction.
……………………………………………………………………………………………
……………………………………………………………………………………………
……………………………………………………………………………………………
……………………………………………………………………………………………
(2)
(Total 10 marks)
7
23.
Define the term standard enthalpy of formation, and write the equation for the standard enthalpy
of formation of ethanol.
……………………………………………………………………………………………………
……………………………………………………………………………………………………
……………………………………………………………………………………………………
……………………………………………………………………………………………………
……………………………………………………………………………………………………
(Total 5 marks)
24.
Throughout this question, use relevant information from the Data Booklet.
(a)
Define the term standard enthalpy change of formation, and illustrate your answer with
an equation, including state symbols, for the formation of nitric acid.
(4)
(b)
Propyne undergoes complete combustion as follows:
C3 H4 (g) + 4O2 (g) → 3CO2 (g) + 2H2 O(l)
Calculate the enthalpy change of this reaction, given the following additional values:
HfӨ of CO2 (g) = –394 kJ mol–1
HfӨ of H2 O(l) = –286 kJ mol–1
(4)
(c)
Predict and explain whether the value of S Ө for the reaction in part (b) would be
negative, close to zero, or positive.
(3)
(Total 11 marks)
`25. (a)
Propyne reacts with hydrogen as follows:
C3 H4 (g) + 2H2 (g) → C3 H8 (g)
HӨ = –287 kJ
Calculate the standard entropy change of this reaction, given the following additional
information:
S Ө of H2 (g) = 131 J K–1 mol–1
(3)
(b)
Calculate the standard free energy change at 298 K, GӨ, for the reaction in part (a). Use
your answer and relevant information from part (d). If you did not obtain an answer to
part (a), use S Ө = –360 J K–1 (this is not the correct value).
(3)
(Total 6 marks)
26.
(a)
The lattice enthalpy of an ionic compound can be calculated using a Born-Haber cycle.
Using lithium fluoride as the example, construct a Born-Haber cycle, labelling the cycle
with the formulas and state symbols of the species present at each stage.
(6)
8
(b)
Two values of the lattice enthalpies for each of the silver halides are quoted in the Data
Booklet. Discuss the bonding in silver fluoride and in silver iodide, with reference to
these values.
(2)
(Total 8 marks)
27.
Hex-1-ene gas, C6 H12 , burns in oxygen to produce carbon dioxide and water vapour.
(a)
Write an equation to represent this reaction.
....................................................................................................................................
(1)
(b)
Use the data below to calculate the values of HcӨ and S cӨ for the combustion of
hex-1-ene.
O2 (g)
C6 H12 (g)
CO2 (g)
H2 O(g)
Standard enthalpy of
formation, HfӨ / kJ −1 mol
0.0
–43
–394
–242
Entropy, SӨ / J K−1 mol−1
205
385
214
189
Substance
(i)
Value of HcӨ
.........................................................................................................................
.........................................................................................................................
.........................................................................................................................
.........................................................................................................................
(2)
(ii)
Value of S cӨ
.........................................................................................................................
.........................................................................................................................
.........................................................................................................................
(2)
(c)
Calculate the standard free energy change for the combustion of hex-1-ene.
....................................................................................................................................
....................................................................................................................................
(2)
(d)
State and explain whether or not the combustion of hex-1-ene is spontaneous at 25C.
....................................................................................................................................
....................................................................................................................................
(1)
(Total 8 marks)
9
28.
Calculate the enthalpy change, H4 for the reaction
C + 2H2 +
1
2
O2 → CH3 OH
H4
using Hess’s Law, and the following information.
CH3 OH + 1 12 O2 → CO2 + 2H2 O
H1 = −676 kJ mol−1
C + O2 → CO2
H2 + 12 O2 → H2 O
H2 = −394 kJ mol−1
H3 = −242 kJ mol−1
..............................................................................................................................................
..............................................................................................................................................
..............................................................................................................................................
..............................................................................................................................................
..............................................................................................................................................
..............................................................................................................................................
(Total 4 marks)
29.
Methylamine can be manufactured by the following reaction.
CH3 OH(g) + NH3 (g) → CH3 NH2(g) + H2 O(g)
(a)
Define the term standard enthalpy change of formation.
....................................................................................................................................
....................................................................................................................................
....................................................................................................................................
(2)
(b)
The values of standard enthalpy changes of formation for some compounds are shown in
the table.
Compound
HfӨ / kJ mol–1
NH3 (g)
– 46
H2 O(g)
– 242
Predict, with a reason, whether the value of HfӨ for H2 O(l) is less than, greater than, or
equal to, the value of HfӨ for H2 O(g).
....................................................................................................................................
....................................................................................................................................
....................................................................................................................................
....................................................................................................................................
(2)
10
(c)
Use information from the table in (b) and from Table 11 of the Data Booklet to calculate
the enthalpy change for the reaction used to manufacture methylamine.
....................................................................................................................................
....................................................................................................................................
....................................................................................................................................
....................................................................................................................................
....................................................................................................................................
....................................................................................................................................
(3)
(Total 7 marks)
30.
(a)
Define the term standard enthalpy change of formation, HfӨ.
(2)
(b)
(i)
Use the information in the following table to calculate the enthalpy change for the
complete combustion of but-1-ene according to the following equation.
C4 H8 (g) + 6O2 (g) → 4CO2 (g) + 4H2 O(g)
Compound
C4 H8 (g)
CO2 (g)
H2 O(g)
HfӨ / kJ mol–1
+1
– 394
– 242
(3)
(ii)
Deduce, giving a reason, whether the reactants or the products are more stable.
(2)
(iii)
Predict, giving a reason, how the enthalpy change for the complete combustion of
but-2-ene would compare with that of but-1-ene based on average bond enthalpies.
(1)
(Total 8 marks)
31.
(i)
Define the term standard enthalpy change of formation, HfӨ.
(2)
(ii)
Construct a simple enthalpy cycle and calculate the value of HfӨ (C2 H5 OH(l)) given the
following data.
Compound
HfӨ / kJ mol–1
H2 O(l)
–286
CO2 (g)
–394
C2 H5 OH(l)
ΔHӨcomb/ kJ mol–1
–1371
(5)
(Total 7 marks)
11
32.
Consider the following reaction:
N2 (g) + 3H2 (g)
(i)
2NH3 (g)
Suggest why this reaction is important for humanity.
(1)
(ii)
Using the average bond enthalpy values in Table 10 of the Data Booklet, calculate the
standard enthalpy change for this reaction.
(4)
(iii)
The absolute entropy values, S, at 238 K for N2 (g), H2 (g) and NH3 (g) are 192, 131 and
193 J K–1 mol–1 respectively. Calculate ∆S ο for the reaction and explain the sign of ∆S ο .
(2)
(iv)
Calculate ∆Gο for the reaction at 238 K. State and explain whether the reaction is
spontaneous.
(3)
(v)
If ammonia was produced as a liquid and not as a gas, state and explain the effect this
would have on the value of ∆Hο for the reaction.
(2)
(Total 12 marks)
33.
(i)
Define the terms lattice enthalpy and electron affinity.
(2)
(ii)
Use the data in the following table and from the data booklet to construct the Born-Haber
cycle for sodium chloride, NaCl, and determine the lattice enthalpy of NaCl(s).
Na(s) +
1
Cl2 (g) → NaCl(g)
2
Na(s) → Na(g)
∆Hο = –411 kJ mol–1
∆Hο = +108 kJ mol–1
(4)
(iii)
Describe the structure of sodium chloride.
(2)
(Total 8 marks)
12
IB Chemistry HL
Topic5 Answers
1.
B
2.
3.
B
B
4.
5.
A
A
6.
C
7.
D
8.
A
9.
B
10.
D
11.
C
12.
C
13.
D
14.
B
15.
B
16.
C
17.
B
18.
D
19.
(a)
(cannot be ο as) conditions are not standard/at 500 K/OWTTE;
(cannot be f as) not formation from elements/is decomposition/OWTTE;
(b)
change in entropy/degree of (dis)order (of system);
(c)
∆G = 177000 – (500×161) = +96500;
reaction is not spontaneous;
∆G is positive;
Allow ECF from calculation for last two marks.
2
1
3
[6]
20.
(a)
C6 H5 OH + 7O2 → 6CO2 + 3H2 O;
Ignore state symbols.
(b)
∆Hrο =Σ∆Hfο products – Σ∆Hfο reactants;
–3050 = (6(–394) +3 (–286) – (∆Hfο phenol + O));
∆Hfο phenol =–172 kJ mol–1 ;
Award [3] for correct final answer.
1
3
13
Apply –1 (U) if appropriate.
Award [2 max] for ∆Hfο phenol = +172 kJ mol –1 .
(c)
(d)
(e)
appropriate conversion of units;
∆G = –172 – 298(– 0.385)
= –57.3 kJ mol–1 /–57 300 J mol–1 ;
Award [3] for correct final answer.
Accept answers in range –57.0 to –57.3 kJ mol–1.
Accept 3 sig. fig. only.
Allow ECF from (b).
Apply –1 (U) if appropriate.
3
spontaneous;
since ∆G is negative;
Allow ECF from (c).
2
reaction becomes less spontaneous;
∆G becomes less negative/more positive;
Accept a suitable calculation.
Allow ECF from (c).
2
[11]
ο
21. a reaction is spontaneous when ∆G is negative;
at high T, ∆Gο is negative;
–T∆S ο is larger/greater than ∆Hο ;
at low T, ∆Gο is positive because –T∆Sο is smaller than ∆Hο /OWTTE;
4
[4]
22.
(i)
∆H = (sum of energies of bonds broken) – (sum of energies of bonds formed);
Can be implied by working.
Correct substitution of values and numbers of bonds broken;
Correct substitution of values and numbers of bonds made;
(∆H = (N≡≡N) + 3(H—H) – 6(N—H) = 944 + 3(436) – 6(388) =) –76 (kJ);
4
Allow ECF.
Do not penalize for SF or units.
(ii)
∆S ο = (sum of entropies of products) – (sum of entropies of reactants);
Can be implied by working.
3
(= 2×192 – (193 + 3×131) =) –202(J K–1 mol–1 );
four molecules make two molecules/fewer molecules of gas;
(iii)
(∆Gο =∆Hο – T∆Sο = –76.0 – 300(–0.202)) = – 15.4 (kJ mol –1 );
Do not penalize for SF.
1
(iv)
∆Hο becomes more negative;
heat released when gas → liquid;
2
[10]
14
23. enthalpy change associated with the formation of one mole of a
compound/substance; from its elements;
in their standard states/under standard conditions;
2C(s) + 3H2 (g) +
1
2
O2 (g) → C2H5OH(l);
5
Award [1] for formulas and coefficients, [1] for state symbols.
[5]
24.
(a)
the enthalpy/energy/heat change for the formation of one mole of a
compound/substance from its elements;
in their standard states/under standard conditions/at 298 K and 1 atm;
1
2
H 2 (g) + 12 N 2 (g) + 1 12 O 2 (g) → HNO 3 (l) ;
4
Award [1] for correctly balanced equation, [1] for all state
symbols correct.
Do not award equation mark if 2HNO3 formed.
(b)
Hr = ∑HfӨ (products) − ∑HfӨ (reactants)/suitable cycle;
= 3( − 394) + 2( − 286) − 185;
Award [1] for correct coefficients of CO 2 and H2 O values, [1]
for correct value for C3 H4 from Data Booklet.
= −1939 or −1940 kJ;
Ignore units.
Award [4] for correct final answer.
Award [3] for +1939 or − 1569.
(c)
negative;
decrease in disorder/increase in order;
5 mol of gas → 3 mol of gas/reduction in number of gas moles;
Award [1] for answer of close to zero based on use of H 2 O(g).
4
3
[11]
25.
(a)
S = S Ө (products) − S Ө (reactants)/suitable cycle;
= 270 − 248 − 2131;
= − 240 (J K−1 );
Units not needed for mark, but penalize incorrect units.
Award [3] for correct final answer.
(b)
3
ΔGӨ = − 287 − (298−0.240);
Award [1] for correct substitution of values and [1] for
conversion of units.
= −215 kJ;
3
Units needed for mark.
Apply ECF from − 360 kJ or incorrect answer from (a).
[6]
15
26.
(a)
6
Li +(g) + e- +F(g)
Li +(g) + e- + 21 F2(g)
Li + (g) + F- (g)
Li(g) + 12 F2(g)
Li(s) + 12 F2(g)
LiF(s)
Award [6] for completely correct cycle, with endothermic
processes in any order.
Deduct [1] for each line in which species symbol and/or state
symbol is incorrect or missing.
Penalize missing electrons once only.
(b)
bonding in AgF more ionic than in AgI/bonding in AgI more covalent than
in AgF;
Accept AgF is ionic and AgI is covalent.
values closer/in better agreement in AgF/big(ger) difference in values for
AgI/OWTTE;
2
[8]
27.
(a)
C6 H12 + 9O2 → 6CO2 + 6H2 O;
(b)
(i)
(HӨ = ∑HfӨ products − ∑HfӨreactants)
HӨ = (6×−394 + 6×−242) − (−43);
HӨc = −3773/−3.8×103 (kJ mol−1 );
Accept 2, 3 or 4 sig. fig..
1
2
Award [1] for + 3773/+ 3.8 ×10 3 (kJ mol−1 ).
Allow ECF from (a) only if coefficients used.
(ii)
S Ө = (Sp Ө − SrӨ) = (6×189 + 6×214) − (385 + 9×205);
SӨ =188 (J K−1 mol−1 );
2
Accept only 3 sig. fig..
Award [1] for –188.
Allow ECF from (a) only if coefficients used.
(c)
(GӨc = HӨc − TSӨc) = −3800 − (298×0.188);
= − 3900 kJ mol−1 .
Accept − 3800 to − 3900.
Accept 2, 3 or 4 sig. fig.
Allow ECF from (b).
Units needed for second mark.
2
16
(d)
spontaneous and GӨ negative;
Allow ECF from (c).
1
[8]
28.
− 1×H1 /676;
1×H2 /− 394;
2×H3 /− 484;
H4 = −202 (kJ mol−1 );
Accept alternative methods.
Correct answers score [4].
4
Award [3] for (+)202 or (+)40 (kJ/kJ mol−1 ).
[4]
29.
(a)
(b)
(c)
enthalpy/energy change for the formation of 1 mol of a compound from
its elements;
Do not accept heat needed to form 1 mol…
in their standard states/under standard conditions/at 298 K and 1 atm;
2
greater value/more negative value;
energy given out when steam condenses/turns to water;
2
HӨ = ∑HfӨ (products) − ∑HfӨ (reactants)/suitable cycle;
= (−28−242)−(−201−46);
= −23 kJ/kJ mol−1 ;
Units needed for 3rd mark.
Correct final answer scores [3].
3
23 or +23 kJ/kJ mol−1 scores [2].
If −239 used instead of −201 for CH3OH, award [2] for +15 kJ.
[7]
30.
(a)
(b)
the enthalpy change when one mole of compound is formed from its elements
in their (standard state);
at (standard conditions of) 298 K/25C and 101 325 Pa/1 atm;
(i)
2
HP = (4×−242 + 4×−394) kJ mol−1 ;
HR = 1 kJ mol−1 ;
HӨ = (∑HӨ p −∑HӨ R) = −2545 /−2.55×103 / −2550 (kJ mol−1 );
3
Allow ECF.
(ii)
(iii)
products more stable than reactants;
bonds are stronger in products than reactants/Hp < HR/enthalpy/stored
energy of products less than reactants;
2
same/equal, because the same bonds are being broken and formed;
1
[8]
17
31.
(i)
change in energy for the formation of (1 mol) of a substance from its
elements; under standard conditions/1 atm pressure or 101 kPa and
298 K/25C;
2
(ii)

C 2 H 5 OH(l) + 3O 2 (g)

Hcomb
2CO 2 (g) + 3H 2O(l)

Hf (C 2 H 5OH(l))

H f (CO2 (g))
2C(s) + 2O2(g)
+
H f (H 2O(l))
3H 2 (g) + 1.5 O2 (g)
States not required.
Correct cycle showing:
Hcomb Ө
HfӨ (C2 H5 OH(l));
2HfӨ (CO2 (g)) and 3HfӨ (H2 O(l));
(HfӨ (C2 H5 OH(l)) = (2HfӨ (CO2 (g)) + 3HfӨ (H2 O(l)) − Hcomb Ө
= 2(−394) + 3(−286) + 1371;
= −275 kJ mol−1 ;
If values are substituted for symbols in the enthalpy cycle
diagram to give correct answer, award last [2] marks.
If no enthalpy cycle drawn but equation written and Hess’s
Law applied or calculated as follows, then [3 max]
5
(Hr = ∑Hf (products) − ∑ Hf (reactants))
−1371 = (−394×2) + (−286×3) −Hf (ethanol);
Hf (ethanol) = −788 − 858 + 1371;
= − 275(kJ mol−1 );
Award [2] for correct answer without enthalpy cycle and
without working and [1] for 275 or + 275.
[7]
32.
(i)
(ii)
fertilizers/increasing crop yields;
production of explosives for mining;
max
1
H = (sum of energies of bonds broken) – (sum of energies of bonds formed);
Can be implied by working.
correct substitution of values and numbers of bonds broken;
correct substitution of values and numbers of bonds made;
(H = (NN) + 3(H–H) – 6(N–H) = 944 + 3(436) – 6(388) =) –76.0 (kJ);
Allow ECF.
Do not penalize for sig. fig. or units.
Award [4] for correct final answer.
4
18
(iii)
(S ο [2×193] – [192 + 3×131]) = –199 (J K–1 mol–1 );
Allow ECF.
four gaseous molecules generating two gaseous
molecules/fewer molecules of gas;
(iv)
(Gο = Hο – TS ο = –76.0 – 298(–0.199)) = –16.7 (kJ);
Spontaneous;
G is negative;
Do not penalize for SF.
(v)
2
3
heat released when gas → liquid;
Hο becomes more negative;
2
[12]
33.
(i)
lattice enthalpy for a particular ionic compound is defined as ΔH for the
process, MX(s) → M+(g) + X– (g);
Accept definition for exothermic process
electron affinity is the energy change that occurs when an electron is added
to a gaseous atom or ion;
(ii)
2
H f = –411 kJ mol –1
Na(s)
+
+108 kJ mol –1
1
2
Cl 2 (g)
+121 kJ mol –1
Cl(g)
Na(g)
+494 kJ mol –1
Na+ (g)
NaCl(s)
–364 kJ mol –1
+
–
Cl (g)
lattice enthalpy = –[(–411) – (+108) – (+494) – (+121) – (–364)]
= 770 (kJ mol–1 )
Award [2] for all correct formulas in correct positions on cycle
diagram.
1 incorrect or missing label award [1].
Award [1] for all correct values in correct positions on cycle
diagram.
calculation of lattice enthalpy of NaCl(s) = 770 (kJ mol–1 );
Allow ECF.
Accept alternative method e.g. energy level diagram.
(iii)
lattice/network/regular structure;
each chloride ion is surrounded by six sodium ions and each sodium ion is
surrounded by six chloride ions/6:6 coordination;
4
2
[8]
19
IB Chemistry – SL
Topic 6 Questions
1.
2.
3.
Which of the following is (are) important in determining whether a reaction occurs?
I.
Energy of the molecules
II.
Orientation of the molecules
A.
I only
B.
II only
C.
Both I and II
D.
Neither I nor II
Consider the reaction between solid CaCO 3 and aqueous HCl. The reaction will be speeded up
by an increase in which of the following conditions?
I.
Concentration of the HCl
II.
Size of the CaCO3 particles
III.
Temperature
A.
I only
B.
I and III only
C.
II and III only
D.
I, II and III
Excess magnesium was added to a beaker of aqueous hydrochloric acid on a balance. A graph
of the mass of the beaker and contents was plotted against time (line 1).
Mass
1
2
Time
1
What change in the experiment could give line 2?
4.
5.
6.
I.
The same mass of magnesium but in smaller pieces
II.
The same volume of a more concentrated solution of hydrochloric acid
III.
A lower temperature
A.
I only
B.
II only
C.
III only
D.
None of the above
The rate of a reaction between two gases increases when the temperature is increased and a
catalyst is added. Which statements are both correct for the effect of these changes on the
reaction?
Increasing the temperature
Adding a catalyst
A.
Collision frequency increases
Activation energy increases
B.
Activation energy increases
Activation energy does not change
C.
Activation energy does not change
Activation energy decreases
D.
Activation energy increases
Collision frequency increases
Which of the following is (are) altered when a liquid at its boiling point is converted to a gas at
the same temperature?
I.
The size of the molecules
II.
The distance between the molecules
III.
The average kinetic energy of the molecules
A.
I only
B.
II only
C.
III only
D.
I and II only
Based on the definition for rate of reaction, which units are used for a rate?
A.
mol dm–3
B.
mol time–1
C.
dm time–1
D.
mol dm–3 time–1
2
7.
Which of the quantities in the enthalpy level diagram below is (are) affected by the use of a
catalyst?
Enthalpy
I
II
III
Time
8.
A.
I only
B.
III only
C.
I and II only
D.
II and III only
In the Haber process for the synthesis of ammonia, what effects does the catalyst have?
A.
B.
C.
D.
9.
10.
Rate of formation of NH 3 (g)
Increases
Increases
Increases
No change
Amount of NH 3 (g) formed
Increases
Decreases
No change
Increases
Which statement is correct for a collision between reactant particles leading to a reaction?
A.
Colliding particles must have different energy.
B.
All reactant particles must have the same energy.
C.
Colliding particles must have a kinetic energy higher than the activation energy.
D.
Colliding particles must have the same velocity.
Which change of condition will decrease the rate of the reaction between excess zinc granules
and dilute hydrochloric acid?
A.
increasing the amount of zinc
B.
increasing the concentration of the acid
C.
pulverize the zinc granules into powder
D.
decreasing the temperature
3
11.
The table shows the concentrations of reactants and products during this reaction.
2A + B → C + 2D
[A] / mol dm–3
[B] / mol dm–3
[C] / mol dm–3
[D] / mol dm–3
at the start
6
3
0
0
after 1 min
4
2
1
2
The rate of reaction can be measured by reference to any reactant or product. Which rates are
correct for this reaction?
I.
II.
III.
12.
13.
rate = –2 mol dm–3 min–1 for A
rate = –1 mol dm–3 min–1 for B
rate = –1 mol dm–3 min–1 for C
A.
I and II only
B.
I and III only
C.
II and III only
D.
I, II and III
In general, the rate of a reaction can be increased by all of the following except
A.
increasing the temperature.
B.
increasing the activation energy.
C.
increasing the concentration of reactants.
D.
increasing the surface area of the reactants.
At 25C, 100 cm3 of 1.0 mol dm–3 hydrochloric acid is added to 3.5 g of magnesium carbonate.
If the sample of magnesium carbonate is kept constant, which conditions will not increase the
initial rate of reaction?
Volume of HCl / cm3
Concentration of HCl / mol dm–3
Temperature / C
A.
200
1.0
25
B.
100
2.0
25
C.
100
1.0
35
D.
200
2.0
25
4
14.
15.
16.
17.
At 25C, 100 cm3 of 1.0 mol dm–3 hydrochloric acid is added to 3.5 g of magnesium carbonate.
If the sample of magnesium carbonate is kept constant, which conditions will not increase the
initial rate of reaction?
Volume of HCl / cm3
Concentration of HCl / mol dm–3
Temperature / C
A.
200
1.0
25
B.
100
2.0
25
C.
100
1.0
35
D.
200
2.0
25
Which statement is correct with regard to the catalysed and uncatalysed pathways for a given
reaction?
A.
The enthalpy change of the catalysed reaction is less than the enthalpy change for the
uncatalysed reaction.
B.
The enthalpy change of the catalysed reaction is greater than the enthalpy change for the
uncatalysed reaction.
C.
The enthalpy change of the catalysed reaction is equal to the enthalpy change for the
uncatalysed reaction.
D.
The activation energy of the catalysed reaction is greater than the activation energy for
the uncatalysed reaction.
Which changes increase the rate of a chemical reaction?
I.
Increase in the concentration of an aqueous solution
II.
Increase in particle size of the same mass of a solid reactant
III.
Increase in the temperature of the reaction mixture
A.
I and II only
B.
I and III only
C.
II and III only
D.
I, II and III
Excess magnesium, was added to a beaker of aqueous hydrochloric acid. A graph of the mass of
the beaker and contents was plotted against time (line 1).
Mass
1
2
Time
5
What change in the experiment could give line 2?
18.
A.
The same mass of magnesium in smaller pieces
B.
The same volume of a more concentrated solution of hydrochloric acid
C.
A lower temperature
D.
A more accurate instrument to measure the time
Which quantities in the enthalpy level diagram are altered by the use of a catalyst?
Enthalpy
I
II
III
Time
19.
A.
I and II only
B.
I and III only
C.
II and III only
D.
I, II and III
The graph below shows the volume of carbon dioxide gas produced against time when excess
calcium carbonate is added to x cm3 of 2.0 mol dm–3 hydrochloric acid.
Volume of CO 2
Time
(i)
Write a balanced equation for the reaction.
…………………………………………………………………………………………..
(1)
(ii)
State and explain the change in the rate of reaction with time. Outline how you would
determine the rate of the reaction at a particular time.
…………………………………………………………………………………………..
…………………………………………………………………………………………..
…………………………………………………………………………………………..
(4)
6
(iii)
Sketch the above graph on an answer sheet. On the same graph, draw the curves you
would expect if:
I.
the same volume (x cm3 ) of 1.0 mol dm–3 HCl is used.
II.
double the volume (2x cm3 ) of 1.0 mol dm–3 HCl is used.
Label the curves and explain your answer in each case.
…………………………………………………………………………………………..
…………………………………………………………………………………………..
…………………………………………………………………………………………..
…………………………………………………………………………………………..
…………………………………………………………………………………………..
…………………………………………………………………………………………..
(5)
(Total 10 marks)
20.
When excess lumps of magnesium carbonate are added to dilute hydrochloric acid the following
reaction takes place.
MgCO3 (s) + 2HCl(aq) → MgCl2 (aq) + CO2 (g) + H2 O(l)
(a)
Outline two ways in which the rate of this reaction could be studied. In each case sketch a
graph to show how the value of the chosen variable would change with time.
……………………………………………………………………………………………
……………………………………………………………………………………………
……………………………………………………………………………………………
……………………………………………………………………………………………
……………………………………………………………………………………………
……………………………………………………………………………………………
……………………………………………………………………………………………
……………………………………………………………………………………………
(4)
(b)
State and explain three ways in which the rate of this reaction could be increased.
……………………………………………………………………………………………
……………………………………………………………………………………………
……………………………………………………………………………………………
……………………………………………………………………………………………
(6)
7
(c)
State and explain whether the total volume of carbon dioxide gas produced would
increase, decrease or stay the same if
(i)
more lumps of magnesium carbonate were used.
……………………………………………………………………………………
……………………………………………………………………………………
……………………………………………………………………………………
(2)
(ii)
the experiments were carried out at a higher temperature.
……………………………………………………………………………………
……………………………………………………………………………………
……………………………………………………………………………………
(2)
(Total 14 marks)
21.
Carbon dioxide gas in the atmosphere reacts slightly with rainwater as shown below.
CO2 (g) + H2 O(l)
(i)
State the meaning of the
H+(aq) + HCO3 – (aq)
sign.
……………………………………………………………………………………………
(1)
(ii)
Predict the effect, if any, of the presence of a catalyst on the acidity of rainwater. Give a
reason for your answer.
……………………………………………………………………………………………
……………………………………………………………………………………………
……………………………………………………………………………………………
(2)
(iii)
Use Le Chatelier’s principle to predict the effect of the addition of a small quantity of an
alkali on the acidity of rainwater. Explain what effect, if any, this would have on the
equilibrium constant, Kc.
……………………………………………………………………………………………
……………………………………………………………………………………………
……………………………………………………………………………………………
……………………………………………………………………………………………
……………………………………………………………………………………………
……………………………………………………………………………………………
(3)
(Total 6 marks)
8
22.
Excess 0.100 mol dm–3 nitric acid is added to a certain mass of powdered calcium carbonate at
20C. The rate of reaction is monitored by measuring the change in mass over time due to the
loss of carbon dioxide.
2HNO3 ( aq) + CaCO3 (s)→ C a ( NO 3) 2 (aq) + H 2O (l) + CO 2 (g)
Mass loss / g
Time / minutes
(a)
Define the term rate of reaction.
....................................................................................................................................
....................................................................................................................................
(1)
(b)
Explain why the mass loss remains constant after a certain time.
....................................................................................................................................
....................................................................................................................................
(1)
(c)
Draw a line on the graph above, to show what the graph would look like if the same mass
of calcium carbonate in larger pieces were reacted with excess 0.100 mol dm–3 nitric acid.
(1)
(d)
Explain in terms of the collision theory what would happen to the rate if the reaction was
conducted at 50C.
....................................................................................................................................
....................................................................................................................................
....................................................................................................................................
....................................................................................................................................
....................................................................................................................................
(3)
9
(e)
Determine the rate of formation of carbon dioxide when the nitric acid reacts at a rate of
2.0010–3 mol cm–3 s–1 .
....................................................................................................................................
....................................................................................................................................
(1)
(Total 7 marks)
23.
(i)
Draw a graph that shows the distribution of molecular energies in a sample of a gas at two
different temperatures, T1 and T2 , such that T2 is greater than T1 .
(2)
(ii)
Define the term activation energy.
(1)
(iii)
State and explain the effect of a catalyst on the rate of an endothermic reaction.
(2)
(Total 5 marks)
24.
(i)
Magnesium is added to a solution of hydrochloric acid. Sketch a graph of acid
concentration on the y-axis against time on the x-axis to illustrate the progress of the
reaction.
(1)
(ii)
Describe how the slope of the line changes with time.
(1)
(iii)
Use the collision theory to state and explain the effect of decreasing concentration on the
rate of the reaction.
(2)
(Total 4 marks)
25.
The reaction between ammonium chloride and sodium nitrite in aqueous solution can be
represented by the following equation.
NH4 Cl(aq) + NaNO2 (aq) → N2 (g) + 2H2 O(l) + NaCl(aq)
The graph below shows the volume of nitrogen gas produced at 30 second intervals from a
mixture of ammonium chloride and sodium nitrite in aqueous solution at 20°C.
140
120
100
Volume of N 2 / cm 3
80
60
40
20
0
0
20 40
60
80 100 120 140 160 180
Time / s
10
(a)
(i)
State how the rate of formation of nitrogen changes with time. Explain your answer
in terms of collision theory.
............................................................................................................................
............................................................................................................................
............................................................................................................................
............................................................................................................................
............................................................................................................................
(2)
(ii)
Explain why the volume eventually remains constant.
............................................................................................................................
............................................................................................................................
(1)
(b)
(i)
State how the rate of formation of nitrogen would change if the temperature were
increased from 20°C to 40°C.
............................................................................................................................
............................................................................................................................
(1)
(ii)
State two reasons for the change described in (b)(i) and explain which of the two is
more important in causing the change.
............................................................................................................................
............................................................................................................................
............................................................................................................................
............................................................................................................................
............................................................................................................................
(3)
(iii)
The reaction between solid ammonium chloride and aqueous sodium nitrite can
be represented by the following equation.
NH4 Cl(s) + NaNO2 (aq) → N2 (g) + 2H2 O(l) + NaCl(aq)
State and explain how the rate of formation of nitrogen would change if the
same amount of ammonium chloride was used as large lumps instead of as a
fine powder.
............................................................................................................................
............................................................................................................................
............................................................................................................................
(2)
(Total 9 marks)
11
26.
(a)
Define the term average bond enthalpy, illustrating your answer with an equation for
methane, CH4 .
.....................................................................................................................................
.....................................................................................................................................
.....................................................................................................................................
(3)
(b)
The equation for the reaction between methane and chlorine is
CH4 (g) + Cl2 (g) → CH3 Cl(g) + HCl(g)
Use the values from Table 10 of the Data Booklet to calculate the enthalpy change for
this reaction.
.....................................................................................................................................
.....................................................................................................................................
.....................................................................................................................................
.....................................................................................................................................
.....................................................................................................................................
.....................................................................................................................................
(3)
(c)
Explain why no reaction takes place between methane and chlorine at room
temperature unless the reactants are sparked, exposed to UV light or heated.
.....................................................................................................................................
.....................................................................................................................................
.....................................................................................................................................
.....................................................................................................................................
(2)
(d)
Draw an enthalpy level diagram for this reaction.
(2)
(Total 10 marks)
12
27.
(a)
Identify two features of colliding molecules that react together in the gas phase.
.....................................................................................................................................
.....................................................................................................................................
.....................................................................................................................................
(2)
(b)
For many reactions, the rate approximately doubles for a 10°C rise in temperature.
State two reasons for this increase and identify which of the two is the more important.
.....................................................................................................................................
.....................................................................................................................................
.....................................................................................................................................
(3)
(Total 5 marks)
28.
(a)
Define the term rate of reaction.
...................................................................................................................................
...................................................................................................................................
(1)
(b)
The reaction between gases C and D is slow at room temperature.
(i)
Suggest two reasons why the reaction is slow at room temperature.
.........................................................................................................................
.........................................................................................................................
.........................................................................................................................
(2)
(ii)
A relatively small increase in temperature causes a relatively large increase in the
rate of this reaction. State two reasons for this.
.........................................................................................................................
.........................................................................................................................
(2)
(iii)
Suggest two ways of increasing the rate of reaction between C and D other than
increasing temperature.
.........................................................................................................................
.........................................................................................................................
.........................................................................................................................
.........................................................................................................................
(2)
(Total 7 marks)
13
29.
The reaction between two substances A and B
A+B →C+D
has the following rate expression:
rate = k [B]
Draw the graphical representation of:
[A] against time
[B] against time
[B]
[A]
time
time
rate against [A]
rate against [B]
rate
rate
[A]
[B]
(Total 3 marks)
14
IB Chemistry – SL
Topic 6 Answers
1.
C
2.
B
3.
B
4.
C
5.
B
6.
D
7.
C
8.
C
9.
C
10.
D
11.
A
12.
B
13.
A
14.
A
15.
C
16.
B
17.
B
18.
A
19.
(i)
CaCO3 (s) + 2H+(aq) → Ca2+(aq) + H2 O(l) + CO2 (g)
States not required, accept molecular equation.
(ii)
rate decreases with time;
as concentration decreases so fewer (successful) collisions;
draw tangent to the curve at time t;
rate = slope or gradient;
1
4
(iii)
15
Volume of
CO2 produced
orig
inal
II
I
Time
20.
(a)
I.
(less CO2 because) amount of HCl is limiting and half the
orginal/OWTTE;
II.
(same amount of CO2 because) amount of HCl is the same;
curve less steep because less frequent (accept fewer) collisions
measure volume of carbon dioxide/CO 2 /gas produced/measure pH;
5
4
starts at origin and levels off
measure mass of chemicals/apparatus;
starts high and decreases
Graph should show increase as reaction progresses
(as HCl is consumed).
(b)
Method 1
use powdered MgCO3 /OWTTE;
particles collide more frequently/increased surface area/OWTTE;
Method 2
increase (reaction) temperature/heat/warm;
more of the collisions are successful/more particles with E > Ea/OWTTE;
Method 3
increase acid concentration;
more frequent (reactant) collisions;
Method 4
16
add catalyst;
lowers activation energy/Ea/OWTTE;
Award [2] each for any three methods
(c)
(i)
(ii)
stays the same;
MgCO3 was already in excess;
6 max
2
stays the same;
same quantities of reactants used;
2
[14]
21.
(i)
(ii)
(iii)
reversible reaction/reaction may proceed in either direction
(depending on reaction conditions) equilibrium/dynamic equilibrium;
1
no effect;
catalyst will speed up both forward and reverse reactions (equally)/
increase the rate at which equilibrium is achieved;
2
acidity: no effect;
equilibrium shifts to the right;
Kc: no change;
3
[6]
22.
(a)
change of concentration/mass/amount/volume/of a reactant/product with time;
Do not accept “substance”.
(b)
all the CaCO3 (s) has been consumed/no further CO 2 (g) is produced/reaction
is complete;
Do not accept reaction has stopped or all reactants used up.
(c)
line on graph should be initially less steep/a smaller gradient and should
plateau at the same mass loss;
1
1
1
Mass loss / g
Time / minutes
17
(d)
there are more particles with KE greater than or equal to Ea ;
collisions more frequent/more collisions per unit time/more
successful/forceful collisions per unit time;
the rate increases;
(e)
1.00×10−3 (mol cm−3 s−1 )
Ignore units even if wrong.
Apply −1(sf).
3
1
[7]
(i)
Number/Fraction of Molecules
23.
T1
T2 > T 1
T2
Energy/speed
T2 peak lower/T1 higher;
T2 peak at higher energies/T1 curve at lower energies;
Maximum [1] if axes not labeled correctly
(ii)
(iii)
2
minimum energy required to react/energy difference between
reactants and transition state;
1
makes the reaction go faster;
because it lowers the activation energy/Ea ;
2
[5]
24.
(i)
a curve showing concentration decreases with time;
1
i.e.
[acid]
Time
No penalty if curve reaches x axis
Do not accept a straight line
18
(ii)
slope decreases;
1
(iii)
rate decreases;
fewer collisions per unit time;
2
[4]
25.
(a)
(b)
(i)
it is decreasing;
less frequent collisions/fewer collisions per second or (unit) time;
2
(ii)
reactant(s) used up/reaction is complete;
Do not accept reaction reaches equilibrium.
(i)
it would increase;
Accept a quantitative answer such as “doubles”.
(ii)
more frequent collisions;
collisions or molecules have more
energy (OWTTE); more molecules with energy  Ea;
3
rate would be lower;
smaller surface area;
2
(iii)
1
1
[9]
26.
(a)
(b)
(c)
energy for the conversion of a gaseous molecule into (gaseous) atoms;
(average values) obtained from a number of similar bonds/compounds/OWTTE;
CH4 (g) → C(g) + 4H(g);
State symbols needed.
(bond breaking) = 1890/654;
(bond formation) = 2005/769;
enthalpy = –115(kJ mol–1 )
Allow ECF from bond breaking and forming.
Award [3] for correct final answer.
Penalize [1] for correct answer with wrong sign.
molecules have insufficient energy to react (at room temperature)/
wrong collision geometry/unsuccessful collisions;
extra energy needed to overcome the activation energy/Ea for the reaction;
3
3
2
(d)
Ea
energy
reactants
products
reaction path
exothermic shown;
activation energy/Ea shown;
Allow ECF from (b).
2
[10]
19
27.
(a)
(b)
molecules must have sufficient/minimum energy/energy ≥ activation energy;
appropriate collision geometry/correct orientation;
2
increased frequency of collisions/collisions more likely;
Not just “more collisions”, there must be a reference to time.
increased proportion of molecules with sufficient energy to react/E ≥ Ea;
Not “activation energy is reduced”.
Proportion of molecules with E ≥ Ea is more important;
(dependent on correct second marking point);
28.
(a)
(b)
increase in product concentration per unit time/decrease in reactant concentration
per unit time;
Accept change instead of increase or decrease.
(i)
(ii)
(iii)
high activation energy/not enough molecules have Ea /OWTTE;
incorrect collision geometry/OWTTE;
infrequent collisions;
Award [1] for any two reasons.
more energetic collisions/more molecules have (energy ) Ea;
more frequent collisions/collide more often;
add a catalyst;
increase the (total) pressure/decrease the volume of the container;
increase the concentration of C (or D);
Do not accept surface area.
Award [1] for any two.
3
1
2
2
2
[7]
29.
([A] against time) - straight line with negative gradient;
Accept any decreasing curve
([B] against time) - decreasing curve;
Award [1] unless half - lives clearly not constant
(rate against [A]) - any horizontal straight line;
(rate against [B]) - straight line through origin;
Award [3] for all four correct, award [2] for any three correct
and [1] for any two correct.
20
IB Chemistry – HL
Topic 6 Questions
1.
2.
3.
The reaction between NO 2 and F2 gives the following rate data at a certain temperature.
What is the order of reaction with respect to NO 2 and F2 ?
[NO2 ]/mol dm–3
[F2 ]/mol dm–3
0.1
0.2
0.1
0.2
0.2
0.4
0.1
0.4
0.2
NO2 order
F2 order
A.
first
first
B.
first
second
C.
second
first
D.
second
second
Rate /mol dm–3
min–1
Which step in a multi-step reaction is the rate determining step?
A.
The first step
B.
The last step
C.
The step with the lowest activation energy
D.
The step with the highest activation energy
The rate expression for a reaction is shown below.
rate = k[A]2 [B]2
Which statements are correct for this reaction?
I.
The reaction is second order with respect to both A and B.
II.
The overall order of the reaction is 4.
III.
Doubling the concentration of A would have the same effect on the rate of reaction
as doubling the concentration of B.
A.
I and II only
B.
I and III only
C.
II and III only
D.
I, II and III
1
4.
5.
Values of a rate constant, k, and absolute temperature, T, can be used to determine the activation
energy of a reaction by a graphical method. Which graph produces a straight line?
A.
k versus T
B.
k versus 1
T
C.
ln k versus T
D.
ln k versus 1
T
The rate expression for a particular reaction is
Rate = k[P][Q]
Which of the units below is a possible unit for k?
6.
A.
mol–2 dm6 min–1
B.
mol–1 dm3 min–1
C.
mol dm3 min–1
D.
mol–2 dm–6 min–1
The reaction 2X(g) + Y(g) → 3Z(g) has the rate expression
rate = k [X]2 [Y]0
The concentration of X is increased by a factor of three and the concentration of Y is increased
by a factor of two. By what factor will the reaction rate increase?
7.
A.
6
B.
9
C.
12
D.
18
A reaction occurs in four steps. The steps and their rates are shown in the table
Step
Rate
1
0.01 mol dm–3 s–1
2
0.10 mol dm–3 s–1
3
0.01 mol dm–3 min–1
4
0.10 mol dm–3 min–1
Which is the rate-determining step?
A.
Step 1
B.
Step 2
C.
Step 3
D.
Step 4
2
8.
The rate expression for a reaction is
rate = k[CH3 Br][OH– ]
Which is a possible unit for k?
9.
10.
A.
mol2 dm–6 min–1
B.
mol dm–3 min–1
C.
mol–1 dm3 min–1
D.
mol–2 dm6 min–1
What happens to the rate constant (k) and activation energy (Ea) of a reaction when the
temperature is increased?
A.
k increases and Ea is unaffected.
B.
k decreases and Ea is unaffected.
C.
Ea increases and k is unaffected.
D.
Ea decreases and k is unaffected.
The mechanism of a reaction is
XY2 + XY2 → X2 Y4
X2 Y4 → X2 + 2Y2
X2 + Y2 → 2XY
What is the overall equation for the reaction?
11.
A.
X2 Y4 → 2XY2
B.
2XY2 → X2 + 2Y2
C.
2XY2 → 2XY + Y2
D.
X2 Y4 → 2XY + Y2
Consider the reaction
2I−(aq) + H2 O2 (aq) + 2H+(aq) → I2 (aq) + 2H2 O(l)
In the presence of S2 O3 2– (aq) and starch solution, the time taken for a blue colour to form was
observed at various reactant concentrations.
Experiment
[I– ] / mol dm–3
[H2 O2 ] / mol dm–3
[H+] / mol dm–3
Time / s
1
0.10
0.12
0.01
25
2
0.05
0.12
0.01
50
3
0.10
0.06
0.01
100
What is the correct order with respect to I– and H2 O2 ?
I–
H2 O2
A.
1
2
B.
1
2
1
4
C.
2
1
D.
2
4
3
12.
13.
Which statement is correct about the rate expression for a chemical reaction?
A.
It can be determined from its chemical equation.
B.
It can be predicted from the value of ΔH Ө for the reaction.
C.
It can be calculated from the effect of temperature on the reaction.
D.
It can be determined by measuring the change in concentration of a reactant or product
over time.
For the reaction 2NO2 (g) + F2 (g) → 2NO2 F(g) the accepted mechanism is
NO2 (g) + F2 (g) → NO2 F(g) + F(g)
NO2 (g) + F(g) → NO2 F(g)
slow
fast
What is the rate expression for this reaction?
14.
15.
A.
rate = k[NO2 ]2 [F2]
B.
rate = k[NO2 ][F2 ]
C.
rate = k[NO2 ][F]
D.
rate = k[NO2 ]2
The activation energy, of a reaction can be obtained from the rate constant, k, and the absolute
temperature, T. Which graph of these quantities produces a straight line?
A.
k against T
B.
k against
C.
ln k against T
D.
ln k against
1
T
1
T
What is the order of reaction with respect to NO 2 (g) and F2 (g) given the following rate data at a
certain temperature?
[NO2 (g)] / mol dm–3
[F2 (g)] / mol dm–3
Rate / mol dm–3 min–1
0.1
0.2
0.1
0.2
0.2
0.4
0.1
0.4
0.2
Order with respect to NO 2(g)
Order with respect to F2 (g)
A.
first
first
B.
first
second
C.
second
first
D.
second
second
4
16.
Nitrogen(II) oxide reacts with hydrogen as shown by the following equation.
2NO(g) + 2H2 (g) → N2 (g) + 2H2 O(g)
The table below shows how the rate of reaction varies as the reactant concentrations vary.
Experiment
1
2
3
4
(a)
Initial [NO] /
mol dm–3
0.100
0.100
0.200
0.300
Initial [H2 ] / mol
dm–3
0.100
0.200
0.100
0.100
Initial rate /
mol N2 dm–3 s –1
2.53×10–6
5.05×10–6
10.10×10–6
22.80×10–6
Determine the order of reaction with respect to NO and with respect to H 2 .
Explain how you determined the order for NO.
NO ..............................................................................................................................
.....................................................................................................................................
H2 ................................................................................................................................
.....................................................................................................................................
(3)
(b)
Write the rate expression for the reaction.
.....................................................................................................................................
(1)
(c)
Calculate the value for the rate constant, including its units.
.....................................................................................................................................
.....................................................................................................................................
.....................................................................................................................................
(2)
(d)
A suggested mechanism for this reaction is as follows.
H2 + NO
X fast step
X + NO → Y + H2 O slow step
Y + H2 → N2 + H2 O fast step
State and explain whether this mechanism agrees with the experimental rate expression
in (b).
.....................................................................................................................................
.....................................................................................................................................
.....................................................................................................................................
.....................................................................................................................................
(4)
5
(e)
Explain why a single step mechanism is unlikely for a reaction of this kind.
.....................................................................................................................................
.....................................................................................................................................
.....................................................................................................................................
.....................................................................................................................................
(2)
(f)
Deduce how the initial rate of formation of H 2 O(g) compares with that of N 2 (g) in
experiment 1. Explain your answer.
.....................................................................................................................................
.....................................................................................................................................
.....................................................................................................................................
.....................................................................................................................................
.....................................................................................................................................
(2)
(Total 14 marks)
17.
The oxidation of nitrogen monoxide takes place as follows:
2NO(g) + O2 (g) → 2NO2 (g)
The following experimental data was obtained at 101.3 kPa and 298 K.
Experiment
Initial [NO] / mol dm–3
Initial [O2 ] / mol dm–3
Initial rate / mol dm–3 s–1
1
3.50×10–2
1.75×10–2
3.75×10–3
2
3.50×10–2
3.50×10–2
7.50×10–3
3
7.00×10–2
7.00×10–2
6.00×10–2
(a)
Deduce the order of reaction with respect to O 2 .
...................................................................................................................................
...................................................................................................................................
(1)
(b)
Deduce the order of reaction with respect to NO.
...................................................................................................................................
...................................................................................................................................
(1)
(c)
State the rate expression for the reaction.
...................................................................................................................................
(1)
6
(d)
Calculate the value of the rate constant and state the units.
...................................................................................................................................
...................................................................................................................................
...................................................................................................................................
...................................................................................................................................
(2)
(e)
Suggest a possible mechanism that is consistent with the rate expression. Indicate which
of the steps is the rate-determining step.
...................................................................................................................................
...................................................................................................................................
...................................................................................................................................
...................................................................................................................................
(3)
(Total 8 marks)
18.
An equation for the decomposition of substance A is
2A → 2B + C
A graph showing the change in concentration of A against time as the reaction proceeds at a
particular temperature is shown below.
0.10
0.09
0.08
0.07
0.06
[A] /
0.05
mol dm–3
0.04
0.03
0.02
0.01
0
1000
2000
3000
4000
5000
Time / s
(a)
Define the term half-life of reaction.
...................................................................................................................................
...................................................................................................................................
(1)
7
(b)
Use the graph to measure values of half-life of reaction, starting from
time = zero ................................................................................................................
time = 1000 s ............................................................................................................
(2)
(c)
Deduce the order of the reaction with respect to A, giving a reason for your choice, and
write the rate expression for the reaction.
...................................................................................................................................
...................................................................................................................................
(3)
(d)
For a different reaction, between compounds D and E, the rate expression is
rate = k[D]2 [E]
Calculate the value of k, including units, for the reaction when the concentrations of both
D and E are 1.35×10–2 mol dm–3 and the reaction rate is 3.75×10–5 mol dm–3 min–1 .
(3)
(Total 9 marks)
In a particular experiment, various concentrations of HI(aq) are reacted with a constant
H2 O2 (aq) concentration according to the following equation:
H2 O2 (aq) + 2HI(aq) → I2 (aq) + 2H2 O(I)
A graph of [HI] against time is as follows:
16.0
14.0
12.0
[HI] / mol dm–3 10 –2
19.
10.0
8.0
6.0
4.0
2.0
0
20
40
60
80
Time / sec
100
120
140
160
8
(a)
Use the graph to deduce the order of reaction with respect to HI. Give a reason for your
answer.
....................................................................................................................................
....................................................................................................................................
....................................................................................................................................
....................................................................................................................................
(2)
(b)
The order with respect to H 2 O2 is the same as HI. Deduce the rate expression for this
reaction.
....................................................................................................................................
....................................................................................................................................
(1)
(c)
Determine the half-life of the reaction from the graph and calculate the value for the rate
constant.
....................................................................................................................................
....................................................................................................................................
....................................................................................................................................
....................................................................................................................................
(2)
(Total 5 marks)
20.
(a)
The table below shows kinetic data for the following reaction
C+D →E+F
(i)
Experiment
[C] / mol dm–3
[D] / mol dm–3
Initial rate
/ mol dm–3 min–1
1
2.0×10–3
3.0×10–3
1.0×10–6
2
4.0×10–3
3.0×10–3
2.0×10–6
3
6.0×10–3
6.0×10–3
3.0×10–6
Deduce the order of reaction with respect to both C and D, giving a reason in each
case.
C ......................................................................................................................
.........................................................................................................................
D ......................................................................................................................
.........................................................................................................................
(4)
9
(ii)
Deduce the rate expression for this reaction.
.........................................................................................................................
.........................................................................................................................
(1)
(iii)
Use data from Experiment 1 to calculate a value for the rate constant for this
reaction and deduce its units.
.........................................................................................................................
.........................................................................................................................
.........................................................................................................................
.........................................................................................................................
(3)
(b)
Define the term half-life and calculate the half-life for a first-order reaction with a rate
constant of 3.310–2 min–1 .
....................................................................................................................................
....................................................................................................................................
....................................................................................................................................
....................................................................................................................................
(2)
(Total 10 marks)
21.
Nitrogen(II) oxide reacts with hydrogen according to the following equation:
2NO(g) + 2H2 (g) → N2 (g) + 2H2 O(g)
The table shows how the rate of reaction varies as the concentrations of the reactants are
changed.
(a)
Experiment
Initial [NO] /
mol dm–3
Initial [H2 ] /
mol dm–3
Initial rate /
mol (N2 ) dm–3 s –1
1
0.100
0.100
253×10–6
2
0.100
0.200
5.05×10–6
3
0.200
0.100
1.01×10–5
4
0.300
0.100
2.28×10–5
Determine the order of reaction with respect to H 2 and with respect to NO.
H2 ................................................................................................................................
NO ..............................................................................................................................
(2)
10
(b)
Write the rate expression for the reaction.
.....................................................................................................................................
(1)
(c)
Calculate the value for the rate constant, and state its units using the data from
experiment 1.
.....................................................................................................................................
.....................................................................................................................................
.....................................................................................................................................
.....................................................................................................................................
(2)
(d)
A suggested mechanism for this reaction is as follows.
H2 + NO
X
X + NO → Y + H2O
Y + H2 → N2 + H2O
fast step
slow step
fast step
State and explain whether this mechanism agrees with the experimental rate expression in
(b).
.....................................................................................................................................
.....................................................................................................................................
.....................................................................................................................................
.....................................................................................................................................
.....................................................................................................................................
.....................................................................................................................................
(4)
(e)
Explain why a single step mechanism is unlikely for a reaction of this kind.
.....................................................................................................................................
.....................................................................................................................................
.....................................................................................................................................
(2)
(f)
Deduce and explain how the initial rate of formation of H 2 O compares with that of N 2 .
.....................................................................................................................................
.....................................................................................................................................
.....................................................................................................................................
(2)
(Total 13 marks)
11
22.
The data below refer to a reaction between X and Y.
Initial concentration /
mol dm–3
(i)
Initial rate of reaction /
mol dm–3 s –1
Experiment
X
Y
1
0.25
0.25
10×10–2
2
0.50
0.25
4.0×10–2
3
0.50
0.50
8.0×10–2
Define the term order of reaction.
…………………………………………………………………………………………..
…………………………………………………………………………………………..
…………………………………………………………………………………………..
(1)
(ii)
Deduce the order of reaction with respect to both X and Y. Explain your reasoning.
…………………………………………………………………………………………..
…………………………………………………………………………………………..
…………………………………………………………………………………………..
…………………………………………………………………………………………..
…………………………………………………………………………………………..
…………………………………………………………………………………………..
(4)
(iii)
Write the rate expression for the reaction and calculate the rate constant, including its
units.
…………………………………………………………………………………………..
…………………………………………………………………………………………..
…………………………………………………………………………………………..
…………………………………………………………………………………………..
(4)
(iv)
Calculate the initial rate of reaction when the initial concentrations of X and Y are 0.40
mol dm–3 and 0.60 mol dm–3 respectively.
…………………………………………………………………………………………..
…………………………………………………………………………………………..
(2)
(Total 11 marks)
12
23.
Oxygen and nitrogen monoxide react together to form nitrogen dioxide.
O2 (g) + 2NO(g) → 2NO2 (g)
The graph below shows how the initial rate of reaction changed during an experiment in which
the initial [NO(g)] was kept constant whilst the initial [O 2 (g)] was varied.
Rate
[O 2 (g)]
(a)
Deduce, giving a reason, the order of reaction with respect to O 2
……………………………………………………………………………………………
……………………………………………………………………………………………
(2)
(b)
In a series of experiments, the initial [O 2 (g)] was kept constant while the initial [NO(g)]
was varied. The results showed that the reaction was second order with respect to NO.
Sketch a graph to show how the rate of reaction would change if the initial
[NO(g)] was increased.
……………………………………………………………………………………………
……………………………………………………………………………………………
……………………………………………………………………………………………
……………………………………………………………………………………………
……………………………………………………………………………………………
(2)
(c)
Deduce the overall order of this reaction.
……………………………………………………………………………………………
……………………………………………………………………………………………
(1)
(d)
State and explain what would happen to the initial rate of reaction if the initial
concentration of NO was doubled and that of O 2 was halved.
……………………………………………………………………………………………
……………………………………………………………………………………………
……………………………………………………………………………………………
(3)
13
(e)
When the initial values are [O 2 (g)] = 1.0×10–2 mol dm–3 and [NO(g)] = 3.0×10–2 mol
dm–3 , the initial rate of reaction is 6.3×10–4 mol dm–3 s–1 . Write the rate expression for
this reaction and calculate the rate constant, stating its units.
……………………………………………………………………………………………
……………………………………………………………………………………………
(4)
(Total 12 marks)
24.
The compound iodine chloride, ICl, reacts with hydrogen to form iodine and hydrogen chloride.
(i)
Deduce the equation for this reaction.
(1)
(ii)
The kinetics of this reaction were studied at a certain temperature, when all the reactants
and products were in the gas phase. The table shows the initial rate of reaction for
different concentrations of reactants.
Experiment
[ICl] / mol dm–3
[H2 ] / mol dm–3
Initial rate /
mol dm–3 s–1
1
0.100
0.0500
5.00×10–3
2
0.200
0.0500
1.00×10–2
3
0.200
0.0250
2.50×10–3
Deduce and explain the order of reaction with respect to ICl and to H2 .
(4)
(iii)
Write the rate expression for the reaction.
(1)
(iv)
Use information from Experiment 1 to determine the value, with units, of the rate
constant for the reaction.
(2)
(v)
Determine the rate of reaction when the concentrations of reactants in Experiment 1 are
both doubled.
(1)
(Total 9 marks)
25.
(a)
The variation of the rate constant, k, for a reaction with temperature is shown by the
Arrhenius equation. Two versions of this equation are shown in Table 1 of the Data
Booklet.
(i)
Explain the significance of the Arrhenius constant, A, in this equation.
(1)
(ii)
Explain what is meant by the term activation energy,Ea.
(1)
(iii)
Describe how, using a graphical method, values of A and Ea can be obtained for a
reaction.
(5)
14
(b)
The equation for a reaction used in industry is
CH2 CH2 + Cl2 → CH2 ClCH2 CL HӨ = –185 kJ
Iron(III) chloride can be used as a catalyst for the reaction.
(i)
Explain the difference between the terms homogeneous and heterogeneous when
applied to a catalyst.
(1)
(ii)
Draw an enthalpy level diagram for this reaction, including labels for HӨ, Ea and
the activation energy when a catalyst is used, Ecat .
(4)
(Total 12 marks)
26.
Nitrogen(II) oxide reacts with bromine according to the following equation.
2NO(g) + Br2 (g) → 2NOBr(g)
H = negative
The data below were obtained for the reaction between NO(g) and Br 2 (g) at a specified
temperature and pressure.
(a)
Experiment
Initial [NO] /
mol dm–3
Initial [Br2 ] /
mol dm–3
Initial rate / mol dm–3 s–1
1
2.00×10–2
5.00×10–3
3.20×10–3
2
2.00×10–2
2.50×10–3
1.60×10–3
3
4.00×10–2
5.00×10–3
1.30×10–2
Determine, giving a reason, the order of reaction with respect to NO and the order of
reaction with respect to Br2 .
....................................................................................................................................
....................................................................................................................................
....................................................................................................................................
....................................................................................................................................
(2)
(b)
Derive the rate expression for the reaction between NO and Br 2 .
....................................................................................................................................
....................................................................................................................................
(1)
(c)
Calculate the rate constant for the rate expression using experiment 1 and state its units.
....................................................................................................................................
....................................................................................................................................
....................................................................................................................................
(2)
15
(d)
If the total volume of the reaction mixture was doubled at constant temperature, state the
effect, if any, on
(i)
the rate constant.
.........................................................................................................................
.........................................................................................................................
(1)
(ii)
the rate of change of the Br2 (g) concentration.
.........................................................................................................................
.........................................................................................................................
(1)
(e)
Draw a labelled enthalpy level diagram for the reaction between NO(g) and Br 2 (g), with
and without the use of a catalyst.
(3)
(Total 10 marks)
27.
(i)
The reaction between propanone, CH 3 COCH3 and bromine, Br2 in the presence of acid,
H+, is found to be second order overall, but the rate is independent of the bromine
concentration. Write three possible rate expressions for the reaction.
(3)
(ii)
The concentration of each of the three reactants was doubled in three separate
experiments. Choose one of the rate expressions in (i) and predict the effect on the rate of
the reaction of each of these changes.
(2)
16
(iii)
The graph below shows how the concentration of propanone changes with time in a
reaction.
1.8
1.6
1.4
1.2
Concentration
/ mol dm–1
1.0
0.8
0.6
0.4
0.2
0.0
0.0
5.0
10.0
15.0
20.0
25.0
30.0
35.0
40.0
Time / s
Use the graph to confirm that the reaction is first order with respect to propanone showing
your working.
(2)
(iv)
The overall reaction is:
+
⎯→ CH3 COCH2Br(aq) + HBr(aq)
CH3 COCH3 (aq) + Br2 (aq) ⎯H⎯(aq)
Describe one observation that would allow you to follow the progress of the reaction.
State and explain the role of the acid in the reaction.
(4)
(Total 11 marks)
17
IB Chemistry – HL
Topic 6 Answers
1.
C
2.
D
3.
4.
D
D
5.
6.
B
B
7.
C
8.
C
9.
A
10.
C
11.
A
12.
D
13.
B
14.
D
15.
C
16.
(a)
(order with respect to) NO = 2;
(order with respect to) H2 = 1;
rate increases×4 when [NO] doubles/OWTTE;
(b)
rate = k[NO]2 [H2 ];
ECF from (a).
(c)
(2.53×10–6 mol dm–3 s–1 = k (0.100 mol dm–3 )2 (0.100 mol dm–3 ))
k = 2.53×10–3 ;
mol –2 dm6 s–1 ;
ECF from (b).
(d)
3
1
1
1
agrees/yes;
slow step depends on X and NO;
X depends on H2 and NO;
(so) NO is involved twice and H 2 once;
Overall equation matches the stoichiometric equation;
Award [1] each for any three of the four above.
OWTTE
ECF for “no”, depending on answer for (b).
18
Or
agrees/yes;
X = constant;
and
H 2 NO
rate of slow step = k [X][NO]
= k [H2 ][NO]2 ;
4
ECF for “no”, depending on answer for (b).
(e)
(f)
reaction involves four molecules;
statistically/geometrically unlikely;
2
the rate of formation of H2 O(g) = 2×rate for N2 (g);
because 2 moles H2O formed with 1 mole N 2 /OWTTE;
2
[14]
17.
(a)
first order (with respect to O 2 );
1
(b)
second order (with respect to NO);
1
(c)
rate = k[NO]2[O2];
Allow ECF from parts (a) and (b).
1
(d)
k=
3.75 10 −3
(3.50 10 − 2 ) 2 (1.75 10 − 2 )
= 1.75 10 2 ;
dm6 mol–2 s–1 ;
Award [1] mark for the answer and [1] mark for units.
Allow ECF from part (c).
(e)
NO + NO
2
N2 O2 ;
N2 O2 + O2 → 2NO2 ;
second step is rate determining step;
Allow ECF from part (c).
OR
NO + O2
NO3 ;
NO3 + NO → 2NO2 ;
second step is rate determining step;
Allow ECF from part (c).
3
[8]
18.
(a)
time for reactant concentration to halve/OWTTE;
Accept “time for mass to halve”.
(b)
1000 s;
1000 s;
1
2
Accept 900-1100 s.
19
(c)
(d)
first order;
constant half-life;
rate = k[A];
Allow ECF for rate expression from stated order.
k=
3
3.75  10 −5
rate
/
;
[D] 2 [E] (1.35  10 − 2 ) 3
= 15.2;
Accept answer in range 15.2 to 15.3.
mol−2 dm6 min−1 ;
3
[9]
19.
(a)
first order;
constant half-life;
(b)
rate = k[HI][H2O2];
ECF from(a).
(c)
47 sec;
2
1
Accept answer in range 45 to 49.


 t 1 = 0.693  = 0.015 ;

k 
 2
Accept answer in range 0.014-0.015.
ECF from half-life.
2
[5]
20.
(a)
(i)
(C)
first order;
doubling [C] doubles rate/OWTTE;
(D)
zero order;
changing [D] has no effect on rate/OWTTE;
(ii)
rate = k[C]/rate = k[C]1 [D]0;
Apply ECF from (a)(i).
(iii)
k=
4
1
rate 1.0 10 −6
/
;
[C ] 2.0 10 −3
= 5.0×10−4 ;
min−1 ;
Apply ECF from (a)(ii).
(b)
time for half of (amount/concentration of) reactant to react/disappear;
1
t ( = 0.693÷0.033) = 21 min;
2
Units needed for second mark.
3
2
[10]
20
21.
(a)
(order with respect to) H2 = 1;
(order with respect to) NO = 2;
(b)
rate = k[H2 ][NO]2;
ECF from (a).
(c)
(2.53×10−6 mol dm–3 s–1 = k(0.100 mol dm−3 )(0.100 mol dm–3 )2 )
k = 2.53×10–3 ;
mol−2 dm6 s–1 ;
ECF from (b).
(d)
2
1
2
agrees/yes;
slow step depends on X and NO;
(so) NO is involved twice and H 2 once;
overall equation matches the stoichiometric equation/OWTTE;
ECF for “no”, depending on answer for (b).
OR
agrees/yes;
[X]
and
= constant;
[H 2 ][ NO]
rate of slow step = k[X][NO];
but X depends on H2 and NO;
rate of slow step = k[H2 ][NO]2 ;
max
Award [1] each for any three of the four above.
ECF for “no”, depending on answer for (b).
(e)
(f)
4
reaction involves four molecules;
statistically/geometrically unlikely;
2
the rate of formation of H2 O = 2×rate for N2 ;
because 2 moles H2 O formed with 1 mole N 2 /OWTTE;
2
[13]
(i)
the power of a reactant’s concentration in the rate equation/sum of
powers of concentration/rate = k[X]n , where n = order of reaction;
Must be in terms of powers of concentration.
(ii)
(iii)
experiment 1—2 : [X] doubles and rate×4;
2nd order for X;
experiment 2—3 : [Y] doubles and rate×2;
1st order for Y;
rate = k[X]2 [Y](ECF from (ii))
for experiment 1, 1.0×10–2 = k (0.25)2 (0.25);
k = 0.64;
mol–2 dm6 s–1 ;
Allow ECF from rate expression.
1
4
4
21
(iv)
rate = 0.64[0.40]2 [0.60];
= 0.061;
Final answer to 2 sig figs only.
Allow ECF from (iii).
2
[11]
23.
(a)
1/first order;
rate is (directly) proportional to concentration of oxygen/OWTTE;
2
(b)
rate
[NO(g)]
correct axes;
correct shape curve;
(c)
3/third order;
Allow ECF from (a) and (b).
(d)
overall effect on rate = 4× 12 /doubled/×2;
[NO(g)] doubled, rate =×4/quadrupled;
[O2 (g)] halved, rate =×1/halved;
Allow ECF from (a) and (b).
(e)
2
1
3
rate = k[NO(g)]2 [O2 (g)];
k=
rate
6.3 10 −4
;
=
[ NO ( g )]2 [O2 ( g )] (3.0 10 − 2 ) 2 (1.0  10 − 2 )
= 70;
mol–2 dm6 s–1 ;
Allow ECF.
State symbols not needed.
4
[12]
24.
(i)
2ICl + H2 → I2 + 2HCl;
(ii)
ICl order
1;
because doubling [ICl] doubles rate (when [H 2 ] constant);
H2 order
2;
because halving [H2 ] quarters rate (when [ICl] constant);
or doubling [H2 ] quadruples rate (when [ICl] constant);
1
4
22
(iii)
(iv)
(v)
rate = k [ICl][H2 ]2 ;
ECF from (ii).
1
k = 5.00×10−3 ÷0.100×0.05002
mol−2 dm6 s−1 ;
ECF from (iii).
= 20;
2
rate = 20×0.200×0.1002 = 4.00×10−2 (mol dm−3 s−1 );
ECF from (iii).
1
[9]
25.
(a)
(b)
(i)
it relates to the geometric requirements of the reaction/orientation
of reactants on collision/OWTTE;
(ii)
minimum energy needed for reactants to react (on collision)/OWTTE;
(iii)
k measured at different values of temperature;
graph plotted of ln k against 1/T;
intercept on y-axis is ln A;
A = eintercept;
measured slope of graph = − Ea/R;
Ea = – R×gradient;
Award [1] each for any five.
(i)
homogeneous catalyst is in same phase as reactants and heterogeneous
catalyst is in different phase from reactants;
(ii)
1
1
5
1
4
Ea
Ecat
CH2CH 2 + Cl 2
H
CH 2ClCH 2Cl
OR
23
Ea
H
Ecat
CH2CH 2 + Cl2
H
CH 2ClCH2 Cl
reactants line higher than product line (labels not needed);
ΔH label;
Ea label;
Ecat label;
[12]
26.
(a)
(b)
(c)
(d)
order of NO: second/2 - [NO] doubled, rate×4/OWTTE;
order of Br2 : first/1 - as [Br2 ] doubled, rate of reaction doubled/OWTTE;
Reason needed for each mark.
2
rate = k [NO]2 [Br2];
Allow ECF from (a).
1
3.20×10−3 = k(2.00×10−2 )2 ×5.00×10−3
k = 1.60×103 ;
dm6 mol−2 s−1 ;
Allow ECF from (b).
(i)
no effect/K changes only with temperature/OWTTE;
(ii)
decrease (by a factor of 2);
2
1
1
24
(e)
Ea
E a (cat)
E a without catalyst
HR/E R
Enthalpy of
reactants
E a with catalyst
HP/EP
Enthalpy of
products
Time
curve clearly showing Ea without catalyst (Ea);
curve clearly showing Ea with catalyst (Ea(cat));
labelling for x axis;
Accept time/progress of reaction/course of reaction/OWTTE.
Award [2 max] if an enthalpy level diagram for an endothermic
reaction has been correctly drawn.
3
[10]
27.
(i)
+
rate = k[CH3 COCH3 ][H ];
rate = k[CH3 COCH3 ]2;
rate = k[H+]2 ;
(ii)
3
[CH3 COCH3 ] doubles, rate doubles and [H+] doubles, rate doubles;
[Br2 ] double, no effect on rate;
OR
[CH3 COCH3 ] doubles, rate quadruples;
[Br2 ] doubles/[H+] doubles, no effect on rate;
OR
25
[H+] doubles, rate quadruples;
[Br2 ] doubles/[CH3 COCH3 ] doubles, no effect on rate;
The answer given must correspond to the selected expression in (i).
(iii)
(iv)
constant half-life;
at least two sets of data to justify statement;
e.g. [ ] from 1.6 to 0.8 mol dm−3 , 10s; 0.8 to 0.4, 10s; 0.4 to 0.2,
10s.
decrease in the colour of the bromine/OWTTE;
catalyst;
increases rate/speeds up reaction;
by lowering Ea/activation energy (by providing an alternate pathway);
2
2
4
[11]
26
IB Chemistry – SL
Topic 7 Questions
1.
I2 (g) + 3Cl2 (g)
2ICl3 (g)
What is the equilibrium constant expression for the reaction above?
A.
Kc =
[ICl 3 ]
[I 2 ][Cl 2 ]
B.
Kc =
2[ICl 3 ]
3[I 2 ][Cl 2 ]
C.
Kc =
2[ICl 3 ]
[I 2 ] + 3[Cl 2 ]
D.
Kc =
2.
2[ICl 3 ] 2
[I 2 ][Cl 2 ]3
2SO2 (g) + O2 (g)
2SO3 (g)
∆Hο = –200 kJ
According to the above information, what temperature and pressure conditions produce the
greatest amount of SO3 ?
3.
Temperature
Pressure
A.
low
low
B.
low
high
C.
high
high
D.
high
low
Which statement(s) is/are true for a mixture of ice and water at equilibrium?
I.
The rates of melting and freezing are equal.
II.
The amounts of ice and water are equal.
III.
The same position of equilibrium can be reached by cooling water and heating ice.
A.
I only
B.
I and III only
C.
II only
D.
III only
1
4.
What will happen to the position of equilibrium and the value of the equilibrium constant when
the temperature is increased in the following reaction?
Br2 (g) + Cl2 (g)
5.
6.
∆H = +14 kJ
2BrCl(g)
Position of equilibrium
Value of equilibrium constant
A.
Shifts towards the reactants
Decreases
B.
Shifts towards the reactants
Increases
C.
Shifts towards the products
Decreases
D.
Shifts towards the products
Increases
Which statement concerning a chemical reaction at equilibrium is not correct?
A.
The concentrations of reactants and products remain constant.
B.
Equilibrium can be approached from both directions.
C.
The rate of the forward reaction equals the rate of the reverse reaction.
D.
All reaction stops.
In the reaction below
N2 (g) + 3H2 (g)
2NH3 (g)
∆H = –92 kJ
which of the following changes will increase the amount of ammonia at equilibrium?
7.
I.
Increasing the pressure
II.
Increasing the temperature
III.
Adding a catalyst
A.
I only
B.
II only
C.
I and II only
D.
II and III only
In the Haber process for the synthesis of ammonia, what effects does the catalyst have?
A.
B.
C.
D.
Rate of formation of NH 3 (g)
Increases
Increases
Increases
No change
Amount of NH 3 (g) formed
Increases
Decreases
No change
Increases
2
8.
What will happen if CO2 (g) is allowed to escape from the following reaction mixture at
equilibrium?
H+(aq) + HCO3 – (aq)
CO2 (g) + H2 O(l)
9.
10.
A.
The pH will decrease.
B.
The pH will increase.
C.
The pH will remain constant.
D.
The pH will become zero.
Which statements are correct for a reaction at equilibrium?
I.
The forward and reverse reactions both continue.
II.
The rates of the forward and reverse reactions are equal.
III.
The concentrations of reactants and products are equal.
A.
I and II only
B.
I and III only
C.
II and III only
D.
I, II and III
The manufacture of sulfur trioxide can be represented by the equation below.
2SO2 (g) + O2 (g)
2SO3 (g)
∆Hο = –197 kJ mol–1
What happens when a catalyst is added to an equilibrium mixture from this reaction?
11.
A.
The rate of the forward reaction increases and that of the reverse reaction decreases.
B.
The rates of both forward and reverse reactions increase.
C.
The value of ∆Hο increases.
D.
The yield of sulfur trioxide increases.
Which changes will shift the position of equilibrium to the right in the following reaction?
2CO2 (g)
2CO(g) +O2 (g)
I.
adding a catalyst
II.
decreasing the oxygen concentration
III.
increasing the volume of the container
A.
I and II only
B.
I and III only
C.
II and III only
D.
I, II and III
3
12.
13.
Which statement is always true for a chemical reaction that has reached equilibrium?
A.
The yield of product(s) is greater than 50.
B.
The rate of the forward reaction is greater than the rate of the reverse reaction.
C.
The amounts of reactants and products do not change.
D.
Both forward and reverse reactions have stopped.
The equation for a reversible reaction used in industry to convert methane to hydrogen is shown
below.
CO(g) + 3H2 (g) HӨ = +210 kJ
CH4 (g) + H2 O(g)
Which statement is always correct about this reaction when equilibrium has been reached?
14.
A.
The concentrations of methane and carbon monoxide are equal.
B.
The rate of the forward reaction is greater than the rate of the reverse reaction.
C.
The amount of hydrogen is three times the amount of methane.
D.
The value of HӨ for the reverse reaction is –210 kJ.
The equation for a reaction used in the manufacture of nitric acid is
4NH3 (g) + 5O2 (g)
4NO(g) + 6H2 O(g)
HӨ = –900 kJ
Which changes occur when the temperature of the reaction is increased?
15.
Position of equilibrium
Value of Kc
A.
shifts to the left
increases
B.
shifts to the left
decreases
C.
shifts to the right
increases
D.
shifts to the right
decreases
Which changes cause an increase in the equilibrium yield of SO 3 (g) in this reaction?
2SO2 (g) + O2 (g)
I.
increasing the pressure
II.
decreasing the temperature
III.
adding oxygen
A.
I and II only
B.
I and III only
C.
II and III only
D.
I, II and III
2SO3 (g)
HӨ = –196 kJ
4
16.
Iron(III) ions react with thiocyanate ions as follows.
Fe3+(aq) + CNS– (aq)
Fe(CNS)2+(aq)
What are the units of the equilibrium constant, Kc, for the reaction?
17.
A.
mol dm–3
B.
mol2 dm–6
C.
mol–1 dm3
D.
mol–2 dm6
Consider the following equilibrium reaction in a closed container at 350C.
SO2 (g) + Cl2 (g)
SO2 Cl2 (g)
HӨ = −85 kJ
Which statement is correct?
18.
19.
A.
Decreasing the temperature will increase the amount of SO 2 Cl2 (g).
B.
Increasing the volume of the container will increase the amount of SO 2 Cl2 (g).
C.
Increasing the temperature will increase the amount of SO 2 Cl2 (g).
D.
Adding a catalyst will increase the amount of SO 2 Cl2 (g).
Which of the following equilibria would not be affected by pressure changes at constant
temperature?
A.
4HCl(g) + O2 (g)
2H2 O(g) + 2Cl2 (g)
B.
CO(g) + H2 O(g)
H2 (g) + CO2 (g)
C.
C2 H4 (g) + H2 O(g)
D.
PF3 Cl2 (g)
C2 H5 OH(g)
PF3 (g) + Cl2 (g)
Consider the following equilibrium reaction in a closed container at 350°C
SO2 (g) + Cl2 (g)
SO2 Cl2 (g)
HӨ = −85 kJ
Which statement is correct?
A.
Decreasing the temperature will increase the amount of SO 2 Cl2 (g).
B.
Increasing the volume of the container will increase the amount of SO 2 Cl2 (g).
C.
Increasing the temperature will increase the amount of SO 2 Cl2 (g).
D.
Adding a catalyst will increase the amount of SO2 Cl2 (g).
5
20.
What is the equilibrium constant expression, Kc, for the reaction below?
N2 (g) + 2O2 (g)
21.
A.
Kc =
NO 2 
N 2 O 2 
B.
Kc =
2NO 2 
3N 2 O 2 
C.
Kc =
NO 2 2
N 2 O 2 2
D.
Kc =
NO 2 2
N 2  + O 2 2
2NO2 (g)
Sulfur dioxide and oxygen react to form sulfur trioxide according to the equilibrium.
2SO2 (g) + O2 (g)
2SO3 (g)
How is the amount of SO 2 and the value of the equilibrium constant for the reaction affected by
an increase in pressure?
22.
A.
The amount of SO3 and the value of the equilibrium constant both increase.
B.
The amount of SO3 and the value of the equilibrium constant both decrease.
C.
The amount of SO3 increases but the value of the equilibrium constant decreases.
D.
The amount of SO3 increases but the value of the equilibrium constant does not change.
The equation for the Haber process is:
N2 (g) + 3H2 (g)
2NH3 (g)
HӨ = −92.2 kJ
Which conditions will favour the production of the greatest amount of ammonia at equilibrium?
A.
High temperature and high pressure
B.
High temperature and low pressure
C.
Low temperature and high pressure
D.
Low temperature and low pressure
6
23.
The sequence of diagrams represents the system as time passes for a gas phase reaction in which
reactant X is converted to product Y.
Diagram 1
t = 7 seconds
Diagram 2
t = 5 minutes
Diagram
t = 10 minutes
Diagram 4
t = 5 days
Time, t
X=
Y=
Which statement is correct?
24.
A.
At t = 5 days the rate of the forward reaction is greater than the rate of the backward
reaction.
B.
At t = 7 seconds the reaction has reached completion.
C.
At t = 10 minutes the system has reached a state of equilibrium.
D.
At t = 5 days the rate of the forward reaction is less than the rate of the backward
reaction.
What changes occur when the temperature is increased in the following reaction at equilibrium?
Br2 (g) + Cl2 (g)
A.
B.
C.
D.
Position of equilibrium
Shifts towards the reactants
Shifts towards the reactants
Shifts towards the products
Shifts towards the products
2BrCl(g)
∆Hο = +14 kJ mol–1
Value of equilibrium constant
Decreases
Increases
Decreases
Increases
7
25.
The table below gives information about the percentage yield of ammonia obtained in the Haber
process under different conditions.
Pressure/
atmosphere
10
100
200
300
400
600
(a)
Temperature/°C
200
50.7
81.7
89.1
89.9
94.6
95.4
300
14.7
52.5
66.7
71.1
79.7
84.2
400
3.9
25.2
38.8
47.1
55.4
65.2
500
1.2
10.6
18.3
24.4
31.9
42.3
From the table, identify which combination of temperature and pressure gives the highest
yield of ammonia.
……………………………………………………………………………………….
(1)
(b)
The equation for the main reaction in the Haber process is
N2 (g) + 3H2 (g)
2NH3 (g)
∆H is negative
Use this information to state and explain the effect on the yield of ammonia of increasing
(i)
pressure: …………………………….………………………………………..
……………………………………………………………..………………….
………………………………………………………………………………..
(2)
(ii)
temperature: ………………………………………………………………….
…………………………………………………………………………….….
………………………………………………………………………………..
………………………………………………………………………………..
(2)
(c)
In practice, typical conditions used in the Haber process are a temperature of 500 °C and
a pressure of 200 atmospheres. Explain why these conditions are used rather than those
that give the highest yield.
……………………………………………………………………………………….
……………………………………………………………………………………….
(2)
(d)
Write the equilibrium constant expression, Kc, for the production of ammonia.
……………………………………………………………………………………….
……………………………………………………………………………………….
(1)
(Total 8 marks)
8
26.
Consider the following equilibrium reaction.
2SO2 (g) + O2 (g)
2SO3 (g)
∆H = –198 kJ
Using Le Chatelier’s Principle, state and explain what will happen to the position of
equilibrium if
(a)
the temperature increases.
.....................................................................................................................................
.....................................................................................................................................
(2)
(b)
the pressure increases.
.....................................................................................................................................
.....................................................................................................................................
(2)
(Total 4 marks)
27.
Ammonia is produced by the Haber process according to the following reaction.
N2 (g) + 3H2 (g)
(a)
2NH3 (g) H is negative
State the equilibrium constant expression for the above reaction.
....................................................................................................................................
....................................................................................................................................
(1)
(b)
Predict, giving a reason, the effect on the position of equilibrium when the pressure in the
reaction vessel is increased.
....................................................................................................................................
....................................................................................................................................
....................................................................................................................................
(2)
(c)
State and explain the effect on the value of Kc when the temperature is increased.
....................................................................................................................................
....................................................................................................................................
....................................................................................................................................
(2)
(d)
Explain why a catalyst has no effect on the position of equilibrium.
....................................................................................................................................
....................................................................................................................................
(1)
(Total 6 marks)
9
28.
(a)
The following equilibrium is established at 1700°C.
CO2 (g) + H2 (g)
H2 O(g) CO(g)
If only carbon dioxide gas and hydrogen gas are present initially, sketch on a graph a
line representing rate against time for (i) the forward reaction and (ii) the reverse reaction
until shortly after equilibrium is established. Explain the shape of each line.
(7)
(b)
Kc for the equilibrium reaction is determined at two different temperatures. At 850°C,
Kc = 1.1 whereas at 1700°C, Kc = 4.9.
On the basis of these Kc values explain whether the reaction is exothermic or
endothermic.
(3)
(Total 10 marks)
29.
The equation for one reversible reaction involving oxides of nitrogen is shown below:
N2 O4 (g)
2NO2 (g) HӨ = +58 kJ
Experimental data for this reaction can be represented on the following graph:
1.0
product
0.8
concentration
/ mol dm –3
0.6
reactant
0.4
0.2
0
(i)
2
4
6
Time / min
8
10
Write an expression for the equilibrium constant, Kc, for the reaction. Explain the
significance of the horizontal parts of the lines on the graph. State what can be deduced
about the magnitude of Kc for the reaction, giving a reason.
(4)
(ii)
Use Le Chatelier’s principle to predict and explain the effect of increasing the
temperature on the position of equilibrium.
(2)
(iii)
Use Le Chatelier’s principle to predict and explain the effect of increasing the pressure on
the position of equilibrium.
(2)
(iv)
State and explain the effects of a catalyst on the forward and reverse reactions, on the
position of equilibrium and on the value of Kc.
(6)
(Total 14 marks)
10
30.
Consider the following reaction in the Contact process for the production of sulfuric acid for
parts (a) to (d) in this question.
2SO2 + O2
(a)
2SO3
Write the equilibrium constant expression for the reaction.
(1)
(b)
(i)
State the catalyst used in this reaction of the Contact process.
(1)
(ii)
State and explain the effect of the catalyst on the value of the equilibrium constant
and on the rate of the reaction.
(4)
(c)
Use the collision theory to explain why increasing the temperature increases the rate of
the reaction between sulfur dioxide and oxygen.
(2)
(d)
Using Le Chatelier’s principle state and explain the effect on the position of equilibrium
of
(i)
increasing the pressure at constant temperature.
(2)
(ii)
removing of sulfur trioxide.
(2)
(iii)
using a catalyst.
(2)
(Total 14 marks)
31.
Consider the following reaction in the Contact process for the production of sulfuric acid for
parts (a) to (c) in this question.
2SO2 + O2
(a)
2SO3
Write the equilibrium constant expression for the reaction.
(1)
(b)
(i)
State the catalyst used in this reaction of the Contact process.
(1)
(ii)
State and explain the effect of the catalyst on the value of the equilibrium constant
and on the rate of the reaction.
(4)
(c)
Using Le Chatelier’s principle explain the effect on the position of equilibrium of
(i)
increasing the pressure at constant temperature.
(2)
(ii)
removing sulfur trioxide.
(2)
(Total 10 marks)
11
32.
Many reversible reactions in industry use a catalyst. State and explain the effect of a catalyst on
the position of equilibrium and on the value of Kc.
(Total 4 marks)
33.
The equation for a reaction used in industry is
CH4 (g) + H2 O(g)
3H2 (g) + CO(g)
HӨ = +210 kJ
Deduce the equilibrium constant expression, Kc, for this reaction.
(Total 1 mark)
34.
Consider the following reaction where colourless bromide ions react with colourless hydrogen
peroxide to form a red-brown bromine solution.
2Br– (aq) + H2 O2 (aq) + 2H+(aq)
(a)
Br2 (aq) + 2H2 O(l)
H = negative
Predict and explain the effect on the position of equilibrium when
(i)
a small amount of sodium bromide solution is added.
(2)
(ii)
a small amount of sodium hydroxide solution is added.
(2)
(iii)
a catalyst is added.
(2)
(b)
State and explain the effect on the value of the equilibrium constant when the
temperature of the reaction is increased.
(2)
(c)
State and explain the colour change when hydrochloric acid is added to the reaction
solution at equilibrium.
(3)
(Total 11 marks)
35.
The equation for the exothermic reaction in the Contact process is given below:
2SO2 (g) + O2 (g)
(i)
2SO2 (g)
Write the equilibrium constant expression for the reaction.
(1)
(ii)
State and explain qualitatively the pressure and temperature conditions that will give the
highest yield of sulfur trioxide.
(4)
(iii)
In practice, conditions used commercially in the Contact process are 450°C and 2
atmospheres of pressure. Explain why these conditions are used rather than those that
give the highest yield.
(3)
(iv)
Name a catalyst used in the Contact process. State and explain its effect on the value of
the equilibrium constant.
(3)
(Total 11 marks)
12
36.
In the gaseous state, methane and steam react to form hydrogen and carbon dioxide.
(i)
Write an equation for the endothermic equilibrium reaction. Deduce the equilibrium
expression for the reaction and state its units.
(4)
(ii)
Deduce and explain the conditions of temperature and pressure under which the forward
reaction is favoured.
(4)
(iii)
Explain, at the molecular level, why the reaction is carried out at high pressure in
industry.
(2)
(Total 10 marks)
37.
The diagrams below represent equilibrium mixtures for the reaction Y + X2
XY + X at
350 K and 550 K respectively. Deduce and explain whether the reaction is exothermic or
endothermic.
X =
550 K
350 K
Y =
(Total 2 marks)
38.
The equation for the main reaction in the Haber process is:
N2 (g) + 3H2 (g)
(i)
2NH3 (g)
∆H is negative
Determine the equilibrium constant expression for this reaction.
(1)
(ii)
State and explain the effect on the equilibrium yield of ammonia with increasing the
pressure and the temperature.
(4)
(iii)
In practice, typical conditions used in the Haber process involve a temperature of 500°C
and a pressure of 200 atm. Explain why these conditions are used rather than those that
give the highest yield.
(2)
(iv)
At a certain temperature and pressure, 1.1 dm3 of N2 (g) reacts with 3.3 dm3 of H2 (g).
Calculate the volume of NH 3 (g), that will be produced.
(1)
(v)
Suggest why this reaction is important for humanity.
(1)
(vi)
A chemist claims to have developed a new catalyst for the Haber process, which increases
the yield of ammonia. State the catalyst normally used for the Haber process, and
comment on the claim made by this chemist.
(2)
(Total 11 marks)
13
IB Chemistry – SL
Topic 7 Answers
1.
D
2.
B
3.
B
4.
D
5.
D
6.
A
7.
C
8.
B
9.
A
10.
B
11.
C
12.
C
13.
D
14.
B
15.
D
16.
C
17.
A
18.
B
19.
A
20.
C
21.
D
22.
C
23.
C
24.
D
25.
(a)
200°C 600 atm. (both for [1], units not needed);
allow the “highest pressure and the lowest temperature”
(b)
(i)
yield increases/equilibrium moves to the right/more ammonia;
2
4 (gas) molecules → 2/decrease in volume/fewer molecules on right hand side;
(ii)
yield decreases/equilibrium moves to the left/less ammonia;
exothermic reaction/OWTTE;
1
2
14
(c)
(d)
high pressure expensive/greater cost of operating at high pressure/reinforced
pipes etc. needed;
lower temperature – greater yield, but lowers rate;
Do not award a mark just for the word “compromise”.
Kc =
[NH 3 ] 2
[N 2 ][H 2 ]3
2
(ignore units);
1
[8]
26.
(a)
(b)
(position of) equilibrium shifts to the left/towards reactants; (forward)
reaction is exothermic/∆H is negative/the reverse reaction is
endothermic/OWTTE;
Do not accept “Le Chatelier’s Principle” without some
additional explanation.
2
(position of) equilibrium shifts to the right/towards products;
fewer gas molecules on the right hand side/volume decreases in forward
reaction/OWTTE;
Do not accept “Le Chatelier’s Principle” without some
additional explanation.
2
[4]
27.
[NH 3 ] 2
(a)
(k c ) =
(b)
equilibrium shifts to the right/products;
4 mol → 2 mol of gas/fewer moles of gas on the right/products;
2
Kc decreases;
equilibrium position shifts to the left/reactants/forward reaction is exothermic
/reverse reaction is endothermic;
2
catalyst increases the rate of the forward and backward reactions equally
/lowers the activation energy of both forward and backward reaction equally
/lowers Ea so rate of forward and backward reactions increase;
1
(c)
(d)
;
[H 2 ]3 [ N 2 ]
Do not allow round brackets unless K p is used.
1
[6]
28.
(a)
forward
Rate
reverse
Time
two curves – one labelled “forward” starting up high up y-axis and one labelled
“reverse” starting from zero;
15
curves merge and become horizontal;
No penalty for failing to label axes.
forward reaction:
highest concentration, thus rate high to begin with;
as reaction proceeds, concentrations decrease, so does rate;
(b)
reverse reaction:
zero rate initially/at t = 0 (since no products present);
rate increases as concentration of products increases;
equilibrium established when rate of forward reaction = rate of reverse reaction;
7
(reaction is) endothermic;
Kc increases with (increasing) temperature;
forward reaction favoured/heat used up/OWTTE;
3
[10]
29.
(i)
( K c =)
[NO 2 ] 2
;
[N 2 O 4 ]
(horizontal line) concentration of reactant and product remains constant/equilibrium
reached;
(magnitude of) Kc greater than 1;
Accept 1.6.
(ii)
(iii)
(iv)
product concentration greater than reactant concentration;
4
increased temperature shifts equilibrium position to right;
(forward) reaction is endothermic/absorbs heat;
2
increased pressure shifts equilibrium to left;
fewer (gas) moles/molecules on left;
2
both/forward and reverse rates increased/increase in forward reverse
rates are equal;
activation energy reduced;
position of equilibrium unchanged;
concentration/amount of reactants and products remain constant;
value of Kc unchanged;
Kc only affected by changes in temperature;
6
[14]
30.
(a)
K/Kc = [SO3 ]2 ÷[SO2 ]2[O2];
Accept correct Kp expression.
1
(b)
(i)
vanadium(V) oxide/(di)vanadium pentaoxide/V 2 O5 ;
Allow just vanadium oxide but not correct formula.
1
(ii)
catalyst does not affect the value of Kc;
forward and reverse rates increase equally/by the same factor;
catalyst increases the rate of the reaction;
(by providing an alternative path for the reaction with) lower
activation energy;
(c)
more energetic collisions/more molecules have energy greater than
activation energy;
more frequent collisions;
4
2
16
(d)
(i)
(ii)
(iii)
shifts equilibrium position to the products/right;
to the side with fewer gas molecules or moles/lower volume of gas;
2
shifts equilibrium position to the products/right;
to compensate for loss of SO 3 /produce more SO3 ;
2
no effect;
forward and backward rates increased equally/by the same factor;
2
[14]
31.
(a)
K / Kc = [SO3 ]2 ÷[SO2 ]2 [O2 ];
Exactly as written.
Accept correct Kp expression.
1
(b)
(i)
vanadium(V) oxide/(di)vanadium pentaoxide/V 2 O5 /Pt;
Allow just vanadium oxide but not incorrect formula.
1
(ii)
catalyst does not affect the value of K c;
forward and reverse rate increase equally/by the same factor;
catalyst increases the rate of the reaction;
(by providing an alternative path for the reaction with) lower
activation energy;
4
shifts equilibrium position to the products/right;
to the side with least gas molecules or moles/lower volume of gas;
2
shifts equilibrium position to the products/right;
to compensate for loss of SO 3 /produce more SO3 ;
2
(c)
(i)
(ii)
[10]
32.
no effect on position of equilibrium;
forward and reverse reactions speeded up equally/affects the rate of reaction
but not the extent of the reaction;
no effect on value of Kc;
no change in concentrations of reactants or products/Kc only changes if
temperature alters;
[4]
33.
kc =
[H 2 ]3 [CO]
;
[CH 4 ][H 2 O]
[1]
34.
(a)
(i)
(ii)
shifts to the right/toward products/forward reaction favoured;
to consume excess Br− added;
Do not accept “due to Le Chatelier’s principle”.
2
shifts to the left/toward reactants/reverse reaction favoured;
NaOH reacts to consume H +/an increase in the amount of H 2 O
resulting from neutralization;
2
(iii) no effect;
catalyst increases the rate of the forward and backward reactions
equally/lowers the activation energy of both forward and backward
reaction equally/lowers EA so rate of forward and backward reactions
increase equally;
2
17
(b)
(c)
equilibrium constant decreases;
forward reaction is exothermic/produces heat/reverse reaction is endothermic
/absorbs heat;
2
colour change from red-brown to darker red-brown of Br2 /red-brown colour
intensifies/OWTTE;
equilibrium position shifts to the right/products;
to consume H+;
3
[11]
35.
[SO 3 ] 2
(i)
KC =
(ii)
pressure
high pressure (will allow system to occupy smaller volume);
[SO 2 ] 2 [O 2 ]
1
;
Vproduct <Vreactant/equilibrium moves to the right to reduce pressure
/reaction proceeds to lower/lowest number of gaseous molecules
/OWTTE;
Temperature
low temperature;
(exothermic reaction) forward reaction favoured to replace some
of the heat removed/equilibrium moves to the right to produce heat
/OWTTE;
No mark for just saying “due to Le Chatelier's principle”
(iii)
(iv)
4
rate is faster at 450C (than at low temperatures);
>95%/90 − 99% yield/(very) high conversion takes place;
unnecessary to use expensive high pressure equipment/(to achieve) high
pressure is very expensive;
3
vanadium pentoxide/vanadium(V) oxide/V 2 O5 /finely divided
platinum/Pt;
no effect on Kc;
forward and reverse rates speeded up (equally);
3
[11]
36.
(i)
CH4 (g) + 2H2 O(g)
4H2 (g) + CO2 (g);
States not required. Award [1] for balanced equation and [1]
for equilibrium sign.
Kc =
[H 2 ] 4 [CO 2 ]
[CH 4 ][H 2 O] 2
;
units: mol2 dm−6 /mol2 L−2 /mol2 l−2 ; do not accept: M 2
(ii)
4
(endothermic reaction) increase in temperature (favours the forward
reaction);
absorbs (some of) the heat supplied/OWTTE;
Award no marks for saying: “because of Le Chatelier’s
principle”.
low pressure (will allow system to occupy more volume);
18
Vproduct > Vreactant/reaction proceeds to greater number of gaseous moles
/molecules/more moles of gases on right/OWTTE;
ECF from (i)
(iii)
4
at high pressure concentration increases/reaction rate faster;
more frequent collisions;
2
[10]
37.
less product is present at higher temperatures;
Therefore the forward reaction is exothermic;
2
[2]
38.
[NH 3 ] 2
(i)
(Kc =)
(ii)
Increasing the pressure:
Yield increases/equilibrium moves to the right/more ammonia;
4 gas molecules → 2/decrease in volume/fewer gas molecules
on right hand side;
[N 2 ][H 2 ]3
(ignore units);
Increasing the temperature:
Yield decreases/equilibrium moves to the left/less ammonia;
Exothermic reaction/OWTTE;
(iii)
(iv)
Higher temperature increases rate;
Lower pressure is less expensive/lower cost of operating at low
pressure/reinforced pipes not needed;
Do not award a mark just for the word “compromise”.
2.2 (dm3 );
1
4
2
1
Penalize incorrect units.
(v)
(vi)
Fertilizers/increasing crop yields;
Production of explosives for mining;
max
1
Fe/iron;
Allow magnetite/iron oxide.
Claim is not valid since catalysts do not alter the yield/position
of equilibrium/only increase the rate of reaction;
2
[11]
19
IB Chemistry – HL
Topic 7 Questions
1.
For the reaction below
H2 (g) + I2 (g)
2HI(g)
at a certain temperature, the equilibrium concentrations are (in mol dm–3 )
[H2 ] = 0.30, [I2 ] = 0.30, [HI] = 3.0
What is the value of K?
A.
2.
5.0
B.
10
C.
15
D.
100
The value of the equilibrium constant for the reaction
2HI(g)
H2 (g) + I2 (g)
is 0.25 at 440°C. What would the value of the equilibrium constant be for the following reaction
at the same temperature?
H2 (g) + I2 (g)
3.
2HI(g)
A.
0.25
B.
0.50
C.
2.0
D.
4.0
Hydrogen and carbon dioxide react as shown in the equation below.
H2 (g) + CO2 (g)
H2 O(g) + CO(g)
For this reaction the values of Kc with different temperatures are
Temperature / K
Kc
500
7.76×10–3
700
1.23×10–1
900
6.01×10–1
Which statement for the reaction is correct?
A.
The forward reaction is endothermic.
B.
H2 O(g) and CO(g) are more stable than H 2 (g) and CO2 (g).
C.
The reaction goes almost to completion at high temperatures.
D.
The reverse reaction is favoured by high temperatures.
1
4.
The expression for the equilibrium constant for a reaction is
Kc =
BC
A 2
At a certain temperature the values of [A], [B] and [C] are all 0.2 mol dm–3 . What happens to
the value of Kc when all three values are doubled to 0.4 mol dm–3 ?
5.
A.
It is halved.
B.
It does not change.
C.
It doubles.
D.
It increases by a factor of four.
A 1.0 dm3 reaction vessel initially contains 6.0 mol of P and 6.0 mol of Q. At equilibrium 4.0
mol of R is present. What is the value of Kc for the following reaction?
P(g) + Q(g)
6.
A.
0.11
B.
0.25
C.
0.44
D.
4.00
R(g) + S(g)
For the reaction below:
H2 (g) + I2 (g)
2HI(g)
at a certain temperature, the equilibrium concentrations, in mol dm–3 , are
[H2 (g)] = 0.30, [I2 (g)] = 0.30, [HI(g)] = 3.0
What is the value of K?
7.
A.
1.0×10–2
B.
10
C.
33
D.
1.0×102
A liquid and its vapour are at equilibrium inside a sealed container. Which change will alter the
equilibrium vapour pressure of the liquid in the container?
A.
Adding more liquid
B.
Adding more vapour
C.
Decreasing the volume of the container
D.
Decreasing the temperature
2
8.
The equilibrium between nitrogen dioxide (dark brown) and dinitrogen tetroxide (colourless) is
represented by the following equation.
2NO2 (g)
(a)
N2 O4 (g)
∆H = negative Kc = 1 at 328K
Write the equilibrium constant expression, Kc.
.....................................................................................................................................
(1)
(b)
State and explain the effect of an increase in temperature on the value of Kc.
.....................................................................................................................................
.....................................................................................................................................
.....................................................................................................................................
(2)
(c)
State and explain the visible change that takes place as a result of a decrease in pressure,
after equilibrium is re-established.
.....................................................................................................................................
.....................................................................................................................................
.....................................................................................................................................
(2)
(d)
Two moles of NO2 (g) and two moles of N2 O4 (g) were placed in an empty 1 dm3
container and allowed to come to equilibrium at 328 K. Predict, with reference to
the value of Kc, whether the equilibrium mixture would contain more or less than two
moles of NO2 (g).
.....................................................................................................................................
.....................................................................................................................................
(2)
(Total 7 marks)
9.
(a)
The equation for the decomposition of hydrogen iodide is
2HI(g)
H2 (g) + I2 (g)
HӨ = +52 kJ
Predict and explain the effect on the position of equilibrium of
(i)
increasing the pressure, at constant temperature.
(2)
(ii)
increasing the temperature, at constant pressure.
(2)
(iii)
adding a catalyst, at constant temperature and pressure.
(2)
(b)
Deduce the expression for Kc for the forward reaction.
(1)
3
(c)
The equilibrium formed during this reaction was investigated in two experiments carried
out at different temperatures. The results are shown in the table below.
Experiment
(i)
Initial concentration / mol
dm–3
number
[HI]
[H2 ]
[I2 ]
1
0.06
0.00
0.00
2
0.00
0.04
0.04
Equilibrium concentration
/ mol dm–3
[HI]
[H2 ]
[I2 ]
0.01
0.04
For each experiment, deduce the concentrations of the other species present at
equilibrium. Calculate the values of Kc for the forward reaction for each
experiment.
(6)
(ii)
Use the two calculated values of Kc to deduce which of the two experiments was
carried out at the higher temperature, and explain your choice. (If you were not
able to calculate the values of Kc in (c)(i), assume that the values are 0.1 for
experiment 1 and 0.2 for experiment 2, although these are not the correct values.)
(2)
(Total 15 marks)
10.
The equation for another reaction used in industry is
CO(g) + H2 O(g)
(i)
H2 (g) + CO2 (g)
HӨ = −42 kJ
Under certain conditions of temperature and pressure, 2.0 mol of carbon monoxide and
3.2 mol of steam were left to reach equilibrium. At equilibrium, 1.6 mol of both hydrogen
and carbon dioxide were present.
Calculate the amounts of carbon monoxide and steam at equilibrium and the value of Kc.
(3)
(ii)
Under the same conditions of temperature and pressure, 2.0 mol of carbon monoxide and
2.0 mol of steam were left to reach equilibrium.
Calculate the amounts of each reactant and product at equilibrium.
(If you were unable to calculate a value for Kc in (i) use the value 9.0, although this is not
the correct value.)
(2)
(Total 5 marks)
4
IB Chemistry – HL
Topic 7 Answers
1.
D
2.
3.
D
A
4.
B
5.
D
6.
D
7.
8.
D
No ECF throughout this question.
[N 2 O 4 ]
(a)
Kc =
(b)
Kc decreases;
forward reaction is exothermic/∆H is negative/equilibrium
moves to left/OWTTE;
(c)
1
[NO 2 ] 2
2
(mixture will get) darker/darker than expected;
equilibrium position moves to the left/towards reactants as there is an
increase in the number of moles of gas from right to left;
(d)
2
(equilibrium mixture contains) less (than 2 moles NO 2 );
given values make
[N 2 O 4 ]
[NO 2 ] 2
=
1
2
i.e. too much NO2 /OWTTE;
2
[7]
9.
(a)
(i)
(ii)
(iii)
(b)
Kc =
no effect;
equal gas moles on each side;
2
shift to right;
forward reaction absorbs heat/endothermic/OWTTE;
2
no effect;
catalyst speeds up both forward and reverse reactions equally;
2
[H 2 ][I 2 ]
1
[HI] 2
;
Ignore state symbols.
(c)
(i)
experiment 1
[HI] = 0.04 (mol dm−3 );
[I2 ] = 0.01 (mol dm−3 );
5
Kc =
(0.01) 2
= 6.25 10 − 2 ;
(0.04 ) 2
ECF from above values.
experiment 2
[H2 ] = 0.02 (mol dm−3 );
[I2 ] = 0.02 (mol dm−3 );
(0.02 ) 2
= 0.25;
(0.04 ) 2
ECF from above values.
Kc =
(ii)
6
experiment 2 (at higher temperature);
higher Kc value/equilibrium shifted to right;
2
[15]
10.
(i)
CO = 0.4 (mol);
H2 O = 1.6 (mol);
Kc (= 1.62 ÷0.4×1.6) = 4.0/4;
Apply ECF from Kc expression.
3
Ignore units.
(ii)
H2 and CO2 /products = 1.33/1.3 (mol);
CO and H2 O/reactants = 0.67/0.7 (mol);
Using Kc = 9.0, values for H2 and CO2 are 1.5 and values for
CO and H2 O are 0.5.
2
[5]
6
IB Chemistry – SL
Topic 8 Questions
1.
2.
3.
An aqueous solution of which of the following reacts with magnesium metal?
A.
Ammonia
B.
Hydrogen chloride
C.
Potassium hydroxide
D.
Sodium hydrogencarbonate
Which of the following is/are formed when a metal oxide reacts with a dilute acid?
I.
A metal salt
II.
Water
III.
Hydrogen gas
A.
I only
B.
I and II only
C.
II and III only
D.
I, II and III
Four aqueous solutions, I, II, III and IV, are listed below.
I.
0.100 mol dm–3 HCl
II.
0.010 mol dm–3 HCl
III.
0.100 mol dm–3 NaOH
IV.
0.010 mol dm–3 NaOH
What is the correct order of increasing pH of these solutions?
4.
A.
I, II, III, IV
B.
I, II, IV, III
C.
II, I, III, IV
D.
II, I, IV, III
Which substance can be dissolved in water to give a 0.1 mol dm–3 solution with a high pH and a
high electrical conductivity?
A.
HCl
B.
NaCl
C.
NH3
D.
NaOH
1
5.
6.
7.
8.
9.
The pH of a solution is 2. If its pH is increased to 6, how many times greater is the [H +] of the
original solution?
A.
3
B.
4
C.
1000
D.
10 000
The pH of solution X is 1 and that of Y is 2. Which statement is correct about the hydrogen ion
concentrations in the two solutions?
A.
[H+] in X is half that in Y.
B.
[H+] in X is twice that in Y.
C.
[H+] in X is one tenth of that in Y.
D.
[H+] in X is ten times that in Y.
Lime was added to a sample of soil and the pH changed from 4 to 6. What was the
corresponding change in the hydrogen ion concentration?
A.
increased by a factor of 2
B.
increased by a factor of 100
C.
decreased by a factor of 2
D.
decreased by a factor of 100
When the following 1.0 mol dm–3 solutions are listed in increasing order of pH (lowest first),
what is the correct order?
A.
HNO3  H2 CO3  NH3  Ba(OH)2
B.
NH3  Ba (OH)2  H2 CO3  HNO3
C.
Ba (OH)2  H2 CO3  NH3  HNO3
D.
HNO3  H2 CO3  Ba (OH)2  NH3
Which change in [H+] causes the biggest increase in pH?
A.
A change in [H+(aq)] from 1×10–3 to 1×10–2 mol dm–3
B.
A change in [H+(aq)] from 1×10–3 to 1×10–4 mol dm–3
C.
A change in [H+(aq)] from 1×10–4 to 1×10–2 mol dm–3
D.
A change in [H+(aq)] from 1×10–4 to 1×10–6 mol dm–3
2
10.
Which methods can distinguish between solutions of a strong monoprotic acid and a weak
monoprotic acid of the same concentration?
I.
II.
III.
11.
12.
13.
14.
Add magnesium to each solution and measure the rate of the formation of gas
bubbles.
Add aqueous sodium hydroxide to each solution and measure the temperature
change.
Use each solution in a circuit with a battery and lamp and see how bright the lamp
glows.
A.
I and II only
B.
I and III only
C.
II and III only
D.
I, II and III
Which species are a conjugate pair according to the Brønsted-Lowry theory?
A.
CH3 COOH and CH3 CHO
B.
NH3 and BF3
C.
H2 NO3 + and NO3 –
D.
H2 SO4 and HSO4 –
Which is not a strong acid?
A.
Nitric acid
B.
Sulfuric acid
C.
Carbonic acid
D.
Hydrochloric acid
Lime is added to a lake to neutralize the effects of acid rain. The pH value of the lake water
rises from 4 to 7. What is the change in concentration of H + ions in the lake water?
A.
An increase by a factor of 3
B.
An increase by a factor of 1000
C.
A decrease by a factor of 3
D.
A decrease by a factor of 1000
Which is a Brønsted-Lowry acid-base pair?
A.
H2 O and O2–
B.
CH3 COOH and CH3 COO–
C.
NH4 + and NH2 –
D.
H2 SO4 and SO4 2–
3
15.
16.
17.
18.
19.
Solutions of hydrochloric acid (HCl(aq)) and ethanoic acid (CH 3 COOH(aq)) of the same
concentration reacted completely with 5.0 g of calcium carbonate in separate containers. Which
statement is correct?
A.
CH3 COOH(aq) reacted slower because it has a lower pH than HCl(aq).
B.
A smaller volume of CO2 (g) was produced with CH3 COOH(aq) than with HCl(aq).
C.
A greater volume of CO2 (g) was produced with CH3 COOH(aq) than with HCl(aq).
D.
The same volume of CO2 (g) was produced with both CH 3 COOH(aq) and HCl(aq).
Solutions of hydrochloric acid (HCl(aq)) and ethanoic acid (CH 3 COOH(aq)) of the same
concentration reacted completely with 5.0 g of calcium carbonate in separate containers. Which
statement is correct?
A.
CH3 COOH(aq) reacted slower because it has a lower pH than HCl(aq).
B.
A smaller volume of CO2 (g) was produced with CH3 COOH(aq) than with HCl(aq).
C.
A greater volume of CO2 (g) was produced with CH3 COOH(aq) than with HCl(aq).
D.
The same volume of CO2 (g) was produced with both CH 3 COOH(aq) and HCl(aq).
Which acids are strong?
I.
HCl(aq)
II.
HNO3 (aq)
III.
H2 SO4 (aq)
A.
I and II only
B.
I and III only
C.
II and III only
D.
I, II and III
The pH of a solution changes from pH = 1 to pH = 3. What happens to the [H +] during this pH
change?
A.
It increases by a factor of 100.
B.
It decreases by a factor of 100.
C.
It increases by a factor of 1000.
D.
It decreases by a factor of 1000.
What is the conjugate base of the HSO 4 – (aq) ion?
A.
H2 SO4 (aq)
B.
SO4 2– (aq)
C.
H2 O(l)
D.
H3 O+(aq)
4
20.
21.
22.
23.
Which species can act as a Lewis acid?
A.
BF3
B.
OH–
C.
H2 O
D.
NH3
Which substance, when dissolved in water, to give a 0.1 mol dm–3 solution, has the highest pH?
A.
HCl
B.
NaCl
C.
NH3
D.
NaOH
Which methods will distinguish between equimolar solutions of a strong base and a strong acid?
I.
Add magnesium to each solution and look for the formation of gas bubbles.
II.
Add aqueous sodium hydroxide to each solution and measure the temperature
change.
III.
Use each solution in a circuit with a battery and lamp and see how bright the lamp
glows.
A.
I and II only
B.
I and III only
C.
II and III only
D.
I, II and III
(a)
Aqueous XO4 3– ions form a precipitate with aqueous silver ions, Ag+. Write a balanced
equation for the reaction, including state symbols.
...........................................................................................................................................
(2)
(b)
When 41.18 cm3 of a solution of aqueous silver ions with a concentration of 0.2040 mol
dm–3 is added to a solution of XO4 3– ions, 1.172 g of the precipitate is formed.
(i)
Calculate the amount (in moles) of Ag+ ions used in the reaction.
(1)
............................................................................................................................
(ii)
Calculate the amount (in moles) of the precipitate formed.
(1)
............................................................................................................................
5
(iii)
Calculate the molar mass of the precipitate.
............................................................................................................................
............................................................................................................................
............................................................................................................................
............................................................................................................................
(2)
(iv)
Determine the relative atomic mass of X and identify the element.
............................................................................................................................
............................................................................................................................
(2)
(Total 8 marks)
24.
(a)
(i)
A solution of hydrochloric acid has a concentration of 0.10 mol dm–3 and a pH
value of 1. The solution is diluted by a factor of 100. Determine the concentration
of the acid and the pH value in the diluted solution.
..........................................................................................................................
..........................................................................................................................
(2)
(ii)
Explain why 0.10 mol dm–3 ethanoic acid solution and the diluted solution in (a) (i)
have similar [H+] values.
..........................................................................................................................
..........................................................................................................................
..........................................................................................................................
..........................................................................................................................
(3)
(b)
Suggest one method, other than measuring pH, which could be used to distinguish
between solutions of a strong acid and a weak acid of the same concentration. State the
expected results.
....................................................................................................................................
....................................................................................................................................
....................................................................................................................................
....................................................................................................................................
....................................................................................................................................
....................................................................................................................................
(2)
(Total 7 marks)
6
25.
Define the terms strong acid and weak acid. Using hydrochloric and ethanoic acid as examples,
write equations to show the dissociation of each acid in aqueous solution.
……………………………………………………………………………………………
……………………………………………………………………………………………
……………………………………………………………………………………………
……………………………………………………………………………………………
……………………………………………………………………………………………
……………………………………………………………………………………………
(Total 4 marks)
26.
(i)
Calcium carbonate is added to separate solutions of hydrochloric acid and ethanoic acid
of the same concentration. State one similarity and one difference in the observations you
could make.
…………………………………………………………………………………………..
…………………………………………………………………………………………..
…………………………………………………………………………………………..
(2)
(ii)
Write an equation for the reaction between hydrochloric acid and calcium carbonate.
…………………………………………………………………………………………..
…………………………………………………………………………………………..
(2)
(iii)
Determine the volume of 1.50 mol dm–3 hydrochloric acid that would react with exactly
1.25 g of calcium carbonate.
…………………………………………………………………………………………..
…………………………………………………………………………………………..
…………………………………………………………………………………………..
…………………………………………………………………………………………..
(3)
(iv)
Calculate the volume of carbon dioxide, measured at 273 K and
1.01×105 Pa, which would be produced when 1.25 g of calcium carbonate reacts
completely with the hydrochloric acid.
…………………………………………………………………………………………..
…………………………………………………………………………………………..
…………………………………………………………………………………………..
(2)
(Total 9 marks)
7
27.
The pH values of solutions of three organic acids of the same concentration were measured.
acid X
acid Y
acid Z
(i)
pH = 5
pH = 2
pH = 3
Identify which solution is the least acidic.
(1)
(ii)
Deduce how the [H+] values compare in solutions of acids Y and Z.
(2)
(iii)
Arrange the solutions of the three acids in decreasing order of electrical conductivity,
starting with the greatest conductivity, giving a reason for your choice.
(2)
(Total 5 marks)
28.
The equilibrium reached when ethanoic acid is added to water can be represented by the
following equation:
CH3 COOH(l) + H2 O(l)
CH3 COO– (aq)+H3 O+(aq)
Define the terms Brønsted-Lowry acid and Lewis base, and identify two examples of each of
these species in the equation.
(Total 4 marks)
29.
Identify one example of a strong acid and one example of a weak acid. Outline three different
methods to distinguish between equimolar solutions of these acids in the laboratory. State how
the results would differ for each acid.
(Total 5 marks)
30.
Vinegar has a pH of approximately 3 and some detergents have a pH of approximately 8. State
and explain which of these has the higher concentration of H + and by what factor.
(Total 1 mark)
31.
Define the terms Brønsted-Lowry acid and Lewis acid. For each type of acid, identify one
example other than water and write an equation to illustrate the definition.
(Total 5 marks)
32.
The pH values of three acidic solutions, X, Y and Z, are shown in the following table:
Solution
(i)
Acid
pH
X
HCl(aq)
2
Y
HCl(aq)
4
Z
CH3 COOH(aq)
4
Solutions X and Z have the same acid concentration. Explain, by reference to both acids,
why they have different pH values.
(2)
(ii)
Deduce by what factor the values of [H +] in solutions X and Y differ.
(1)
(Total 3 marks)
8
33.
State and explain two methods, other than measuring pH, which could be used to distinguish
between 1.0 mol dm–3 solutions of nitric acid and ethanoic acid.
(Total 4 marks)
34.
Propanoic acid is classified as a weak acid.
(a)
State the meaning of the term weak acid.
....................................................................................................................................
....................................................................................................................................
(1)
(b)
State, giving a reason in each case, two methods other than measuring pH, that could be
used to distinguish between 0.100 mol dm–3 propanoic acid and 0.100 mol dm–3 nitric
acid.
....................................................................................................................................
....................................................................................................................................
....................................................................................................................................
....................................................................................................................................
(2)
(Total 3 marks)
35.
State an equation for the reaction of propanoic acid with water. Identify one conjugate
Brønsted-Lowry pair.
..............................................................................................................................................
..............................................................................................................................................
..................................................................................................................................... .........
..............................................................................................................................................
..............................................................................................................................................
..............................................................................................................................................
(Total 2 marks)
9
IB Chemistry – SL
Topic 8 Answers
1.
B
2.
B
3.
B
4.
D
5.
D
6.
D
7.
D
8.
A
9.
D
10.
D
11.
D
12.
C
13.
D
14.
B
15.
D
16.
D
17.
D
18.
B
19.
B
20.
A
21.
D
22.
A
23.
(a)
(b)
3Ag+(aq) + XO4 3– (aq) → Ag3 XO4 (s);
states;
[1] for balanced equation and [1] for states.
(i)
(ii)
nAg+ = cV = 0.2040 mol dm–3 ×0.04118 dm3
= 0.008401/8.401×10–3 mol (–1 SF)
Ignore units even if wrong, do not award mark unless 4 sig fig.
n Ag3XO 4 =
1
3
nAg+ =
1
3
2
1
×0.008401 mol
= 0.002800/2.800×10–3 mol
1
ECF from (a) and (b)(i)
10
(iii)
(iv)
0.002800 mol weighs 1.172 g
1.172 g
1 mol weighs
= 418.6 g mol–1
0.002800 mol
418.6;
Accept answer in range 418 to 419.
No penalty for too many sig figs.
ECF from (b) (ii)
g mol–1
Do not accept g.
2
(3×107.87) + x + 4 (16.0) = 418.6
therefore, x = 30.99 (accept 31.0/31);
P/phosphorous;
2
[8]
24.
(a)
(i)
(ii)
(b)
0.0010 / 1.0×10−3 (mol dm−3 );
pH = 3;
2
HCl: strong acid/fully dissociated;
CH3 COOH : weak acid/partially dissociated;
HCl less concentrated/CH 3 COOH more concentrated;
only one molecule in 100 dissociates in ethanoic acid so [H +]
1/100/OWTTE
3
measure electrical conductivity;
strong acids are good conductors/weak acids are poor conductors;
OR
react with magnesium or a named active metal/(metal) carbonate;
hydrogen carbonate/bicarbonate;
strong acids have a faster reaction/more gas bubbles (per unit time)
/more heat produced/weak acids have a slower reaction/less gas
bubbles (per unit time)/less heat produced;
2
titration curves: namely strong acid and strong base will have an equivalence
point pH of 7 and a weak acid and strong base will have an equivalence point
pH of >7.
OR
temperature change: on neutralization for temperature change: namely,
neutralization (H+ + OH−) is exothermic, weak acid is partially dissociated
so some energy used up in dissociation of weak acid − net result, weak acid
would produce less energy/less temperature increase compared to
neutralization of strong acid.
[7]
25.
strong acid completely dissociated/ionized;
weak acid only partially dissociated/ionized;
HCl(aq) → H+(aq) + Cl+(aq);
CH3 COOH(aq)
CH3 COO– (aq) + H+(aq);
Insist on both arrows as shown, state symbols not needed.
Also accept H2 O(1) and H3 O+(aq) in equations
4
[4]
11
26.
(i)
bubbling/effervescence/dissolving of CaCO 3 /gas given off
(do not accept CO2 produced);
more vigorous reaction with HCl/OWTTE;
(ii)
2HCl(aq) + CaCO3 (s) → CaCl2 (aq) + CO2 (g) + H2 O(1);
[1] for correct formulas, [1] for balanced, state symbols not
essential.
(iii)
amount of CaCO3 = 1.25 (no penalty for use of 100);
100.09
amount of HCl = 2×0.0125 = 0.0250 mol (allow ECF);
volume of HCl = 0.0167 dm3 /16.7 cm3 (allow ECF);
(iv)
2
2
3
1:1 ratio of CaCO3 to CO2 /use 0.0125 moles CO2 (allow ECF);
(0.0125×22.4) = 0.28 dm3 /280 cm3 /2.8×10–4 m3 (allow ECF);
Accept calculation using pV=nRT.
1
[9]
27.
(i)
X;
1
(ii)
greater in Y/smaller in Z;
by a factor of 10;
2
Y > Z > X;
most ions/greatest concentration of ions in Y/OWTTE;
2
(iii)
[5]
28.
Brønsted-Lowry acid
proton donor/OWTTE;
CH3 COOH and H3 O+;
Lewis base
electron pair donor/OWTTE;
H2 O and CH3 COO−;
4
[4]
29.
HCl/H2 SO4 /HNO3 /any strong acid;
CH3 COOH/H2CO3/any weak acid;
Measure pH − the strong acid has the lower pH;
Accept universal indicator and two correct colours.
Measure (electrical) conductivity − this is greater for the stronger acid;
Add magnesium/carbonate − more gas bubbles with the stronger acid/Mg or
carbonate would disappear faster with stronger acid;
[5]
30.
vinegar and factor of 105 ;
[1]
31.
Brønsted-Lowry acid
a proton donor;
Lewis acid
electron pair acceptor;
Brønsted-Lowry acid
Any suitable equation;
12
Lewis acid − BF3 /AlCl3 /transition metal ions that form complex ion
with ligands;
For example
BF3 + NH3 → BF3 NH3 /Cu2+ + 4NH3 → [Cu(NH3)4 ]2+/AlCl3 + Cl− → AlCl 4− ;
Or any suitable equation.
5
[5]
32.
(i)
(ii)
HCl/X is strong and CH3 COOH/Z is weak;
HCl/X is fully dissociated and CH 3 COOH is slightly dissociated;
[H+ ] is greater in HCl/X than in CH 3 COOH/Z;
Any two for [1] each.
a factor of 100;
2
1
[3]
33.
conductivity;
nitric acid will contain more ions and have a higher conductivity/ethanoic acid
will have fewer ions and have a lower conductivity;
rate of reaction with metal/carbonate/hydrogencarbonate;
nitric acid will react more rapidly/produce bubbles faster/ethanoic
acid will react less rapidly/produce bubbles more slowly;
reaction with alkali;
temperature change will be less for ethanoic acid;
Accept any two methods and explanations from above.
4
[4]
34.
(a)
an acid that partially dissociates/ionizes/doesn’t fully dissociate/ionize;
(b)
conductivity - propanoic acid will be lower because lower ion concentration
/less dissociated;
reaction with metal/metal carbonate/metal hydrogencarbonate - propanoic
acid will react slower/less vigorously because lower [H +]/less dissociated;
reaction with alkali - temperature change will be less for propanoic acid
because lower [H+]/less dissociated;
Award [1] mark each for two.
1
2
[3]
35.
CH3 CH2 COOH + H2 O
CH3 CH2COO− + H3 O+/CH3CH2COOH
CH3 CH2 COO− + H+;
required for mark.
CH3 CH2 COOH and CH3 CH2COO−/H3 O+ and H2 O;
2
[2]
13
IB Chemistry – HL
Topic 8 Questions
1.
2.
3.
The Ka value for an acid is 1.0×10–2 . What is the Kb value for its conjugate base?
A.
1.0×10–2
B.
1.0×10–6
C.
1.0×10–10
D.
1.0×10–12
Separate 20.0 cm3 solutions of a weak acid and a strong acid of the same concentration are
titrated with NaOH solution. Which will be the same for these two titrations?
I.
Initial pH
II.
pH at equivalence point
III.
Volume of NaOH required to reach the equivalence point
A.
I only
B.
III only
C.
I and II only
D.
II and III only
When the following 1.0 mol dm–3 aqueous solutions are arranged in order of increasing pH,
which is the correct order?
I.
Ammonium chloride
II.
Ammonium ethanoate
III.
Sodium ethanoate
A.
I, II, III
B.
II, I, III
C.
III, I, II
D.
III, II, I
1
4.
An acid-base indicator, HIn, dissociates according to the following equation.
HIn(aq)
colour A
H+(aq) + In– (aq)
colour B
Which statement about this indicator is correct?
5.
6.
I.
In a strongly acidic solution colour B would be seen.
II.
In a neutral solution the concentrations of HIn(aq) and In– (aq) must be equal.
III.
It is suitable for use in titrations involving weak acids and weak bases.
A.
I only
B.
II only
C.
III only
D.
None of the above
What is the concentration of OH – ions (in mol dm–3 ) in an aqueous solution in which
[H+] = 2.0×10–3 mol dm–3 ? (Kw = 1.0×10–14 mol2 dm–6 )
A.
2.0×10–3
B.
4.0×10–6
C.
5.0×10–12
D.
2.0×10–17
What is the relationship between Ka and pKa?
A.
pKa = –log Ka
B.
−14
pKa = 1.0  10
Ka
C.
pKa = log Ka
D.
pKa = 1.0
Ka
2
7.
Which curve is produced by the titration of a 0.1 mol dm–3 weak base with 0.1 mol dm–3 strong
acid?
12
A.
12
B.
8
8
pH
pH
4
4
Volume of titrant
12
C.
Volume of titrant
12
D.
8
8
pH
pH
4
4
Volume of titrant
8.
9.
10.
Volume of titrant
The acid dissociation constant of a weak acid HA has a value of 1.0×10–5 mol dm–3 .
What is the pH of a 0.10 mol dm–3 aqueous solution of HA?
A.
2
B.
3
C.
5
D.
6
Which mixture would produce a buffer solution when dissolved in 1.0 dm3 of water?
A.
0.50 mol of CH3 COOH and 0.50 mol of NaOH
B.
0.50 mol of CH3 COOH and 0.25 mol of NaOH
C.
0.50 mol of CH3 COOH and 1.00 mol of NaOH
D.
0.50 mol of CH3 COOH and 0.25 mol of Ba(OH)2
Which compound, when dissolved in aqueous solution, has the highest pH?
A.
NaCl
B.
Na2 CO3
C.
NH4 Cl
D.
NH4 NO3
3
11.
Which values are correct for a solution of NaOH of concentration 0.010 mol dm–3 at 298 K?
(Kw = 1.0×10–14 mol2 dm–6 at 298 K)
12.
13.
14.
15.
A.
[H+] = 1.0×10–2 mol dm–3
and
pH = 2.00
B.
[OH– ] = 1.0×10–2 mol dm–3
and
pH = 12.00
C.
[H+] = 1.0×10–12 mol dm–3
and
pOH = 12.00
D.
[OH– ] = 1.0×10–12 mol dm–3
and
pOH = 2.00
Which solution, of concentration 0.10 mol dm–3 , has the highest pH value?
A.
HCl(aq)
B.
MgCl2 (aq)
C.
NaCl(aq)
D.
AlCl3 (aq)
Which statement about indicators is always correct?
A.
The mid-point of an indicator’s colour change is at pH = 7.
B.
The pH range is greater for indicators with higher pK4 values.
C.
The colour red indicates an acidic solution.
D.
The pKa value of the indicator is within its pH range.
Which compound will dissolve in water to give a solution with a pH greater than 7?
A.
sodium chloride
B.
potassium carbonate
C.
ammonium nitrate
D.
lithium sulfate
An aqueous solution has a pH of 10. Which concentrations are correct for the ions below?
[H+(aq)] mol dm–3
[OH– (aq)] mol dm–3
A.
104
10–10
B.
10–4
10–10
C.
10–10
10–4
D.
10–10
10–4
4
16.
Which graph shows how the pH changes when a weak base is added to a strong acid?
pH
14
pH
14
A
B
7
7
0
0
volume of base
volume of base
pH
14
pH
14
C
D
7
7
0
0
volume of base
17.
volume of base
When the following acids are listed in decreasing order of acid strength (strongest first), what is
the correct order?
Ka
18.
benzoic
6.31×10–5
chloroethanoic
1.38×10–3
ethanoic
1.74×10–5
A.
chloroethanoic  benzoic  ethanoic
B.
benzoic  ethanoic  chloroethanoic
C.
chloroethanoic  ethanoic  benzoic
D.
ethanoic  benzoic  chloroethanoic
The strengths of organic acids can be compared using Ka and pKa values. Which acid is the
strongest?
A.
Acid A
pKa = 6
B.
Acid B
pKa = 3
C.
Acid C
Ka = 1×10–5
D.
Acid D
Ka = 1×10–4
5
19.
20.
21.
22.
23.
Which is the correct statement about the pH and pOH values of an aqueous solution at 25°C?
A.
pH + pOH =14.0
B.
pH + pOH =1.0 ×10–14
C.
pH × pOH =14.0
D.
pH × pOH =1.0 ×10–14
Which salt, when dissolved in water to form a 1.0 mol dm–3 solution, produces the lowest pH
value?
A.
Ammonium chloride
B.
Ammonium ethanoate
C.
Sodium ethanoate
D.
Sodium chloride
Which solution has the lowest pH value?
A.
Aluminium sulfate
B.
Sodium nitrate
C.
Potassium chloride
D.
Sodium ethanoate
Which neutralization reaction could use phenolphthalein (pKa = 9.3) and not methyl orange (pKa
= 3.7) as an indicator?
A.
NaOH(aq) and HNO3 (aq)
B.
NH3 (aq) and CH3 COOH(aq)
C.
NaOH(aq) and CH3 COOH(aq)
D.
NH3 (aq) and HNO3 (aq)
Water dissociates according to the equation
H2 O(l)
H+(aq) + OH– (aq)
H = +56 kJ
At 25C water has a pH of 7. Which of the following occurs when water is heated to 30C?
A.
It remains neutral and its pH decreases.
B.
It becomes acidic and its pH decreases.
C.
It remains neutral and its pH increases.
D.
It becomes acidic and its pH increases.
6
24.
25.
Which mixture would produce a buffer solution when dissolved in 1.0 dm3 of water?
A.
0.30 mol of NH3 (aq) and 0.30 mol of HCl(aq)
B.
0.30 mol of NH3 (aq) and 0.15 mol of HCl(aq)
C.
0.30 mol of NH3 (aq) and 0.60 mol of HCl(aq)
D.
0.30 mol of NH3 (aq) and 0.15 mol of H2 SO4 (aq)
Ammonia (NH3 ) is a weak base in aqueous solution with an ionization constant Kb. What
expression is equal to the ionization constant for the following reaction?
NH4 +(aq) + H2 O(l)
26.
A.
Kw
Ka
B.
Ka
Kw
C.
Kw
Kb
D.
Kb
Kw
NH3 (aq) + H3 O+(aq)
The pKa values of four acids are as follows.
W
X
Y
Z
4.87
4.82
4.86
4.85
What is the correct order when these acids are arranged in order of increasing acid strength?
27.
A.
X, Z, Y, W
B.
X, Y, Z, W
C.
W, Z, Y, X
D.
W, Y, Z, X
10 cm3 of 0.01 mol dm–3 nitric acid (HNO3 ) is diluted with 90 cm3 of water. What is the pH of
the resulting solution?
A.
1
B.
2
C.
3
D.
4
7
28.
29.
30.
31.
32.
A base of concentration 0.10 mol dm–3 is titrated with 25 cm3 of an acid of concentration 0.10
mol dm–3 . Which base-acid pair would have the highest pH at the equivalence point?
A.
NaOH(aq) and CH3 COOH(aq)
B.
NaOH(aq) and HNO3 (aq)
C.
NH3 (aq) and HNO3 (aq)
D.
NH3 (aq) and CH3 COOH(aq)
What is the value of [H+] in a buffer solution in which [CH 3 COOH] = 2.0 mol dm–3 and
[CH3 COO– ] 1.0 mol dm–3 ? For CH3 COOH, Ka = 1.8×10–5 mol dm–3 .
A.
6.0×10–3
B.
3.6×10–5
C.
1.8×10–5
D.
9.1×10–6
Which salt forms the most acidic solution when added to water?
A.
NaCl
B.
MgSO4
C.
Al(NO3 )3
D.
KHCO3
An acid-base indicator has a pKa value of 4.0. At what pH will this indicator change colour?
A.
2.0
B.
4.0
C.
8.0
D.
12.0
Which values are correct for a 0.010 mol dm–3 solution of NaOH(aq) at 298 K?
(Kw = 1.0×10–14 mol2 dm–6 at 298 K)
A.
[H+] = 1.0×10–12 mol dm–3 and pH = 12.00
B.
[OH– ] = 1.0×10–12 mol dm–3 and pH = 12.00
C.
[H+] = 1.0×10–12 mol dm–3 and pOH = 12.00
D.
[OH– ] = 1.0×10–12 mol dm–3 and pOH = 12.00
8
33.
34.
35.
At 25°C, Ka for an acid is 1.0×10–2 . What is the value of Kb for its conjugate base?
A.
1.0×102
B.
1.0×10–2
C.
1.0×1012
D.
1.0×10–12
Which statement about indicators is always correct?
A.
The mid-point of the pH range of an indicator is 7.
B.
The pH range is greater for indicators with higher pKa values.
C.
The colour red indicates an acidic solution.
D.
The pKa value of the indicator is within its pH range.
(a)
(i)
Calculate the Ka value of methanoic acid, HCOOH, using table 16 in the Data
Booklet.
........................................................................................................................
........................................................................................................................
(1)
(ii)
Based on its Ka value, state and explain whether methanoic acid is a strong or weak
acid.
...........................................................................................................................
...........................................................................................................................
...........................................................................................................................
...........................................................................................................................
(2)
(iii)
Calculate the hydrogen ion concentration and the pH of a 0.010 mol dm–3
methanoic acid solution. State one assumption made in arriving at your answer.
(4)
...........................................................................................................................
...........................................................................................................................
...........................................................................................................................
...........................................................................................................................
...........................................................................................................................
...........................................................................................................................
...........................................................................................................................
9
(b)
Explain how you would prepare a buffer solution of pH 3.75 starting with methanoic
acid.
......................................................................................................................................
......................................................................................................................................
......................................................................................................................................
......................................................................................................................................
......................................................................................................................................
......................................................................................................................................
(3)
(Total 10 marks)
36.
The indicator bromophenol blue, HIn(aq), has a form that is yellow and an In– (aq) form that
is blue.
(a)
Write an equation to show how bromophenol blue acts as an indicator.
.....................................................................................................................................
(1)
(b)
State and explain the colour of bromophenol blue
(i)
on the addition of a strong acid.
...........................................................................................................................
...........................................................................................................................
(ii)
at the equivalence point of a titration.
...........................................................................................................................
...........................................................................................................................
(3)
(Total 4 marks)
37.
(a)
The dissociation of water takes place as follows:
H2 O(l)
(i)
H+(aq) + OH– (aq)
State the expression for the ionic product constant of water, Kw.
.........................................................................................................................
.........................................................................................................................
(1)
(ii)
The value of Kw is 2.410–14 mol2 dm–6 at 310 K. Calculate the [H +] at 310 K.
.........................................................................................................................
.........................................................................................................................
(1)
10
(b)
Lactic acid CH3 CH(OH)COOH is a weak monoprotic acid
(pKa = 3.85 and Ka = 1.4×10–4 mol dm–3 ).
(i)
Write an equation for the reaction of lactic acid with water.
.........................................................................................................................
.........................................................................................................................
(1)
(ii)
State the ionization constant expression, Ka, for lactic acid.
.........................................................................................................................
.........................................................................................................................
(1)
(iii)
Calculate the pH of a 0.20 mol dm–3 solution of lactic acid.
.........................................................................................................................
.........................................................................................................................
(2)
(iv)
Determine the pH of a solution containing 0.10 mol dm–3 of lactic acid and 0.10
mol dm–3 of sodium lactate.
.........................................................................................................................
.........................................................................................................................
(1)
(Total 7 marks)
38.
(a)
(i)
Write the equation for the reaction of ammonia with water.
.........................................................................................................................
.........................................................................................................................
(1)
(ii)
Derive the expression for Kb for this reaction.
.........................................................................................................................
.........................................................................................................................
(1)
(b)
Using information from Table 16 in the Data Booklet, determine the pOH of a
0.20 mol dm–3 solution of ammonia.
....................................................................................................................................
....................................................................................................................................
....................................................................................................................................
(3)
(Total 5 marks)
11
39.
Benzoic acid, C6 H5 COOH, is a weak acid.
(a)
Deduce the equation for the ionization of benzoic acid in water.
....................................................................................................................................
....................................................................................................................................
(1)
(b)
Use information from Table 16 in the Data Booklet to calculate a value for the
dissociation constant, Ka, for benzoic acid.
....................................................................................................................................
....................................................................................................................................
(1)
(c)
Derive the ionization constant expression for benzoic acid and use it to determine the pH
of a 0.20 mol dm–3 aqueous solution of benzoic acid.
....................................................................................................................................
....................................................................................................................................
....................................................................................................................................
....................................................................................................................................
....................................................................................................................................
....................................................................................................................................
(3)
(Total 5 marks)
40.
The hydrogen ion concentration in pure water varies with temperature. At a particular
temperature [H+] =1.7×10–7 mol dm–3 .
(a)
State the expression for the ionic product constant of water, Kw, and calculate the value of
Kw at this temperature.
....................................................................................................................................
....................................................................................................................................
....................................................................................................................................
....................................................................................................................................
(2)
(b)
Calculate the pH of water at this temperature.
....................................................................................................................................
....................................................................................................................................
(1)
12
(c)
State and explain whether water at this temperature is acidic, neutral or alkaline.
....................................................................................................................................
....................................................................................................................................
(2)
(Total 5 marks)
41.
Predict whether each of the following solutions would be acidic, alkaline or neutral. In each
case explain your reasoning.
(i)
0.1 mol dm–3 FeCl3 (aq)
…………………………………………………………………………………………..
…………………………………………………………………………………………..
(ii)
0.1 mol dm–3 NaNO3 (aq)
…………………………………………………………………………………………..
…………………………………………………………………………………………..
(iii)
0.1 mol dm–3 Na2 CO3 (aq)
…………………………………………………………………………………………..
…………………………………………………………………………………………..
(Total 6 marks)
42.
The following graph shows how the pH changes during the titration of 10 cm3 of a solution of a
weak acid (HA) with 0.10 mol dm–3 NaOH.
14
13
12
11
10
9
8
pH
7
6
5
4
3
2
1
0
0
1
2
3
4
5
6
7
8
Volume of 0.10 mol dm
9 10 11 12 13 14
–3
NaOH / cm 3
13
(i)
State the pH at the equivalence point and explain why the pH changes rapidly in this
region.
…………………………………………………………………………………………..
…………………………………………………………………………………………..
…………………………………………………………………………………………..
(2)
(ii)
Calculate the initial concentration of the acid (HA).
…………………………………………………………………………………………..
…………………………………………………………………………………………..
…………………………………………………………………………………………..
…………………………………………………………………………………………..
…………………………………………………………………………………………..
(3)
(iii)
Calculate the [H+] of the acid before any sodium hydroxide is added. Use this value to
determine the Ka value and the pKa value of the acid.
…………………………………………………………………………………………..
…………………………………………………………………………………………..
…………………………………………………………………………………………..
…………………………………………………………………………………………..
…………………………………………………………………………………………..
…………………………………………………………………………………………..
…………………………………………………………………………………………..
…………………………………………………………………………………………..
(5)
(Total 10 marks)
43.
A buffer solution can be made by dissolving 0.25 g of sodium ethanoate in 200 cm3 of 0.10 mol
dm–3 ethanoic acid. Assume that the change in volume is negligible.
(i)
Define the term buffer solution.
…………………………………………………………………………………………..
…………………………………………………………………………………………..
…………………………………………………………………………………………..
…………………………………………………………………………………………..
(2)
14
(ii)
Calculate the concentration of the sodium ethanoate in mol dm–3 .
…………………………………………………………………………………………..
…………………………………………………………………………………………..
…………………………………………………………………………………………..
…………………………………………………………………………………………..
(3)
(iii)
Calculate the pH of the resulting buffer solution by using information from Table 16 of
the Data Booklet.
…………………………………………………………………………………………..
…………………………………………………………………………………………..
…………………………………………………………………………………………..
…………………………………………………………………………………………..
(3)
(Total 8 marks)
44.
An experiment was carried out to determine the concentration of an aqueous solution of
ammonia by titrating it with a solution of sulfuric acid of concentration 0.150 mol dm–3 . It was
found that 25.0 cm3 of the ammonia solution required 20.1 cm3 of the sulfuric acid solution for
neutralization.
(a)
Write the equation for the reaction and calculate the concentration, in mol dm–3 , of the
ammonia solution.
.....................................................................................................................................
.....................................................................................................................................
.....................................................................................................................................
.....................................................................................................................................
.....................................................................................................................................
.....................................................................................................................................
(4)
(b)
Several acid-base indicators are listed in Table 17 of the Data Booklet. State and explain
which one of the following indicators should be used for this experiment: bromocresol
green, phenol red or phenolphthalein.
.....................................................................................................................................
.....................................................................................................................................
.....................................................................................................................................
.....................................................................................................................................
(3)
15
(c)
Determine the pOH of a solution with an ammonia concentration of 0.121 mol dm–3 .
(pKb of ammonia is 4.75.)
.....................................................................................................................................
.....................................................................................................................................
.....................................................................................................................................
.....................................................................................................................................
.....................................................................................................................................
.....................................................................................................................................
.....................................................................................................................................
.....................................................................................................................................
(4)
(Total 11 marks)
45.
(i)
State what is meant by the term buffer solution, and describe the composition of an acid
buffer solution in general terms.
.....................................................................................................................................
.....................................................................................................................................
.....................................................................................................................................
.....................................................................................................................................
(3)
(ii)
Calculate the pH of a mixture of 50 cm3 of ammonia solution of concentration 0.10 mol
dm–3 and 50 cm3 of hydrochloric acid solution of concentration 0.050 mol dm–3 .
.....................................................................................................................................
.....................................................................................................................................
.....................................................................................................................................
.....................................................................................................................................
(4)
(Total 7 marks)
46.
The pH of a solution is 4.8. Using information from Table 17 of the Data Booklet, deduce and
explain the colours of the indicators bromophenol blue and phenol red in this solution.
(Total 3 marks)
47.
Calculate the pH of a buffer solution containing 0.0500 mol dm–3 of ethanoic acid
(Ka = 1.7410–5 ) and 0.100 mol–3 of sodium ethanoate.
(Total 3 marks)
48.
Describe the composition and behaviour of a buffer solution.
(Total 3 marks)
16
49.
(i)
Define the term pH.
(1)
(ii)
A 25.0 cm3 sample of 0.100 mol dm–3 hydrochloric acid was placed in a conical flask,
and 0.100 mol dm–3 sodium hydroxide is added until a total of 50.0 cm3 had been added.
Sketch a graph of pH against volume of NaOH(aq) added, clearly showing the volume of
NaOH(aq) needed for complete reaction and the pH values at the start, the equivalence
point and finish.
(4)
(iii)
The experiment in (ii) was repeated, but with a 25.0 cm3 sample of 0.100 mol dm–3
ethanoic acid in the conical flask instead of the hydrochloric acid. Use information from
Table 16 of the Data Booklet to calculate the pH at the start of the experiment. State the
approximate pH value at the equivalence point.
(5)
(Total 10 marks)
50.
(i)
Describe how an indicator, HIn, works.
(3)
(ii)
Name a suitable indicator for the reaction between ethanoic acid and sodium hydroxide.
Use information from Table 17 in the Data Booklet to explain your choice.
(2)
(Total 5 marks)
51.
(i)
Identify two substances that can be added to water to form a basic buffer solution.
(1)
(ii)
Describe what happens when a small amount of acid solution is added to the buffer
solution prepared in (i). Use an equation to support your explanation.
(2)
(Total 3 marks)
52.
Predict and explain whether an aqueous solution of 0.10 mol dm–3 AlCl3 will be acidic, alkaline
or neutral.
(Total 2 marks)
53.
A titration was carried out to determine the concentration of 25.0 cm3 of an aqueous solution of
nitric acid. The pH value of the liquid in the flask was measured as 0.100 mol dm–3 , aqueous
sodium hydroxide was added. The results are shown on the graph below.
14
12
10
pH
8
6
4
2
0
30
20
10
Volume of NaOH(aq) / cm 3
40
17
(i)
Use the graph to determine the value of [H +] of the nitric acid solution.
(1)
(ii)
Determine the pH value when the value of [H+] has decreased to 1×10–3 mol dm–3 .
(1)
(iii)
Use the graph to determine the volume of 0.100 mol dm–3 aqueous sodium hydroxide
solution needed to exactly neutralize the nitric acid.
(1)
(iv)
Calculate the concentration, in mol dm–3 , of the nitric acid.
(2)
(Total 5 marks)
54.
In aqueous solution at 298 K, ammonia is a weak base with a pKb value of 4.75 and a Kb value
of 1.7×10–5 mol dm–3 .
(a)
Write an equation for the reaction of ammonia with water.
....................................................................................................................................
....................................................................................................................................
(1)
(b)
State the ionization constant expression, Kb, for ammonia.
....................................................................................................................................
....................................................................................................................................
....................................................................................................................................
(1)
(c)
Calculate the pH of a 0.25 mol dm–3 solution of ammonia.
....................................................................................................................................
....................................................................................................................................
....................................................................................................................................
....................................................................................................................................
....................................................................................................................................
....................................................................................................................................
(3)
(Total 5 marks)
55.
Nitric acid and ammonia may be used to make a buffer solution.
(i)
Describe the behaviour of a buffer solution.
(2)
(ii)
Describe how you could prepare a buffer solution using 0.100 mol dm–3 solutions of
nitric acid and ammonia.
(3)
(Total 5 marks)
18
56.
With reference to Table 16 in the Data Booklet, determine the pH of a 0.100 mol dm–3 solution
of propanoic acid.
..............................................................................................................................................
..............................................................................................................................................
..............................................................................................................................................
..............................................................................................................................................
..............................................................................................................................................
..............................................................................................................................................
(Total 3 marks)
57.
0.100 mol dm–3 hydrochloric acid solution is added to 25.0 cm3 0.100 mol dm–3 ammonia
solution and the pH is recorded until a total of 35.0 cm3 hydrochloric acid has been added.
(i)
Sketch a graph to show how the pH changes as hydrochloric acid is added to the
ammonia solution. Use a pH scale of 0–14, and an acid volume scale of 0–35 cm3 .
Explain the shape of the curve.
(6)
(ii)
Use table 17 of the Data Booklet to suggest an indicator that could be used in the titration,
explaining your choice.
(2)
(Total 8 marks)
58.
(i)
State the composition of an acidic buffer solution.
(1)
(ii)
Suggest the identity of an acid and its amount that could be added to a solution containing
0.10 mol ammonia in order to prepare a buffer.
(2)
(iii)
Explain how the solution you prepare in (ii) can act as a buffer solution when a strong
acid and a strong base are added to separate portions of it. Write an equation to illustrate
the buffer action in each case.
(4)
(iv)
Write an equation for the reaction of ammonia with water, and write its Kb expression.
(2)
(Total 9 marks)
59.
(a)
Predict and explain, using equations where appropriate, whether the following solutions
are acidic, alkaline or neutral.
(i)
0.1 mol dm–3 FeCl3 (aq)
...........................................................................................................................
...........................................................................................................................
...........................................................................................................................
...........................................................................................................................
(1)
19
(ii)
0.1 mol dm–3 NaNO3 (aq)
...........................................................................................................................
...........................................................................................................................
...........................................................................................................................
(1)
(iii)
0.1 mol dm–3 Na2 CO3 (aq)
...........................................................................................................................
...........................................................................................................................
...........................................................................................................................
...........................................................................................................................
(1)
(b)
Acidic gases can be released into the atmosphere that have an environmental impact when
they are deposited as acid rain. State two elements that form the acidic gases and describe
two impacts they have on the natural environment.
.....................................................................................................................................
.....................................................................................................................................
.....................................................................................................................................
(3)
(Total 6 marks)
60.
An experiment was carried out to determine the concentration of aqueous ammonia by titrating
it with a 0.150 mol dm–3 sulfuric acid solution. It was found that 25.0 cm3 of the aqueous
ammonia required 20.1 cm3 of the sulfuric acid solution for neutralization.
(a)
Write the equation for the reaction and calculate the concentration, in mol dm–3 , of the
aqueous ammonia.
(4)
(b)
Several acid-base indicators are listed in Table 16 of the Data Booklet. Identify one
indicator that could be used for this experiment. Explain your answer.
(3)
(c)
(i)
Determine the pOH of 0.121 mol dm–3 aqueous ammonia (pKb = 4.75).
(4)
(ii)
State what is meant by the term buffer solution, and describe the composition of an
acid buffer solution in general terms.
(3)
(iii)
Calculate the pH of a mixture of 50.0 cm3 of 0.100 mol dm–3 aqueous ammonia
and 50.0 cm3 of 0.0500 mol dm–3 hydrochloric acid solution.
(4)
(Total 18 marks)
20
IB Chemistry – HL
Topic 8 Answers
1.
D
2.
B
3.
A
4.
5.
D
C
6.
A
7.
C
8.
B
9.
B
10.
B
11.
B
12.
C
13.
14.
D
B
15.
C
16.
C
17.
A
18.
B
19.
A
20.
A
21.
A
22.
C
23.
A
24.
B
25.
C
26.
D
27.
C
28.
A
29.
B
21
30.
C
31.
B
32.
A
33.
D
34.
D
35.
(a)
(i)
pKa = 3.75, therefore Ka = 1.78×10–4 (accept 1.8×10–4 )
No units required.
(ii)
weak acid;
less [H+]/partial dissociation/more reactants/less products/
Ka << 1/small Ka;
(iii)
(HCOOH(aq)
Ka =
1
2
H+(aq) + HCOO– (aq))
2
[H + ][HCOO – ]
= x ;
[HCOOH]
0.010
(x2 = 1.78×10–6 )
x = 1.33×10–3 mol dm–3 = [H+] (no mark without units);
ECF from (a)(i).
No penalty for incorrect significant figures.
pH = 2.88/2.9 (ECF);
assume x << 0.010/25°C/negligible dissociation;
(b)
add strong base/sodium hydroxide or other named alkali/salt of methanoic
acid/HCOONa to methanoic acid;
in equimolar amounts/quantities/so that [HCOOH] = [HCOO – ];
(from Ka expression) pH = pKa (= 3.75);
4
3
[10]
36.
H+(aq) + In– (aq);
needed for mark. State symbols not essential.
(a)
HIn(aq)
(b)
(i)
yellow as equilibrium shifts to left to remove (added) H +(aq);
Colour and explanation needed for the mark.
(ii)
green/blue-yellow;
both HIn(aq) and In– (aq) are present;
1
1
2
[4]
37.
(a)
(b)
(i)
Kw = [H+][OH−];
1
(ii)
[H+] = 1.5×10−7 (mol dm−3 );
Accept answer in range 1.5 to 1.55.
1
(i)
CH3 CH(OH)COOH + H2 O
CH3 CH(OH)COO− + H3 O+;
1
22
Ignore state symbols even if incorrect.
The double arrow is necessary for the mark.
(ii)
ka =
[CH 3 CH(OH)COO − ][H +
;
[CH 3 CH(OH)COOH ]
1
Allow [H3 O+] for [H+] in the expression.
(iii)
(iv)
5.3×10−3 = [H+];
pH = 2.3;
Allow ECF pH based on wrong [H +] in the value, award [1].
Award [2] for correct pH.
2
pH = 3.85;
Accept answer in range 3.8 to 3.9.
1
[7]
38.
(a)
(i)
NH3 + H2 O
NH4 + + OH−;
Do not penalise →
Do not accept NH4 OH
1
+
(ii)
(b)
Kb =
[NH 4 ][OH − ]
;
[NH 3 ]
1
Kb = 10−4.75 = 1.78×10−5 ;
[OH − ] = 1.78 10 −5  0.2 = (1.89 10 −3 );
pOH = −log[OH−] = 2.72;
Accept answer in range 2.68 to 2.76.
Correct answer scores [3].
Apply ECF throughout this part.
3
[5]
39.
(a)
C6 H5 COOH + H2 O
C6 H5 COO− + H3 O+;
Ignore state symbols.
Accept C6 H5 COOH
1
C6 H5 COO− + H+.
needed for mark.
(b)
Ka (= 10−4.20 ) = 6.31×10−5 (mol dm−3 );
Units not needed for mark, but penalize incorrect units.
(c)
[C 6 H 5 COO − ][H + ] [C 6 H 5 CO 2− ][H 3 O + ]
/
;
Ka =
[C 6 H 5 COOH]
[C 6 H 5 CO 2 H]
[H+] =  Ka[C6 H5 COOH]/3.55×10−3 (mol dm−3 );
pH = 2.45;
Apply ECF from (b) and from [H +] to pH.
1
3
23
Correct final answer scores [3].
[5]
40.
(a)
(b)
(Kw = ) [H+][OH−]/[H3 O+] [OH−];
= 2.89×10−14 / 2.9×10−14 ;
Units not needed.
2
pH = 6.8;
1
Accept answer in range 6.7 to 6.8.
(c)
neutral;
[H+] = [OH−]/[H3 O+] = [OH−]/OWTTE;
2
[5]
41.
(i)
(ii)
(iii)
acidic;
Fe(H2 O)6 3+ is a weak acid/Fe 3+ reacts with OH– /equation to show
formation of HCl or H+;
“FeCl3 is acidic” is not acceptable.
2
neutral;
NaNO3 / sodium nitrate is formed from strong base and strong
acid/ions do not hydrolyse;
2
alkaline;
As CO3 2– is weak base/combines with H +/equation showing
formation of OH– ;
Acidic, neutral, alkali mark in each case is independent of
reason.
2
[6]
42.
(i)
(ii)
(iii)
8.7 ± 0.7;
low [H+] thus small addition of OH – has great effect/OH– increases
rapidly as NaOH is a strong base/logarithmic nature of pH;
volume of NaOH = 8.2 cm3 (exact);
amount of NaOH = 8.2 ×0.1 = 0.00082 mol;
1000
0
.
00082
[HA] =
×0.1 = 0.082 mol dm–3 /0.082 M;
0.010
Correct answer [3], units needed for last mark.
2
3
correct pH reading from graph (2.9) (allow 2.8 or 3.0);
thus [H+] = 1.26×10–3 (mol dm–3 )
–2.9
–2.9
Ka = 10 10
0.082
= 1.9×10–5 (mol dm–3 );
pKa = 4.71
Accept 4.7 and allow ECF from (ii).
5
If pH given as 2.8, Ka = 3.06×10 –5 and pKa = 4.51
If pH given as 2.8, Ka = 1.22×10 –5 and pKa = 4.91
24
If half equivalence method used:
Volume = 4.1 cm3
pKa = 4.75
Award [2] out of last [4].
[10]
43.
(i)
(ii)
a solution that resists pH change/maintains a (nearly) constant pH;
when small amounts of acid or alkali are added;
2
M r of sodium ethanoate;
moles of sodium ethanoate = 0.25 = (0.0030);
82
[CH3 COO– ] = 0.0030 = 0.015 (mol dm–3 ) 2 sig figs only;
0 .2
(iii)
3
[H + ][CH 3 COO – ]
(or with substituted values);
[CH 3 COOH]
May be assumed from later work.
Ka =
–4.76
 0.10 = (1.159×10–4 );
[H+] = 10
0.015
pH = 3.9(4);
3
Allow ECF throughout (ii) and (iii).
[8]
44.
(a)
2NH3 + H2 SO4 → (NH4 )2 SO4
Accept correct equation with NH 4OH instead of NH3.
.
mol H2 SO4 = 0.0201×0.150;
2NH3 = H2 SO4 /mol NH3 = 6.03×10–3 ;
[NH3 ] = 0.241 (mol dm–3 );
Apply –1(SF) if appropriate.
Award [3] for the correct final answer for the
concentration calculation.
(b)
(c)
bromocresol green;
reaction of weak base and strong acid/OWTTE;
pH range of bromocresol green is 3.8 to 5.4 / occurs at pH < 7;
4
3
Kb = 10–4.75 = 1.78×10–5 ;
Kb =
[NH 4+ ][OH − ]
[OH– ] =
[NH 3 ]
/[OH– ] =
K b [ NH 3 ] ;
1.78  10 –5  0.121 ;
pOH = 2.83;
4
Award [4] for the correct final answer.
Allow ECF, for example any correct conversion of [OH– ]
[11]
25
45.
(i)
(ii)
a solution which resists change in pH/changes pH very slightly/
keeps pH constant/OWTTE;
when small amounts of acid or base are added;
weak acid and its salt/weak acid and its conjugate base;
mol NH3 = 0.0050 and mol HCl = 0.0025;
[NH4 +] = [NH3 ];
[OH– ] = Kb = 1.78×10–5
(pOH = 4.75 so) pH = 9.25 (allow 9.2 to 9.3);
Award [4] for correct final answer.
Accept other valid methods such as Henderson-Hasselbach
equation.
3
4
[7]
46.
bromophenol blue is blue and phenol red is yellow;
pH of 4.8 is above range of bromophenol blue/bromphenol blue shows
its alkaline colour/OWTTE;
pH of 4.8 is below range of phenol red/phenol red shows its acidic
colour/OWTTE;
3
[3]
47.
Ka =
−
+
[CH 3 COO ][H ]
/rearrangement for [H+];
[CH 3 COOH]
[H + ] =
1.74 10 −5  0.0500
= 8.70 10 −6 (mol dm −3 ) ;
0.100
pH ( = −log[H+]) = 5.06;
3
OR
pH = p k a + log
[CH 3 COO − ]
;
[CH 3 COOH]
 0.10 
pH = 4.76 + log 
;
 0.05 
pH = 5.06;
Accept answer in range 5.0 to 5.1.
ECF from [H+].
Award [3] for correct final answer.
[3]
48.
weak acid + salt of weak acid/weak acid + conjugate base.
Accept equivalent descriptions of a basic buffer.
the solution resists pH change;
Do not accept pH does not change.
when small amounts of acid or base are added;
Only award if previous answer correct.
[3]
26
49.
(i)
pH = − log[H+];
(ii)
curve should include the following:
starting pH = 1;
equivalence point: 25.0 cm3 of NaOH;
pH at equivalence point = 7;
pH to finish = 12−13;
1
pH
13
7
1
25
Volume / cm3
4
Penalise [1] if profile incorrect.
(iii)
Ka = 10−4.76 /1.74×10−5 ;
Ka = [H+]2 ÷[CH3 COOH]/1.74×10−5 =
[H+] = 1.32×10−3 (mol dm−3 );
starting pH = 2.88;
Accept 3 sig. fig.
Award [4] for correct pH.
Allow ECF.
[H + ] 2
;
0.100
pH at equivalence point: 8−9;
5
[10]
50.
(i)
HIn is a weak acid;
HIn
H+ + In− and two colours indicated;
In acid equilibrium moves left or vice versa;
(ii)
phenolphthalein/phenol red/bromothymol blue;
colour change of indicator occurs within the range of pH at equivalence
point/on vertical part of graph;
3
2
[5]
51.
(i)
(ii)
specific examples of weak base and its salt/specific strong acid and
weak base;
e.g. NH3 and NH4 Cl.
1
pH changes very little/most acid neutralized by base;
equation from (i);
e.g. NH3 + H+ → NH 4+ /NH4 OH + H+ → NH 4+ + H2 O.
2
[3]
27
52.
acidic;
[Al(H2 O6 )]3+ is (weak) acid due to the formation of H +/
[Al(H2 O)6 ]3+
[Al(H2 O)5 (OH)]2+ + H+;
2
[2]
53.
(i)
0.1 (mol dm−3 );
1
(ii)
3;
1
(iii)
28(.0) (cm3 );
1
(iv)
nNaOH/HNO3 ( = 0.100×0.0280) = 2.80×10−3 (mol);
ECF from value in (iii).
[HNO3 ] (= 2.80×10−3 ÷0.025) = 0.112 (mol dm−3 );
ECF from n above.
Correct final answer scores [2].
2
[5]
54.
(a)
(b)
(c)
NH3 (aq) + H2 O(l)
N H 4+ (aq) + OH−(aq);
Ignore state symbols and accept →.
Kb =
[NH 4+ ][OH − ]
[NH 3 ]
[OH−] = 2.1×10−3
pOH = 2.7/[H+] = 4.8×10−12
pH = 11.3;
Allow ECF for the value of pOH and pH.
1
1
3
[5]
55.
(i)
(ii)
a solution which resists change in pH;
when a small amount of strong acid or base is added to it;
react excess ammonia with nitric acid;
stated volumes with about 50% more ammonia solution;
gives a solution containing the weak base and its salt with the acid/
NH4 + and NH3 ;
Accept suitable volumes from about 20 cm3 to about 500 cm3
for 2 nd mark.
2
3
[5]
56.
(pKa (propanoic) = 4.87)
ka =
[CH 3 CH 2 COO − ][H 3 O + ]
;
[CH 3 CH 2 COOH]
[H3 O+]= 1.16×10−3 (mol dm−3 );
pH = 2.94;
3
Award [3] for correct answer.
[3]
28
57.
(i)
13
pH
7
1
25
/ cm 3 Hcl
graph starting at pH < 13;
Award [0] for pH=13.
equivalence point pH < 7;
Accept anything between 4 and 6
bottom end of graph: pH between 3 and 1;
NH3 is a weak base/partially dissociated/[OH −]</<<0.10mol dm−3
(therefore, pH < 13);
(ii)
NH4 + formed is a weak acid/NH 4 +
NH3 + H+/NH4 + dissociates into
a weak base and a strong acid (thus acidic at equivalence point);
HCl is a strong acid, thus graph finishes close to pH = 1;
6
methyl orange/bromocresol green/bromophenol blue/methyl red;
pKa of indicator centred around pH at equivalence/end point/indicator
pH range falls where there is a sharp pH change/OWTTE;
2
[8]
58.
(i)
weak acid and salt of the weak acid/its conjugate base;
(ii)
HCl/HNO3 /H2 SO4 ;
Amount < 0.10 mol for HCl/HNO 3 /< 0.05 mol for H2 SO4 ;
(iii)
1
2
(added) OH− reacts with NH4 + present/acid of buffer;
(added) H+ reacts with NH3 present/base of buffer;
OH− + NH4 + → NH3 + H2 O (strong base replaced by weak base);
H+ + NH3 → NH4 + (strong acid replaced by weak acid);
(iv)
4
NH3 (aq) + H2 O(l)
NH4 +(aq) + OH−(aq);
States not required for mark
29
+
Kb =
[NH 4 ][OH − ]
;
[NH 3 ]
2
[9]
59.
(a)
(i)
(ii)
(iii)
(b)
acidic and [Fe(H2 O)6 ]3+ is a weak acid
[Fe(H2 O)6 ]3+(aq) → [Fe(OH)(H2O)5]2+(aq) + H+(aq);
“FeCl3 is acidic” is not acceptable.
neutral and NaNO3 /sodium nitrate is formed from strong base
and strong acid/ions do not hydrolyze;
alkaline and CO3 2– is a weak base/
CO3 2– (aq) + H2 O(l) → HCO3 –(aq) + OH– (aq);
Award [1] only for correct identification of solutions as acidic,
neutral and alkaline only, without explanation.
nitrogen and sulfur;
kills/harms fish/aquatic life in lakes/rivers;
leaching of soils damages plant life/trees;
1
1
1
3
[6]
60.
(a)
2NH3 (aq) + H2 SO4 (aq) → (NH4 )2SO4(aq);
Accept correct equation with NH 4OH instead of NH3.
n(H2 SO4 ) = 0.0201×0.150 (mol);
n(NH3 ) = 6.03×10–3 (mol);
[NH3 ] = 0.241 (mol dm–3 );
Award [3] for the correct final answer for the concentration
calculation.
(b)
bromocresol green;
reaction of weak base and strong acid;
pH range of bromocresol green is 3.8 to 5.4/occurs at pH < 7;
(c)
(i)
4
3
Kb = 10–4.75 = 1.78×10–5 ;
[ NH 4+ ][OH − ]
Kb =
/[ OH − ] = K b [ NH 3 ];
[ NH 3 ]
[OH– ] = 1.78  10 −5  0.121 ;
pOH = 2.83;
Award [4] for the correct final answer.
4
Allow ECF, for example any correct conversion of [OH– ] to
pOH.
(ii)
(iii)
a solution which resists change in pH/changes pH very slightly;
when small amounts of acid or base are added;
weak acid and its salt/weak acid and its conjugate base;
3
n(NH3 ) = 0.00500 (mol) and n(HCl) = 0.00250 (mol);
[ NH 4+ ] = [ NH 3 ];
[OH– ] = Kb = 1.78×10–5 ;
(pOH = 4.75 so) pH = 9.25 (allow 9.2 to 9.3);
Award [4] for correct final answer.
4
[18]
30
IB Chemistry – SL
Topic 9 Questions
1.
Which statement is correct?
A.
Oxidation involves loss of electrons and a decrease in oxidation state.
B.
Oxidation involves gain of electrons and an increase in oxidation state.
C.
Reduction involves loss of electrons and an increase in oxidation state.
D.
Reduction involves gain of electrons and a decrease in oxidation state.
(Total 1 mark)
2.
What occurs during the operation of a voltaic cell based on the following reaction?
Ni(s) + Pb2+(aq) → Ni2+(aq) + Pb(s)
Ion movement in solution
2+
External circuit
electrons move from Ni to Pb
A.
Pb (aq) move toward Pb(s)
electrons move from Ni to Pb
B.
Ni2+(aq) move away from Ni(s)
electrons move from Pb to Ni
C.
Ni2+(aq) move toward Ni(s)
electrons move from Pb to Ni
D.
Pb (aq) move away from Pb(s)
2+
(Total 1 mark)
3.
The oxidation number of chromium is the same in all the following compounds except
A.
Cr(OH)3
B.
Cr2 O3
C.
Cr2 (SO4 )3
D.
CrO3
(Total 1 mark)
4.
Magnesium is a more reactive metal than copper. Which is the strongest oxidizing agent?
A.
Mg
B.
Mg2+
C.
Cu
D.
Cu2+
(Total 1 mark)
Page 1 of 19
5.
Which processes occur during the electrolysis of molten sodium chloride?
I.
Sodium and chloride ions move through the electrolyte.
II.
Electrons move through the external circuit.
III.
Oxidation takes place at the positive electrode (anode).
A.
I and II only
B.
I and III only
C.
II and III only
D.
I, II and III
(Total 1 mark)
6.
What happens to the Cr3+(aq) ion when it is converted to CrO 4 2– (aq)?
A.
Its oxidation number decreases and it undergoes reduction.
B.
Its oxidation number decreases and it undergoes oxidation.
C.
Its oxidation number increases and it undergoes reduction.
D.
Its oxidation number increases and it undergoes oxidation.
(Total 1 mark)
7.
The following reactions are spontaneous as written.
Fe(s) + Cd2+(aq) → Fe2+(aq) + Cd(s)
Cd(s) + Sn2+(aq) → Cd2+(aq) + Sn(s)
Sn(s) + Pb2+(aq) → Sn2+(aq) + Pb(s)
Which of the following pairs will react spontaneously?
I.
Sn(s) + Fe 2+(aq)
II.
Cd(s) + Pb2+(aq)
III.
Fe(s) + Pb2+(aq)
A.
I only
B.
II only
C.
III only
D.
II and III only
(Total 1 mark)
Page 2 of 19
8.
What species are produced at the positive and negative electrodes during the electrolysis of
molten sodium chloride?
Negative electrode
Cl2 (g)
Positive electrode
+
Na (l)
–
Cl (l)
B.
Cl2 (g)
Na(l)
C.
Na(l)
Cl2 (g)
D.
Na+(l)
A.
(Total 1 mark)
9.
Consider the following reaction.
H2 SO3 (aq) + Sn4+(aq) + H2 O(l) → Sn2+(aq) + HSO4 – (aq) + 3H+(aq)
Which statement is correct?
A.
H2 SO3 is the reducing agent because it undergoes reduction.
B.
H2 SO3 is the reducing agent because it undergoes oxidation.
C.
Sn4+ is the oxidizing agent because it undergoes oxidation.
D.
Sn4+ is the reducing agent because it undergoes oxidation.
(Total 1 mark)
10.
In which change does oxidation occur?
A.
CH3 CHO → CH3 CH2OH
B.
CrO4 2– → Cr2 O7 2–
C.
SO4 2– → SO3 2–
D.
NO2 – → NO3 –
(Total 1 mark)
11.
What happens at the positive electrode in a voltaic cell and in an electrolytic cell?
Electrolytic cell
Reduction
Oxidation
Oxidation
Reduction
Voltaic cell
Oxidation
Reduction
Oxidation
Reduction
A.
B.
C.
D.
(Total 1 mark)
Page 3 of 19
12.
What are the oxidation numbers of the elements in sulfuric acid, H 2 SO4 ?
Oxygen
Sulfur
Hydrogen
–2
+6
+1
A.
–2
+4
+1
B.
+4
+1
+2
C.
–8
+6
+2
D.
(Total 1 mark)
13.
A voltaic cell is made from copper and zinc half-cells. The equation for the reaction occurring
in the cell is
Zn(s) + Cu2+(aq) → Zn2+(aq) + Cu(s)
Which statement is correct when the cell produces electricity?
A.
Electrons are lost from zinc atoms.
B.
The mass of the copper electrode decreases.
C.
Electrons flow from the copper half-cell to the zinc half-cell.
D.
Negative ions flow through the salt bridge from the zinc half-cell to the copper half-cell.
(Total 1 mark)
14.
What happens when molten sodium chloride is electrolysed in an electrolytic cell?
A.
Chlorine is produced at the positive electrode.
B.
Sodium ions lose electrons at the negative electrode.
C.
Electrons flow through the liquid from the negative electrode to the positive electrode.
D.
Oxidation occurs at the negative electrode and reduction at the positive electrode.
(Total 1 mark)
15.
Which equations represent reactions that occur at room temperature?
I.
2Br– (aq) + Cl2 (aq) → 2Cl– (aq) + Br2 (aq)
II.
2Br– (aq) + I2 (aq) → 2I– (aq) + Br2 (aq)
III.
2I– (aq) + Cl2 (aq) → 2Cl– (aq) + I2 (aq)
A.
I and II only
B.
I and III only
C.
II and III only
D.
I, II and III
(Total 1 mark)
Page 4 of 19
16.
Which equation represents a redox reaction?
A.
KOH(aq) + HCl(aq) → KCl(aq) + H2 O(l)
B.
Mg(s) + 2HCl(aq) → MgCl2 (aq) + H2 (g)
C.
CuO(s) + 2HCl(aq) → CuCl2 (aq) + H2 O(l)
D.
ZnCO3 (s) + 2HCl(aq) → ZnCl2 (aq) + CO2 (g) + H2 O(l)
(Total 1 mark)
17.
The following information is given about reactions involving the metals X, Y and Z and
solutions of their sulfates.
X(s) + YSO4 (aq) → no reaction
Z(s) + YSO4 (aq) → Y(s) + ZSO4 (aq)
When the metals are listed in decreasing order of reactivity (most reactive first), what is the
correct order?
A.
Z Y X
B.
XYZ
C.
YXZ
D.
Y Z X
(Total 1 mark)
18.
What are the oxidation numbers of the elements in the compound phosphoric acid, H 3 PO4 ?
Oxygen
Phosphorus
Hydrogen
–2
+1
+1
A.
–2
+5
+1
B.
–4
+1
+3
C.
–8
+5
+3
D.
(Total 1 mark)
19.
A voltaic cell is made from magnesium and iron half-cells. Magnesium is a more reactive metal
than iron. Which statement is correct when the cell produces electricity?
A.
Electrons are lost from magnesium atoms.
B.
The concentration of Fe 2+ ions increases.
C.
Electrons flow from the iron half-cell to the magnesium half-cell.
D.
Negative ions flow through the salt bridge from the magnesium half-cell to the iron
half-cell.
(Total 1 mark)
Page 5 of 19
20.
Which are examples of reduction?
I.
Fe3+ becomes Fe 2+
II.
Cl– becomes Cl2
III.
CrO3 becomes Cr3+
A.
I and II only
B.
I and III only
C.
II and III only
D.
I, II and III
(Total 1 mark)
21.
Which statement is correct for the electrolysis of a molten salt?
A.
Positive ions move toward the positive electrode.
B.
A gas is produced at the negative electrode.
C.
Only electrons move in the electrolyte.
D.
Both positive and negative ions move toward electrodes.
(Total 1 mark)
22.
Which statement about the following reaction is correct?
2Br– (aq) + Cl2 (aq) → Br2 (aq) + 2Cl– (aq)
A.
Br– (aq) is reduced and gains electrons.
B.
Cl2 (aq) is reduced and loses electrons.
C.
Br– (aq) is oxidized and loses electrons.
D.
Cl2 (aq) is oxidized and gains electrons.
(Total 1 mark)
Page 6 of 19
23.
Consider the following spontaneous reactions.
Fe(s) + Cu2+(aq) → Fe2+(aq) + Cu(s)
Cu(s) + 2Ag+(aq) → Cu2+(aq) + 2Ag(s)
Zn(s) + Fe 2+(aq) → Zn2+(aq) + Fe(s)
Which is the correct combination of strongest oxidizing agent and strongest reducing agent?
Strongest reducing agent
Strongest oxidizing agent
Zn(s)
Ag(s)
A.
Zn(s)
Ag+(aq)
B.
Ag(s)
Zn2+(aq)
C.
Ag+(aq)
Zn(s)
D.
(Total 1 mark)
24.
In which change does nitrogen undergo oxidation?
A.
NO2 → N2 O4
B.
NO3 – → NO2
C.
N2 O5 → NO3 –
D.
NH3 → N2
(Total 1 mark)
25.
Which statement is correct?
A.
Spontaneous redox reactions produce electricity in an electrolytic cell.
B.
Electricity is used to carry out a non-spontaneous redox reaction in a voltaic cell.
C.
Oxidation takes place at the negative electrode in a voltaic cell and the positive electrode
in an electrolytic cell.
D.
Oxidation takes place at the negative electrode in a voltaic cell and reduction takes place
at the positive electrode in an electrolytic cell.
(Total 1 mark)
26.
The compound [Co(NH3 )5 Br]SO4 is isomeric with the compound [Co(NH 3 )5 SO4]Br. What is
the oxidation state of cobalt in these compounds?
[Co(NH3 )5 SO4 ]Br
[Co(NH3 )5 Br]SO4
+3
+3
A.
+1
+2
B.
+2
+3
C.
+3
+2
D.
(Total 1 mark)
Page 7 of 19
27.
What happens to vanadium during the reaction VO 2+(aq) → VO3 – (aq)?
A.
It undergoes oxidation and its oxidation number changes from +4 to +5.
B.
It undergoes oxidation and its oxidation number changes from +2 to +4.
C.
It undergoes reduction and its oxidation number changes from +2 to –1.
D.
It undergoes reduction and its oxidation number changes from +4 to +2.
(Total 1 mark)
28.
What occurs during the electrolysis of a molten salt?
A.
Electricity is produced by a spontaneous redox reaction.
B.
Electricity is used to cause a non-spontaneous redox reaction to occur.
C.
Electrons flow through the molten salt.
D.
Electrons are removed from both ions of the molten salt.
(Total 1 mark)
29.
Which statement is correct about an oxidizing agent in a chemical reaction?
A.
It reacts with oxygen.
B.
It reacts with H+ ions.
C.
It loses electrons.
D.
It undergoes reduction.
(Total 1 mark)
30.
Which formula represents an aldehyde?
A.
CH3 CH2 CHO
B.
CH3 COCH3
C.
CH3 CH2 COOH
D.
CH3 COOCH3
(Total 1 mark)
31.
What is the reducing agent in this reaction?
Cu(s) + 2 NO 3− (aq) + 4H+(aq) → Cu2+(aq) + 2NO2 (g) + 2H2 O(l)
A.
Cu(s)
B.
NO 3− (aq)
C.
Cu2+(aq)
D.
H+(aq)
(Total 1 mark)
Page 8 of 19
32.
A particular voltaic cell is made from magnesium and iron half-cells. The overall equation for
the reaction occurring in the cell is
Mg(s) + Fe2+(aq) → Mg2+(aq) + Fe(s)
Which statement is correct when the cell produces electricity?
A.
Magnesium atoms lose electrons.
B.
The mass of the iron electrode decreases.
C.
Electrons flow from the iron half-cell to the magnesium half-cell.
D.
Negative ions flow through the salt bridge from the magnesium half-cell to the iron
half-cell.
(Total 1 mark)
33.
What process occurs at the cathode in a voltaic cell and at the anode in an electrolytic cell?
Anode of
Electrolytic cell
Reduction
Cathode of
voltaic cell
Oxidation
A.
Oxidation
Oxidation
B.
Oxidation
Reduction
C.
Reduction
Reduction
D.
(Total 1 mark)
34.
Consider the following reaction:
H2 SO3 (aq) + Sn4+(aq) + H2 O(l) → Sn2+(aq) + HSO4 – (aq) + 3H+(aq)
Which statement is correct?
A.
H2 SO3 is the reducing agent because it undergoes reduction.
B.
H2 SO3 is the reducing agent because it undergoes oxidation.
C.
Sn4+ is the oxidizing agent because it undergoes oxidation.
D.
Sn4+ is the reducing agent because it undergoes oxidation.
(Total 1 mark)
Page 9 of 19
35.
Which processes occur during the electrolysis of molten sodium chloride?
I.
II.
III.
Sodium and chloride ions move through the electrolyte.
Electrons move through the external circuit.
Oxidation takes place at the anode.
A.
I and II only
B.
I and III only
C.
II and III only
D.
I, II and III
(Total 1 mark)
36.
Tin(II) ions can be oxidized to tin(IV) ions by acidified potassium permanganate(VII) solution
according to the following unbalanced equation.
__ Sn2+ + __MnO4 – + __ H+ →__ Sn4+ + __Mn2+ + __ H2 O
(a)
Identify the oxidizing agent and the reducing agent.
Oxidizing agent .........................................................................................................
Reducing agent ..........................................................................................................
(1)
(b)
Balance the equation above.
....................................................................................................................................
....................................................................................................................................
(1)
(Total 2 marks)
37.
Consider the following redox equation.
5Fe2+(aq) +MnO4 – (aq) +8H+(aq) → 5Fe3+(aq) + Mn2+(aq) + 4H2 O(l)
(i)
Determine the oxidation numbers for Fe and Mn in the reactants and in the products.
……………………………………………………………………………………………
……………………………………………………………………………………………
……………………………………………………………………………………………
……………………………………………………………………………………………
(2)
(ii)
Based on your answer to (i), deduce which substance is oxidized.
……………………………………………………………………………………………
(1)
Page 10 of 19
(iii)
The compounds CH3 OH and CH2 O contain carbon atoms with different oxidation
numbers. Deduce the oxidation numbers and state the kind of chemical change needed to
make CH2 O from CH3 OH.
……………………………………………………………………………………………
……………………………………………………………………………………………
……………………………………………………………………………………………
(3)
(Total 6 marks)
38.
A part of the reactivity series of metals, in order of decreasing reactivity, is shown below.
magnesium
zinc
iron
lead
copper
silver
If a piece of copper metal were placed in separate solutions of silver nitrate and zinc nitrate
(i)
determine which solution would undergo reaction.
……………………………………………………………………………………………
(1)
(ii)
identify the type of chemical change taking place in the copper and write the
half-equation for this change.
……………………………………………………………………………………………
……………………………………………………………………………………………
(2)
(iii)
state, giving a reason, what visible change would take place in the solutions.
……………………………………………………………………………………………
……………………………………………………………………………………………
……………………………………………………………………………………………
……………………………………………………………………………………………
(2)
(Total 5 marks)
Page 11 of 19
39.
(i)
Solid sodium chloride does not conduct electricity but molten sodium chloride does.
Explain this difference, and outline what happens in an electrolytic cell during the
electrolysis of molten sodium chloride using carbon electrodes.
……………………………………………………………………………………………
……………………………………………………………………………………………
……………………………………………………………………………………………
……………………………………………………………………………………………
……………………………………………………………………………………………
……………………………………………………………………………………………
……………………………………………………………………………………………
……………………………………………………………………………………………
(4)
(ii)
State the products formed and give equations showing the reactions at each electrode.
……………………………………………………………………………………………
……………………………………………………………………………………………
……………………………………………………………………………………………
……………………………………………………………………………………………
(4)
(iii)
State what practical use is made of this process.
……………………………………………………………………………………………
……………………………………………………………………………………………
(1)
(Total 9 marks)
40.
Electrolysis can be used to obtain chlorine from molten sodium chloride. Write an equation for
the reaction occurring at each electrode and describe the two different ways in which electricity
is conducted when the cell is in operation.
(Total 4 marks)
41.
Two reactions occurring in the manufacture of bromine from sea water are
I
II
(i)
Cl2 (g) + 2Br– (aq) → 2Cl– (aq) + Br2 (g)
Br2 (g) + SO2 (g) + 2H2 O(l) → 2HBr(g) + H2 SO4 (g)
Explain, by reference to electrons, why reaction I is referred to as a redox reaction.
(2)
(ii)
State and explain whether SO2 is reduced or oxidized in reaction II by referring to the
oxidation numbers of sulfur in this reaction.
(3)
(Total 5 marks)
Page 12 of 19
42.
In terms of electron transfer define:
(i)
oxidation
....................................................................................................................................
....................................................................................................................................
(1)
(ii)
oxidizing agent
....................................................................................................................................
....................................................................................................................................
(1)
(Total 2 marks)
43.
Deduce the change in oxidation number of chromium in the below reaction. State with a reason
whether the chromium has been oxidized or reduced.
CrO7 2− + 14H+ + 6Fe2+ → 2Cr3+ + 6Fe3+ + 7H2 O
..............................................................................................................................................
..............................................................................................................................................
..............................................................................................................................................
..............................................................................................................................................
(Total 2 marks)
44.
(a)
(i)
Define oxidizing agent in terms of electron transfer.
.........................................................................................................................
.........................................................................................................................
(1)
(ii)
Deduce the change in oxidation number of chromium in the reaction below. State
with a reason whether the chromium has been oxidized or reduced.
Cr2 O7 2− + 14H+ + 6Fe2+ → 2Cr3+ + 7H2 O
.........................................................................................................................
.........................................................................................................................
.........................................................................................................................
.........................................................................................................................
(2)
(Total 3 marks)
Page 13 of 19
45.
Iron in food, in the form of Fe 3+, reacts with ascorbic acid (vitamin C), C 6 H8 O6 , to form
dehydroascorbic acid, C6 H6 O6 .
(i)
Write an ionic half-equation to show the conversion of ascorbic acid to
dehydroascorbic acid in aqueous solution.
.........................................................................................................................
.........................................................................................................................
(1)
(ii)
In the other ionic half-equation Fe 3+ is converted to Fe 2+. Deduce the overall
equation for the reaction between C6 H8 O6 and Fe3+.
.........................................................................................................................
.........................................................................................................................
(1)
(Total 2 marks)
46.
(i)
Draw a diagram of apparatus that could be used to electrolyse molten potassium bromide.
Label the diagram to show the polarity of each electrode and the product formed.
(3)
(ii)
Describe the two different ways in which electricity is conducted in the apparatus.
(2)
(iii)
Write an equation to show the formation of the product at each electrode.
(2)
(Total 7 marks)
47.
Iodide ions, I– (aq), react with iodate ions, IO 3 – (aq), in an acidic solution to form molecular
iodine and water.
(i)
Determine the oxidation number of iodine in each iodine-containing species in the
reaction.
....................................................................................................................................
....................................................................................................................................
....................................................................................................................................
(2)
(ii)
Identify, with a reason, the species that undergoes:
oxidation ....................................................................................................................
....................................................................................................................................
reduction ....................................................................................................................
....................................................................................................................................
(2)
(Total 4 marks)
Page 14 of 19
IB Chemistry – SL
Topic 9 Answers
1.
D
[1]
2.
B
[1]
3.
D
[1]
4.
D
[1]
5.
D
[1]
6.
D
[1]
7.
D
[1]
8.
D
[1]
9.
B
[1]
10.
D
[1]
11.
B
[1]
12.
A
[1]
13.
A
[1]
14.
A
[1]
15.
B
[1]
16.
B
[1]
17.
A
[1]
18.
B
[1]
19.
A
[1]
20.
B
[1]
21.
D
[1]
Page 15 of 19
22.
C
[1]
23.
B
[1]
24.
D
[1]
25.
C
[1]
26.
A
[1]
27.
A
[1]
28.
B
[1]
29.
D
[1]
30.
A
[1]
31.
A
[1]
32.
A
[1]
33.
C
[1]
34.
B
[1]
35.
D
[1]
36.
(a)
(b)
oxidizing agent: (acidified) potassium permanganate(VII)/(H +) and Mn O 4−
and reducing agent: Sn2+;
Both oxidizing agent and reducing agent required for [1].
1
5Sn2+ + 2MnO4 – + 16H+ → 5Sn4+ + 2Mn2+ + 8H2 O;
1
[2]
37.
(i)
(ii)
Fe reactant +2 AND Fe product +3 AND Mn product +2;
Mn reactant +7;
Do not accept Roman numerals.
Fe2+/iron(ii) ions/ferrous ions;
Do not accept “iron”.
2
1
Page 16 of 19
(iii)
CH3 OH oxidation state –2;
CH2 O oxidation state 0;
(change is) oxidation/dehydrogenation;
3
[6]
38.
(i)
silver nitrate;
(ii)
oxidation;
Cu → Cu2+ + 2e;
2
(silver nitrate) solution turns blue/grey or black or silver solid forms;
copper ions form/Cu2+ ions form/silver deposited;
2
(iii)
1
[5]
39.
(i)
(ii)
(iii)
sodium chloride crystals consist of ions in a rigid lattice/ions can not
move about;
when melted the ions are free to move or ions move when a voltage
is applied;
in electrolysis positive sodium ions or Na + ions move to the negative
electrode or cathode; and negative chloride ions or Cl– move to the
positive electrode or anode;
4
sodium formed at cathode or negative electrode;
Na+ + e → Na;
chlorine formed at anode or positive electrode;
2Cl– → Cl2 + 2e;
1st and 3rd marks can be scored in (i).
4
manufacture of sodium and chlorine/one stated use of chlorine
or sodium;
1
[9]
40.
at negative electrode (cathode)
Na+ + e− → Na;
at positive electrode (anode)
2Cl− → Cl2 + 2e−;
If both equations correct but electrodes incorrect or not stated,
then deduct [1].
electrons flow through the external circuit or wires;
ions gain/lose electrons at electrodes/ions move to electrodes.
4
[4]
41.
(i)
(ii)
chlorine/Cl2 gains electrons and is reduced;
bromide (ions)/Br− loses electrons and is oxidized;
Award [1] max if no mention of reduced and oxidized.
2
S in SO2
+4;
S in H2 SO4
Page 17 of 19
+6;
Award only [1] for 4 + and 6 + or 4 and 6.
SO2 oxidized because oxidation number (of sulfur) increases;
3
[5]
42.
(i)
loss of electrons;
1
(ii)
(a species that) gains electrons (from another species)/causes electron loss;
1
[2]
43.
changes by 3;
reduced because its oxidation number decreased / +6 → +3 / 6+ → 3+ / it has gained
electrons;
[2]
44.
(i)
(a species that) gains electrons (from another species)/causes electron loss;
1
(ii)
changes by 3;
reduced because its oxidation number decreased / +6 → +3 / 6+ → 3+ /
it has gained electrons;
2
[3]
45.
(i)
C6 H8 O6 → C6 H6O6 +2H+ + 2e;
1
(ii)
C6 H8 O6 + 2Fe3+ → C6 H6 O + 2H+ + 2Fe2+;
1
[2]
46.
(i)
(ii)
(iii)
(diagram showing)
container, liquid, electrodes and power supply;
bromine formed at + electrode;
potassium formed at − electrode;
Award [1] for both correct products shown at wrong electrodes,
or if no polarity indicated.
electrons flow through connecting wires;
ions move (through liquid) to electrodes (and lose/gain electrons);
K+ + e− → K;
2Br− → Br2 + 2e−;
No need to indicate polarity of electrodes.
3
2
2
Accept e instead of e −.
[7]
47.
(i)
I− = −1 / 1−
IO3 − = +5 / 5+
I2 = 0
2
Award [2] for all three correct, [1] for any two correct,
Signs must be included
Do not accept Roman numerals
(ii)
oxidation
Page 18 of 19
I− (to I2 ), increase in oxidation number/loss of electron(s);
reduction
IO3 – (to I2 ), decrease in oxidation number/gain of electron(s);
2
[4]
Page 19 of 19
IB Chemistry HL
Topic9 Questions and Answers
1.
2.
Aqueous solutions of AgNO 3 , Cu(NO3 )2 and Cr(NO3 )3 are electrolyzed using the same quantity
of electricity. How do the number of moles of metal formed compare?
A.
Ag = Cu = Cr
B.
Ag > Cu > Cr
C.
Ag < Cu < Cr
D.
Cu > Ag > Cr
The standard electrode potentials for two half-cells involving iron are given below.
Fe2+(aq) + 2e– → Fe(s)
Fe3+(aq) + e– → Fe2+(aq)
Eο = –0.44 V
Eο = +0.77 V
What is the equation and the cell potential for the spontaneous reaction that occurs when the two
half-cells are connected?
3.
A.
3Fe2+(aq) → Fe(s) + 2Fe 3+(aq)
Eο = +1.21 V
B.
Fe2+(aq) + Fe3+(aq) → 2Fe(s)
Eο = +0.33 V
C.
Fe(s) + 2Fe 3+(aq) → 3Fe2+(aq)
Eο = +0.33 V
D.
Fe(s) + 2Fe 3+(aq) → 3Fe2+(aq)
Eο = +1.21 V
Metallic tin can be produced by the electrolysis of a molten salt containing Sn2+ ions. Which
change(s) would double the amount of tin produced?
I.
Doubling the current passed during electrolysis
II.
Doubling the time used for electrolysis
III.
Using Sn4+ ions instead of Sn2+ ions
A.
I only
B.
II only
C.
I and II only
D.
I, II and III
1
4.
5.
Which of the following factors affect the amount of product formed during electrolysis?
I.
The current used
II.
The duration of electrolysis
III.
The charge on the ion
A.
I and II only
B.
I and III only
C.
II and III only
D.
I, II and III
The cyanide ion, CN– , can form two complex ions with iron ions. The formulas of these ions are
[Fe(CN)6 ]4– and [Fe(CN)6]3– . What is the oxidation number of iron in the two complex ions?
4–
[Fe(CN)6 ]
6.
3–
[Fe(CN)6 ]
A.
–4
–3
B.
+2
+3
C.
+3
+2
D.
–3
–4
Consider the following reactions.
Cu2+(aq) + 2e–
Cu(s)
Eο = +0.34 V
Mg2+(aq) + 2e–
Mg(s) Eο = –2.36 V
Zn2+(aq) + 2e–
Zn(s)
Eο = –0.76 V
Which statement is correct?
A.
Cu2+(aq) will oxidize both Mg(s) and Zn(s).
B.
Zn(s) will reduce both Cu2+(aq) and Mg2+(aq).
C.
Mg2+(aq) will oxidize both Cu(s) and Zn(s).
D.
Cu(s) will reduce both Mg2+(aq) and Zn2+(aq).
2
7.
Consider the standard electrode potentials of the following reactions.
Cr3+(aq) + 3e– →Cr(s)
–0.75 V
Cd2+(aq) + 2e– → Cd(s)
–0.40 V
What is the value of the cell potential (in V) for the following reaction?
2Cr(s) + 3Cd2+(aq) → 2Cr3+(aq) + 3Cd(s)
8.
9.
10.
A.
–0.35
B.
–1.15
C.
+0.30
D.
+0.35
Aqueous solutions containing different concentrations of NaCl were electrolysed using platinum
electrodes. What is the major product at the positive electrode in each case?
0.001 mol dm–3 NaCl(aq)
1.0 mol dm–3 NaCl(aq)
A.
H2
Na
B.
H2
H2
C.
O2
Cl2
D.
Cl2
O2
Which is a feature of the standard hydrogen electrode?
A.
hydrogen gas at 1.01×105 Pa (1 atm) pressure
B.
1.0 mol dm–3 sulfuric acid
C.
a temperature of 273 K
D.
a magnesium electrode
Which pair of factors both affect the amount (in mol) of chlorine produced in the electrolysis of
aqueous sodium chloride?
A.
current and temperature
B.
temperature and chloride ion concentration
C.
chloride ion concentration and length of time of electrolysis
D.
pressure and length of time of electrolysis
3
11.
From the given standard electrode potentials which statement is correct?
Ca2+(aq) + 2e–
Ni2+(aq) + 2e–
Fe3+(aq) + e–
12.
13.
14.
A.
Ca2+(aq) can oxidize Ni(s)
B.
Ni2+(aq) can reduce Ca 2+(aq)
C.
Fe3+(aq) can oxidize Ni(s)
D.
Fe3+(aq) can reduce Ca 2+(aq)
Ca(s)
Ni(s)
Fe2+(aq)
EӨ = –2.87 V
EӨ = –0.23 V
EӨ = +0.77 V
Which statement is correct about the electrolysis of copper(II) sulfate solution using graphite
electrodes?
A.
A colourless gas is produced at the negative electrode.
B.
The electrolyte does not change colour.
C.
The negative electrode decreases in mass.
D.
A colourless gas is produced at the positive electrode.
A metallic object is electroplated with copper using a solution of copper(II) sulfate. Which
statement is correct?
A.
The positive electrode increases in mass.
B.
The concentration of Cu2+ ions in the solution decreases.
C.
Reduction occurs at the positive electrode.
D.
The reaction occurring at the negative electrode is Cu2+ + 2e– → Cu.
Two half-equations and their standard electrode potentials are shown in the table.
EӨ / V
Half-equation
Pb2+(aq) + 2e–
Pb(s)
–0.13
Ag2+(aq) + e–
Ag(s)
+0.80
What is the cell potential, in V, for the reaction below?
Pb(s) + 2Ag+(aq) → Pb2+(aq) + 2Ag(s)
A.
0.67
B.
0.93
C.
1.47
D.
1.73
4
15.
Two electrolytic cells are connected in series so that the same current flows through both cells
for the same length of time.
SnSO4(aq)
-
+
-
+
CuSO4 (aq)
The amount of tin deposited is 0.01 mol. How much copper is deposited?
16.
17.
A.
0.005 mol
B.
0.01 mol
C.
0.02 mol
D.
0.05 mol
Which are used for the electroplating of a metal spoon with copper?
I.
an electrolyte containing aqueous copper(II) ions
II.
a copper anode (positive electrode)
III.
a copper cathode (negative electrode)
A.
I and II only
B.
I and III only
C.
II and III only
D.
I, II and III
Consider these standard electrode potentials.
Cu2+(aq) + e– → Cu+(aq)
EӨ = +0.15 V
Cu+(aq) + e– → Cu(s)
EӨ = +0.52 V
What is the standard cell potential when the two half-cells are connected?
A.
– 0.67 V
B.
– 0.37 V
C.
+ 0.37 V
D.
+ 0.67 V
5
18.
Consider the standard electrode potentials of the following reactions:
Sn4+(aq) + 2e– → Sn2+(aq)
+0.15 V
Fe3+(aq) + e– → Fe2+(aq)
+0.77 V
What is the value of the cell potential (in volts) for the spontaneous reaction?
19.
20.
A.
+1.69
B.
+1.39
C.
+0.92
D.
+0.62
In the electrolysis of acidified water, if 8.4 cm3 of hydrogen gas is evolved, what volume of
oxygen gas is evolved?
A.
4.2 cm3
B.
8.4 cm3
C.
12.6 cm3
D.
16.8 cm3
Which factors affect the amount of metal formed during electrolysis?
I.
II.
III.
Charge on the metal ion
Current
Time
A.
I and II only
B.
I and III only
C.
II and III only
D.
I, II and III
6
21.
22.
23.
Which changes lead to the production of more moles of metal during the electrolysis of a
molten salt?
I.
using a metal ion with a higher charge
II.
increasing the current
III.
using a longer time
A.
I and II only
B.
I and III only
C.
II and III only
D.
I, II and III
Which equation represents the reduction process occurring in the standard hydrogen electrode?
A.
H2 (g) → 2H+(aq) + 2e–
B.
H+(aq) + OH– (aq) → H2 O(l)
C.
2H+(aq) + 2e– → H2 (g)
D.
O2 (g) + 4H+(aq) + 4e– → 2H2 O(l)
Which statement is correct about the value of Eο ?
A.
The more positive the value of Eο , the greater the driving force for reduction.
B.
The more negative the value of Eο , the greater the driving force for reduction.
C.
The more positive the value of Eο , the greater the rate of reaction.
D.
The more negative the value of Eο , the greater the rate of reaction.
24.
Direction of electron flow
A
Salt bridge
Zn(s)
Cu(s)
Zn 2+(aq)
Cu 2+ (aq)
7
(a)
The apparatus shown above may be used to carry out a redox reaction.
(i)
State the function of the salt bridge.
……………………………………………………………………………………
……………………………………………………………………………………
(1)
(ii)
Write a half-equation for the oxidation reaction.
……………………………………………………………………………………
……………………………………………………………………………………
(1)
(iii)
The above reactions are carried out under standard conditions.
State what the standard conditions are for the cell.
……………………………………………………………………………………
……………………………………………………………………………………
……………………………………………………………………………………
(2)
(iv)
Using the Data Booklet, calculate the cell potential for the above cell.
……………………………………………………………………………………
……………………………………………………………………………………
……………………………………………………………………………………
……………………………………………………………………………………
(2)
(v)
State and explain what happens to the concentration of the copper(II) ions when the
cell is producing an electric current.
……………………………………………………………………………………
……………………………………………………………………………………
(2)
(vi)
State two observations that could be made if the zinc rod were placed in a solution
of copper(II) ions.
……………………………………………………………………………………
……………………………………………………………………………………
(2)
8
(b)
The standard electrode potentials for three electrode systems are given below.
(i)
Ti3+(aq) + e– → Ti2+(aq)
Eο = –0.37 V
Fe3+ (aq) + e– → Fe2+(aq)
Eο = +0.77 V
Ce4+(aq) + e– → Ce3+(aq)
Eο = +1.45 V
Using the data above, deduce which species is the best reducing agent, giving a
reason in terms of electrons for your answer.
……………………………………………………………………………………
……………………………………………………………………………………
(2)
(ii)
Write an equation, including state symbols, for the overall reaction with the
greatest cell potential.
……………………………………………………………………………………
……………………………………………………………………………………
……………………………………………………………………………………
(2)
(iii)
State and explain the sign of Gο for the reaction in (b) (ii).
……………………………………………………………………………………
……………………………………………………………………………………
……………………………………………………………………………………
(2)
(c)
(i)
State the name of a solution that would produce only hydrogen and oxygen when
electrolyzed using platinum electrodes.
……………………………………………………………………………………
(1)
(ii)
Draw a diagram of apparatus that would allow the gases produced in the reaction in
(c) (i) to be collected separately. Annotate your diagram to show the polarity of
each electrode and the names and relative volumes of each gas.
………………………………………………………………………………………
………………………………………………………………………………………
………………………………………………………………………………………
………………………………………………………………………………………
………………………………………………………………………………………
………………………………………………………………………………………
(3)
(Total 20 marks)
9
25.
– +
A
X
Y
CuSO 4 (aq)
Two copper strips X and Y are placed in an aqueous solution of copper(II) sulfate and
electrolyzed for a certain time. X was then dried and weighed.
(i)
State and explain what would happen to the mass of X.
…………………………………………………………………………………………..
…………………………………………………………………………………………..
(3)
(ii)
State two ways in which the change in the mass of X could be increased.
…………………………………………………………………………………………..
(2)
(Total 5 marks)
26.
(a)
Some standard electrode potentials are shown in Table 15 of the Data Booklet.
(i)
State three conditions under which the hydrogen electrode is assigned a potential of
zero.
………………………………………………………………………………………
………………………………………………………………………………………
(3)
(ii)
Calculate the cell potential of a cell made by connecting standard copper and zinc
electrodes. State the direction of electron flow in the external circuit when the cell
produces current. Outline the changes occurring at the electrodes and in the
solutions during the process.
………………………………………………………………………………………
………………………………………………………………………………………
………………………………………………………………………………………
………………………………………………………………………………………
………………………………………………………………………………………
(5)
10
(b)
Using information from Table 15, determine whether or not there is a spontaneous
reaction between copper metal and a solution containing hydrogen ions.
……………………………………………………………………………………………
……………………………………………………………………………………………
……………………………………………………………………………………………
……………………………………………………………………………………………
(2)
(c)
Using information from Table 15, identify a substance that will oxidize bromide ions but
not chloride ions. Explain your choice, and write an equation for the redox reaction you
have chosen.
……………………………………………………………………………………………
……………………………………………………………………………………………
……………………………………………………………………………………………
……………………………………………………………………………………………
……………………………………………………………………………………………
(5)
(Total 15 marks)
27.
(a)
In one experiment involving the electrolysis of molten sodium chloride, 0.1 mol of
chlorine was formed. Deduce, giving a reason, the amount of sodium formed at the same
time.
(2)
(b)
In another experiment involving the electrolysis of molten sodium chloride, the time of
the electrolysis was halved and the current increased from 1 amp to 5 amp, compared to
the experiment in (a). Deduce the amount of chlorine formed, showing your working.
(2)
(c)
If dilute aqueous sodium chloride is electrolyzed, a different product is obtained at each
electrode. Identify the product formed at each electrode and write an equation showing its
formation.
(4)
(Total 8 marks)
28.
Define the term standard electrode potential of an element. Table 15 of the Data Booklet
contains EӨ values for two reactions involving O 2 (g). Identify the EӨ value of the reaction that
could be used to oxidize bromide ions and explain your reasoning. Construct a balanced
equation for the oxidation of bromide ions using this reaction and calculate the cell potential.
(Total 8 marks)
11
29.
Use information from Table 15 of the Data Booklet, where relevant, in answering this part.
(i)
The diagram shows the apparatus used in an experiment involving half-cells.
Voltmeter
Nickel
Solution containing
Ni 2+ (aq)
Lead
Solution containing
Pb 2+ (aq)
The reading on the voltmeter is 0.10 V. State two standard conditions that apply for this
reading to be obtained.
(2)
(ii)
The voltmeter is replaced by a 2 volt power supply, so that non-spontaneous reactions
occur. Deduce the half-equations for these reactions.
(2)
(iii)
Chlorine gas is formed when potassium manganate(VII) is added to concentrated
hydrochloric acid. Calculate the cell potential for this reaction and deduce the equation
for the reaction.
(3)
(iv)
Explain why potassium dichromate(VI) does not react with concentrated hydrochloric
acid.
(1)
(Total 8 marks)
30.
(a)
When a concentrated aqueous solution of sodium chloride is electrolyzed using inert
electrodes, a different gas is produced at each electrode.
(i)
Write equations for the oxidation and reduction half-reactions.
Oxidation half-reaction: ...................................................................................
...........................................................................................................................
Reduction half-reaction: ...................................................................................
...........................................................................................................................
(2)
(ii)
Explain why sodium is not formed during the electrolysis of aqueous NaCl
solution.
...........................................................................................................................
...........................................................................................................................
(1)
12
(b)
Deduce the products formed during the electrolysis of an aqueous solution of sodium
fluoride. Write an equation for the reaction at the positive electrode (the anode) and give
your reasoning.
......................................................................................................................................
......................................................................................................................................
......................................................................................................................................
......................................................................................................................................
......................................................................................................................................
......................................................................................................................................
......................................................................................................................................
(4)
(Total 7 marks)
31.
The following are standard electrode potentials.
Half-equation
2+
Zn (aq) + 2e–
Zn(s)
3+
–
Cr (aq) + 3e
Cr(s)
2+
–
Fe (aq) + 2e
Fe(s)
2+
–
Sn (aq) + 2e
Sn(s)
2+
–
Cu (aq) + 2e
Cu(s)
3+
–
Fe (aq) + e
Fe2+(aq)
(a)
Eο / V
–0.76
–0.74
–0.44
–0.14
+0.34
+0.77
These values were obtained using a standard hydrogen electrode. Describe the
materials and conditions used in the standard hydrogen electrode. (A suitably
labelled diagram is acceptable.)
……………………………………………………………………………………………
……………………………………………………………………………………………
……………………………………………………………………………………………
……………………………………………………………………………………………
……………………………………………………………………………………………
……………………………………………………………………………………………
……………………………………………………………………………………………
……………………………………………………………………………………………
(5)
13
(b)
Define the term oxidizing agent in terms of electron transfer and identify the strongest
oxidizing agent in the list above.
……………………………………………………………………………………………
……………………………………………………………………………………………
……………………………………………………………………………………………
(2)
(c)
A cell was set up using zinc in zinc sulfate solution and copper in copper(II) sulfate
solution, both solutions being under standard conditions.
(i)
Calculate the cell potential.
……………………………………………………………………………………
……………………………………………………………………………………
(1)
(ii)
Write an equation for the spontaneous cell reaction.
……………………………………………………………………………………
……………………………………………………………………………………
……………………………………………………………………………………
(2)
(d)
Both zinc and tin are used to coat iron to prevent it from rusting. Once the surface is
scratched, oxygen and water containing dissolved ions come into contact with the iron
and the coating metal.
(i)
State and explain whether zinc or tin would be more effective in preventing iron
from rusting under these conditions.
……………………………………………………………………………………
……………………………………………………………………………………
……………………………………………………………………………………
……………………………………………………………………………………
(2)
(ii)
Electroplating may be used to coat one metal with another metal. Identify the three
factors affecting the amount of metal discharged during electroplating.
……………………………………………………………………………………
……………………………………………………………………………………
……………………………………………………………………………………
(3)
14
(iii)
Explain why electrolysis of aqueous zinc sulfate is not used for coating with zinc
metal.
……………………………………………………………………………………
……………………………………………………………………………………
……………………………………………………………………………………
……………………………………………………………………………………
(2)
(e)
Another cell was set up as shown below.
V
Pt(s)
Cr(s)
Y
Cr 3+ (aq)
Fe 3+(aq), Fe 2+(aq)
half-cell A
(i)
half-cell B
Identify the part of the cell labelled Y and outline its function.
……………………………………………………………………………………
……………………………………………………………………………………
……………………………………………………………………………………
(2)
(ii)
Write an equation for the initial reactions at each electrode and hence write an
equation for the cell reaction.
……………………………………………………………………………………
……………………………………………………………………………………
……………………………………………………………………………………
……………………………………………………………………………………
(4)
(iii)
Describe the direction of electron flow in the external circuit.
15
……………………………………………………………………………………
(1)
(iv)
Calculate the cell potential.
……………………………………………………………………………………
……………………………………………………………………………………
(1)
(Total 25 marks)
32.
The following diagram shows a voltaic cell.
voltmeter
V
salt bridge
iron electrode
silver electrode
Fe 2+(aq)
(a)
Ag +(aq)
State an equation to represent the spontaneous reaction occurring in the cell.
....................................................................................................................................
....................................................................................................................................
(1)
(b)
Define the term standard electrode potential.
....................................................................................................................................
....................................................................................................................................
(1)
(c)
Use Table 15 from the Data Booklet to calculate the standard cell potential for the
spontaneous reaction in (a).
....................................................................................................................................
....................................................................................................................................
(1)
16
(d)
Draw arrows on the above diagram to indicate the direction of electron flow.
(1)
(Total 4 marks)
33.
(a)
Iodide ions, I– (aq), react with iodate ions, IO 3 – (aq), in an acidic solution to form
molecular iodine and water.
(i)
Determine the oxidation number of iodine in I– and in IO3 – .
.........................................................................................................................
(1)
(ii)
Identify, with a reason, the species that undergoes:
oxidation ..........................................................................................................
..........................................................................................................................
reduction ..........................................................................................................
..........................................................................................................................
(2)
(iii)
Write an ionic equation for the reaction of I– with IO3 – in an acidic solution.
..........................................................................................................................
..........................................................................................................................
(2)
(b)
Use information from Table 15 of the Data Booklet to calculate the cell potential for the
following reaction and state whether or not the reaction is spontaneous.
Cu(s) + Cu2+(aq) → 2Cu+(aq)
....................................................................................................................................
....................................................................................................................................
....................................................................................................................................
....................................................................................................................................
(3)
(Total 8 marks)
34.
A table of standard electrode potentials can be found in Table 14 of the Data Booklet.
(a)
Describe the materials and conditions used in the standard hydrogen electrode.
.....................................................................................................................................
.....................................................................................................................................
.....................................................................................................................................
.....................................................................................................................................
(5)
17
(b)
Define the term oxidizing agent in terms of electron transfer and identify the strongest
oxidizing agent in Table 14 of the Data Booklet.
.....................................................................................................................................
.....................................................................................................................................
.....................................................................................................................................
(2)
(c)
A cell was set up using tin in tin(II) sulfate solution and copper in copper(II) sulfate
solution, with both solutions under standard conditions.
(i)
Calculate the cell potential.
...........................................................................................................................
...........................................................................................................................
(1)
(ii)
Write an equation for the spontaneous cell reaction.
...........................................................................................................................
...........................................................................................................................
(2)
(Total 10 marks)
18
1.
B
[1]
2.
D
[1]
3.
C
4.
D
[1]
[1]
5.
B
[1]
6.
A
[1]
7.
D
[1]
8.
C
[1]
9.
A
[1]
10.
C
[1]
11.
C
[1]
12.
D
[1]
13.
D
[1]
14.
B
[1]
15.
B
[1]
16.
A
[1]
17.
C
[1]
18.
D
[1]
19.
A
[1]
19
20.
D
[1]
21.
C
[1]
22.
C
[1]
23.
A
[1]
24.
(a)
(i)
(c)
1
(ii)
Zn(s) → Zn2+(aq) + 2e– (state symbols not needed);
1
(iii)
298 K/25°C, 1 atm/1.01×105 Pa, 1 mol dm–3 solutions;
(all 3 for [2], 2 for [1])
2
(iv)
0.34 – (–0.76) = 1.10 V;
[1] for finding correct data, [1] for answer with unit (ECF).
2
(v)
decreases;
Cu2+ ions are converted to Cu metal/Cu deposited on electrode;
Allow ECF from (iv).
2
(vi)
(b)
ionic conductor/allows movement of ions between electrolytes
/completes circuit;
(i)
Cu deposited on Zn rod/rod goes pink/brown;
blue colour of solution → paler;
gets hotter/temperature increase/exothermic;
Ti2+ (no ECF to explanation);
Ti2+ has greatest tendency to lose electrons/Ti3+ has least tendency
to gain electrons;
(ii)
Ce4+(aq) + Ti2+(aq) → Ce3+(aq) + Ti3+(aq)
[1] for equation, [1] for state symbols. If wrong equation is
given, award [1]
for state symbols.
(iii)
Gο negative;
reaction spontaneous/corresponds to positive cell potential;
Positive [0], non-spontaneous [1].
(i)
(aqueous) sodium hydroxide/dilute sulfuric acid/sodium sulfate;
Accept correct formulas.
Any combination of K+/Na +/H+ and NO3–/SO4 2–.
Halides not acceptable. (“water” is not a solution)
2 max
2
2
2
1
(ii)
20
hydrogen / H 2
oxygen / O 2
–
+
Or similar suitable diagram.
gas collection method;
names of gases correct way round at electrodes;
2:1 volume ratio correct way round;
3
[20]
25.
(i)
(ii)
mass increases;
copper deposited;
because X is negative and attracts Cu2+ ions/reduction occurs at X/
Cu2+ + 2e– → Cu;
3
increase time;
increase current;
2
[5]
26.
(a)
(i)
(ii)
(b)
(c)
([H+]=)1 mol dm–3 ;
298 K/25°C;
1 atm/101.3 or 101 kPa
Accept 100 kPa;
3
Eο (=0.76 +0.34) = (+)1.1(0)(V);
from zinc/Zn to copper/Cu;
copper/Cu deposited/electrode becomes larger/thicker/heavier;
zinc/Zn electrode becomes smaller/thinner/lighter;
Cu2+ solution becomes paler/colourless;
Allow ECF for –1.1 V, all answers must be consistent with
the error.
no (spontaneous) reaction;
appropriate use of Table 15/Eο = –0.34 + 0.00 = –0.34 V/Eο value for
the reaction would be negative;
5
2
O2 /Cr2 O7 2– ;
Accept names.
Eο value for the reaction with Br– is positive/suitable calculation to show this;
Eο value for the reaction with Cl– is negative/Cl2 stronger oxidizing agent than
O2 /Cr2 O72– ;
5
4Br– + O2 + 4H+ → 2Br2 + 2H2 O/Cr2 O7 2– + 14H + 6Br– → 2Cr3+ + 7H2 O + 3Br2;
Award [1] for all formulas correct, [1] if coefficients correct.
[15]
21
27.
(a)
(b)
0.2 mol;
the Na:Cl2 mol ratio is 2 : 1;
0.1 12 5;
= 0.25 mol;
(c)
2
2
negative electrode
hydrogen/H2 ;
2H+ + 2e− → H2 /2H2 O + 2e− → H2 + 2OH−;
positive electrode
oxygen/O2 ;
4OH− → 2H2 O + O2 + 4e−/2H2 O → O2 + 4H+ + 4e−.
4
[8]
28.
the potential/voltage difference between the element and its ions;
(and) a hydrogen electrode;
under standard conditions/ion concentration at 1 mol dm−3 /298 K/25C;
(+)1.23 (V);
EӨ value more positive or less negative than bromine/bromide system/
EӨ value of combined half-cells is positive/OWTTE;
2Br− +
1
2
O2 + 2H+ → Br2 + H2 O;
Award [1] for all formulas correct, [1] for correct balancing.
Award [1] for correct equation reversed.
EӨ (= 1.23 − 1.09) = (+)0.14 (V);
Ignore state symbols.
8
[8]
29.
(i)
(ii)
(iii)
(iv)
(metal ions at) 1 mol dm−3 concentration;
25C/298 K;
Do not accept 1 atm pressure.
Ni2+ + 2e– →Ni;
Pb → Pb2+ + 2e−;
No penalty for using e instead of e −.
No penalty for
instead of →.
0.15 V;
10Cl− + 2Mn O 4− + 16H + → 5Cl2 + 2Mn2+ + 8H2 O;
Ignore state symbols.
Correct reactants and products = [1]
Correct balancing = [1]
not sufficiently good oxidizing agent/cell potential would have
negative EӨ value;
2
2
3
1
[8]
22
30.
(a)
(i)
(ii)
(b)
oxidation half-reaction: 2Cl– → Cl2 + 2e– ;
reduction half-reaction: 2H2 O + 2e– → H2 + 2OH– /2H+ + 2e– → H2
Award [1] only if equations are interchanged.
States not required.
2
Na has high Eο red / Na+ not readily reduced (in comparison to H 2 O)/if
formed, Na would (immediately) react with water to form Na +
H(g) and O2 (g)/accept names;
2H2 O → O2 + 4H+ + 4e– /4OH– → 2H2 O + O2 + 4e– ;
water is oxidized (instead of the halide);
since Eο OX for F– is very negative/Eο red for F2 is very high;
Accept answer based on oxidizing/reducing strengths.
1
4
[7}
31.
(a)
Hydrogen gas at 1 atm
1 mol dm–3 [H +(aq)]
Platinum (coated
with platinum black)
Accept suitable diagram with the following indicated:
Pt electrode;
1 mol dm–3 [H+(aq)];
H2 gas;
at 1 atm/1.01×105 Pa;
298 K/25°C;
(b)
(c)
(d)
5
electron acceptor;
Fe3+(aq)/iron(III) ions/Fe 3+;
Do not accept iron/Fe 2+/iron ion.
(i)
(+)1.10;
(ii)
Cu2+(aq) + Zn(s) → Zn2+(aq) + Cu(s);
Award [1] for correct reactants and products from (c)(i),
and [1] for state symbols.
(i)
zinc;
zinc is more readily oxidised than iron and so protects it by reacting
preferentially/OWTTE or tin is less readily oxidised than iron and
so iron reacts preferentially/OWTTE;
2
1
2
2
23
(ii)
(e)
charge on the ion discharged;
size/magnitude of the current;
time/duration of the electrolysis;
3
(iii)
positive ions/cations in solution = H +(aq), Zn2+(aq);
H+(aq) discharged preferentially;
(i)
salt bridge;
allows movement of ions between the solutions/to complete
the circuit/to maintain electrical neutrality;
(ii)
2
A: Fe 3+(aq) + e– → Fe2+(aq);
B: Cr(s) → Cr3+(aq) + 3e– ;
Allow [2] for correct equation for the cell reaction if
equations for A and B are reversed.
3 Fe3+(aq) + Cr(s) → 3 Fe2+(aq) + Cr3+(aq);
Award [1] for correct reactants and products, with state
symbols, and [1] for correct balancing.
4
(iii)
from B to A/from Cr to Pt/from right to left;
Allow ECF from (ii).
1
(iv)
(+)1.51;
Allow ECF from (ii).
1
[25]
32.
(a)
Fe + 2Ag+ → Fe2+ + 2Ag;
Ignore state symbols.
1
Accept Fe + 3Ag + → Fe3+ + 3Ag
(b)
the potential difference/EMF/Voltage between a standard half-cell and
standard hydrogen electrode/OWTTE;
1
(c)
(+) 1.24 (V);
ECF from (a).
1
(d)
electron flow indicated on wires;
ECF from (a).
1
voltmeter
V
salt bridge
iron electrode
silver electrode
Fe 2+(aq)
Ag +(aq)
24
[4]
33.
(a)
(i)
(ii)
I− = −1/1− and IO3 − = +5/5+;
Both answers needed for [1] mark,
Signs must be included
Do not accept Roman numerals
1
oxidation
I− (to I2 ), increase in oxidation number/loss of electron(s);
reduction
IO3 – (to I2 ), decrease in oxidation number/gain of electron(s);
(iii)
5I−(aq) + IO3 – (aq) + 6H+(aq) → 3I2 (aq) + 3H2 O(l);
Award [2] if correctly balanced
Award [1] for correct reactants and products.
States not required for mark.
2
2
[1 max] if HCl on left and Cl− on right side
(b)
Cu2+(aq) + e−
Cu+(aq) : EӨ = (+) 0.15(V), Cu(s)
Cu+(aq) + e− :
EӨ = − 0.52(V)/(+) 0.15 and − 0.52(V);
No mark if 0.34 or 0.52 quoted, but then ECF
(Cu2+(aq) + Cu(aq)
2 Cu+(aq)) EӨ value for reaction = − 0.37(V);
Award [2] for correct Erxn Ө even if equations are not given,
states not required.
(negative value means) not spontaneous;
Allow ECF: if positive value, can score third mark for
‘spontaneous’.
3
[8]
34.
(a)
(b)
(c)
Pt electrode;
1 mol dm–3 [H+(aq)];
H2 gas;
at 1 atm/1.01×105 Pa;
298 K/25°C;
Accept suitable labelled diagram with the above.
5
electron acceptor;
F2 /fluorine;
2
(i)
(+)0.48 (V);
1
(ii)
Cu2+(aq) + Sn(s) → Sn2+(aq) + Cu(s)
Award [1] for correct reactants and products from (c)(i), and
[1] for state symbols.
2
[10]
25
IB Chemistry – SL
Topic 10 Questions
1.
Which of the structures below is an aldehyde?
O
A.
CH 3 CH 2CH
O
B.
CH 3 CCH 3
O
C.
CH 3 CH 2COH
O
D.
CH 3 COCH3
(Total 1 mark)
2.
What product results from the reaction of CH 2 ==CH2 with Br2 ?
A.
CHBrCHBr
B.
CH2 CHBr
C.
CH3 CH2 Br
D.
CH2 BrCH2 Br
(Total 1 mark)
3.
What is the final product formed when CH 3 CH2 OH is refluxed with acidified potassium
dichromate(VI)?
A.
CH3 CHO
B.
CH2 ==CH2
C.
CH3 COOH
D.
HCOOCH3
(Total 1 mark)
4.
Which substance(s) could be formed during the incomplete combustion of a hydrocarbon?
I.
Carbon
A.
I only
B.
I and II only
C.
I and III only
D.
II and III only
II.
Hydrogen
III.
Carbon monoxide
(Total 1 mark)
1
5.
Which formulas represent butane or its isomer?
I.
CH3 (CH2 )2 CH3
II.
CH3 CH(CH3 )CH3
III.
(CH3 )3 CH
A.
I and II only
B.
I and III only
C.
II and III only
D.
I, II and III
(Total 1 mark)
6.
Which statement about neighbouring members of all homologous series is correct?
A.
They have the same empirical formula.
B.
They differ by a CH2 group.
C.
They possess different functional groups.
D.
They differ in their degree of unsaturation.
(Total 1 mark)
7. What is the IUPAC name for CH 3 CH2 CH(CH3)2?
A.
1,1-dimethylpropane
B.
2-methylbutane
C.
isopentane
D.
ethyldimethylmethane
(Total 1 mark)
8.
Which compound has the lowest boiling point?
A.
CH3 CH2 CH(CH3)CH3
B.
(CH3 )4 C
C.
CH3 CH2 CH2CH2CH3
D.
CH3 CH2 OCH2 CH3
(Total 1 mark)
2
9.
What type of reaction does the equation below represent?
CH2 =CH2 + Br2 → BrCH2 CH2Br
A.
substitution
B.
condensation
C.
reduction
D.
addition
(Total 1 mark)
10.
Which compound is a member of the same homologous series as 1-chloropropane?
A.
1-chloropropene
B.
1-chlorobutane
C.
1-bromopropane
D.
1,1-dichloropropane
(Total 1 mark)
11.
Which formula is a correct representation of pentane?
A.
CH3 CH2 CHCH2 CH3
B.
(CH3 CH2 )2 CH3
C.
CH3 (CH2 )3 CH3
D.
CH3 (CH3 )3 CH3
(Total 1 mark)
12.
How many structural isomers are possible with the molecular formula C 6 H14 ?
A.
4
B.
5
C.
6
D.
7
(Total 1 mark)
13.
Which compound is a member of the aldehyde homologous series?
A.
CH3 COCH3
B.
CH3 CH2 CH2OH
C.
CH3 CH2 COOH
D.
CH3 CH2 CHO
(Total 1 mark)
3
14.
Which type of compound can be made in one step from a secondary alcohol?
A.
an aldehyde
B.
an alkane
C.
a carboxylic acid
D.
a ketone
(Total 1 mark)
15.
Which formula represents a tertiary alcohol?
A.
CH 3
CH
CH 2
CH 3
B.
CH3
CH 2OH
C.
C
D.
CH 2
CH 2
CH 2
OH
CH
CH3
CH 3
CH 3
CH 3
CH
CH 3
OH
CH 3
CH3
CH
OH
(Total 1 mark)
16.
Which reaction type is typical for halogenoalkanes?
A.
nucleophilic substitution
B.
electrophilic substitution
C.
electrophilic addition
D.
nucleophilic addition
(Total 1 mark)
17.
Which substance is not readily oxidized by acidified potassium dichromate(VI) solution?
A.
propan-1-ol
B.
propan-2-ol
C.
propanal
D.
propanone
(Total 1 mark)
4
18.
What is the correct name of this compound?
CH3
CH3
CH
CH2
CH 2
CH 3
A.
1,3-dimethylbutane
B.
2,4-dimethylbutane
C.
2-methylbutane
D.
2-methylpentane
(Total 1 mark)
19.
Propane, C3 H8 , undergoes incomplete combustion in a limited amount of air. Which products
are most likely to be formed during this reaction?
A.
Carbon monoxide and water
B.
Carbon monoxide and hydrogen
C.
Carbon dioxide and hydrogen
D.
Carbon dioxide and water
(Total 1 mark)
20.
What is/are the product(s) of the reaction between ethene and hydrogen bromide?
A.
CH3 CH2 Br
B.
CH3 CH2 Br and H2
C.
CH2 BrCH2 Br
D.
CH3 BrCH2 Br and H2
(Total 1 mark)
21.
Which are characteristics typical of a free radical?
I.
II.
III.
It has a lone pair of electrons.
It can be formed by the homolytic fission of a covalent bond.
It is uncharged.
A.
I and II only
B.
I and III only
C.
II and III only
D.
I, II and III
(Total 1 mark)
5
22.
Which of the following products could be formed from the oxidation of ethanol?
I.
ethanal
II.
ethanoic acid
III.
ethane
A.
I and II only
B.
I and III only
C.
II and III only
D.
I, II and III
(Total 1 mark)
23.
What is the reaction type when (CH 3 )3 CBr reacts with aqueous sodium hydroxide to form
(CH3 )3 COH and NaBr?
A.
Addition
B.
Elimination
C.
SN1
D.
SN2
(Total 1 mark)
24.
Which species is a free radical?
A.
•CH3
B.
+
C.
D.
CH3
–
CH3
:CH3
(Total 1 mark)
25.
Which compound is a tertiary halogenoalkane?
A.
(CH3 CH2 )2 CHBr
B.
CH3 (CH2 )3 CH2Br
C.
(CH3 )2 CHCH2CH2 Br
D.
CH3 CH2 C(CH3 )2Br
(Total 1 mark)
6
26.
Which species reacts most readily with propane?
A.
Br2
B.
Br•
C.
Br–
D.
Br+
(Total 1 mark)
27.
An organic compound X reacts with excess acidified potassium dichromate(VI) to form
compound Y, which reacts with sodium carbonate to produce CO 2 (g).
What is a possible formula for compound X?
A.
CH3 CH2 COOH
B.
CH3 CH2 CH2OH
C.
CH3 CH(OH)CH3
D.
(CH3 )3 COH
(Total 1 mark)
28.
Which statement about successive members of all homologous series is correct?
A.
They have the same empirical formula.
B.
They differ by a CH2 group.
C.
They have the same physical properties.
D.
They differ in their degree of unsaturation.
(Total 1 mark)
29.
The following is a three-dimensional representation of an organic molecule.
7
Which statement is correct?
A.
The correct IUPAC name of the molecule is 2-methylpentane.
B.
All the bond angles will be approximately 90°.
C.
One isomer of this molecule is pentane.
D.
The boiling point of this compound would be higher than that of pentane.
(Total 1 mark)
30.
Which compound forms when hydrogen bromide is added to but-2-ene?
A.
2-bromobutane
B.
2,3-dibromobutane
C.
1-bromobutane
D.
1,2-dibromobutane
(Total 1 mark)
31.
Which products can be potentially obtained from crude oil and are economically important?
I.
II.
III.
Plastics
Margarine
Motor fuel
A.
I and II only
B.
I and III only
C.
II and III only
D.
I, II and III
(Total 1 mark)
32.
Propane, C3 H8 , undergoes incomplete combustion in a limited amount of air. Which products
are most likely to be formed during this reaction?
A.
Carbon monoxide and water
B.
Carbon monoxide and hydrogen
C.
Carbon dioxide and hydrogen
D.
Carbon dioxide and water
(Total 1 mark)
8
33.
Two reactions of an alkene, B, are shown below.
C 4 H 10
A
H2
CH 3
CH 3
Ni
C
C
C
H
C 4 H 9 Br
H
B
(i)
State the name of A and write an equation for its complete combustion. Explain why the
incomplete combustion of A is dangerous.
…………………………………………………………………………………………..
…………………………………………………………………………………………..
…………………………………………………………………………………………..
…………………………………………………………………………………………..
…………………………………………………………………………………………..
(5)
(ii)
Outline a test to distinguish between A and B, stating the result in each case.
…………………………………………………………………………………………..
…………………………………………………………………………………………..
…………………………………………………………………………………………..
…………………………………………………………………………………………..
(3)
(iii)
Write an equation for the conversion of B to C. State the type of reaction taking place and
draw the structure of C.
…………………………………………………………………………………………..
…………………………………………………………………………………………..
…………………………………………………………………………………………..
…………………………………………………………………………………………..
…………………………………………………………………………………………..
…………………………………………………………………………………………..
…………………………………………………………………………………………..
…………………………………………………………………………………………..
(3)
(Total 11 marks)
9
34.
For the two compounds HCOOCH 2 CH3 and HCOOCHCH2:
I
II
(i)
State and explain which of the two compounds can react readily with bromine.
…………………………………………………………………………………………..
…………………………………………………………………………………………..
…………………………………………………………………………………………..
(2)
(ii)
Compound II can form polymers. State the type of polymerization compound II
undergoes, and draw the structure of the repeating unit of the polymer.
(2)
(Total 4 marks)
35.
The compound C2 H4 can be used as a starting material for the preparation of many substances.
(a)
Name the compound C2 H4 and draw its structural formula.
……………………………………………………………………………………………
……………………………………………………………………………………………
……………………………………………………………………………………………
……………………………………………………………………………………………
(2)
(b)
In the scheme below, state the type of reaction and identify the reagent needed for each
reaction.
A
B
C2 H4 → CH3 CH2 OH → CH3COOH
……………………………………………………………………………………………
……………………………………………………………………………………………
……………………………………………………………………………………………
……………………………………………………………………………………………
(4)
(c)
C2 H4 can be converted into one of the compounds below in a single step reaction.
C2 H3 Cl
C2 H4 Cl2
Draw the structural formula for each of these compounds and identify the compound
which can be formed directly from C2 H4 .
……………………………………………………………………………………………
……………………………………………………………………………………………
……………………………………………………………………………………………
(3)
10
(d)
One of the two compounds in (c) has an isomer. Draw the structural formula of the
isomer and explain why it can not be formed directly from C 2 H4 .
……………………………………………………………………………………………
……………………………………………………………………………………………
……………………………………………………………………………………………
……………………………………………………………………………………………
……………………………………………………………………………………………
……………………………………………………………………………………………
(2)
(e)
C2 H4 can also react to form a polymer. Name this type of polymer and draw the
structural formula of a section of this polymer consisting of three repeating units.
……………………………………………………………………………………………
……………………………………………………………………………………………
……………………………………………………………………………………………
……………………………………………………………………………………………
……………………………………………………………………………………………
……………………………………………………………………………………………
(2)
(f)
Polymers can also be formed in a different type of reaction. Identify this type of reaction
and name two different types of such polymers.
……………………………………………………………………………………………
……………………………………………………………………………………………
……………………………………………………………………………………………
……………………………………………………………………………………………
(3)
(Total 16 marks)
36.
The plastic PVC, poly(chloroethene), is made from the monomer chloroethene, C 2 H3 Cl, by a
polymerization reaction.
(i)
Draw the structural formula of chloroethene.
……………………………………………………………………………………………
……………………………………………………………………………………………
……………………………………………………………………………………………
……………………………………………………………………………………………
(1)
11
(ii)
State the type of polymerization reaction that occurs to make poly(chloroethene) and
identify the structural feature needed in the monomer.
……………………………………………………………………………………………
……………………………………………………………………………………………
(2)
(iii)
Draw the structure of the repeating unit of poly(chloroethene).
……………………………………………………………………………………………
……………………………………………………………………………………………
……………………………………………………………………………………………
……………………………………………………………………………………………
(1)
(iv)
Explain why monomers are often gases or volatile liquids, whereas polymers are solids.
……………………………………………………………………………………………
……………………………………………………………………………………………
……………………………………………………………………………………………
……………………………………………………………………………………………
(2)
(Total 6 marks)
37.
The hydrolysis of 2-iodo-2-methylpropane by 0.10 mol dm–3 KOH(aq) to form
2-methylpropan-2-ol is an example of nucleophilic substitution.
Give equations to illustrate the S N1 mechanism for this reaction.
..............................................................................................................................................
..............................................................................................................................................
..............................................................................................................................................
..............................................................................................................................................
(Total 2 marks)
38.
The molecular formula C4 H9 Br represents four structural isomers, all of which can undergo
nucleophilic substitution reactions with aqueous sodium hydroxide. An equation to represent all
these reactions is
C4 H9 Br + NaOH → C4 H9 OH + NaBr
(a)
Explain what is meant by the term nucleophilic substitution.
(2)
(b)
The main mechanism for a tertiary halogenoalkane is SN1. Give the equations for this
substitution reaction of the tertiary isomer of C4 H9 Br. Show the structures of the organic
reactant and product and use curly arrows to show the movement of electron pairs.
(4)
12
(c)
The main mechanism for a primary halogenoalkane is SN2. Give the mechanistic equation
for this substitution reaction of the straight-chain primary isomer of C4 H9 Br, showing the
structures of the organic reactant and product, and using curly arrows to show the
movement of electron pairs.
(4)
(d)
Give a structural formula for the secondary isomer and for the other primary isomer. State
the name of each isomer.
(4)
(Total 14 marks)
39.
Write equations to show the mechanisms of the following reactions. In each case, show the
structure of the intermediate and organic product, and use curly arrows to show the movement
of electron pairs.
(i)
the reaction between KOH and CH 3 CH2 CH2 CH2Cl.
(3)
(ii)
the reaction between KOH and (CH 3 )3 CCl.
(2)
(Total 5 marks)
40.
Some alcohols are oxidized by heating with acidified potassium dichromate(VI). If oxidation
does occur, identify the possible oxidation products formed by each of the alcohols below.
Indicate if no oxidation occurs.
Butan-1-ol
..............................................................................................................................................
..............................................................................................................................................
Butan-2-ol
..............................................................................................................................................
..............................................................................................................................................
2-methylpropan-2-ol
..............................................................................................................................................
..............................................................................................................................................
(Total 4 marks)
13
41.
Chlorine and ethane react together to form chloroethane.
(a)
State the condition needed for the reaction to occur.
....................................................................................................................................
....................................................................................................................................
(1)
(b)
Write equations to represent initiation, propagation and termination steps in the reaction.
....................................................................................................................................
....................................................................................................................................
....................................................................................................................................
....................................................................................................................................
....................................................................................................................................
....................................................................................................................................
....................................................................................................................................
....................................................................................................................................
....................................................................................................................................
....................................................................................................................................
(4)
(Total 5 marks)
42.
CH3 COCH3 is the first member of the ketone homologous series. Draw the full structural
formula of the next member of this homologous series and predict how its melting point
compares with that of CH3 COCH3 .
(Total 2 marks)
43.
(i)
Write an equation for the reaction between but-2-ene and bromine, showing the structure
of the organic product.
...................................................................................................................................
...................................................................................................................................
...................................................................................................................................
...................................................................................................................................
(2)
(ii)
State the type of reaction occurring.
.........................................................................................................................
.........................................................................................................................
(1)
(Total 3 marks)
14
44.
CH3 COCH3 can be prepared in the laboratory from an alcohol. State the name of this alcohol,
the type of reaction occurring and the reagents and conditions needed for the reaction.
(Total 5 marks)
45.
2-bromobutane can be converted into butan-2-ol by a nucleophilic substitution reaction. This
reaction occurs by two different mechanisms.
(i)
Give the structure of the transition state formed in the S N2 mechanism.
(2)
(ii)
Write equations for the S N1 mechanism.
(2)
(Total 4 marks)
46.
Ethene is an unsaturated hydrocarbon used as a starting material for many organic chemicals.
(a)
Draw the structural formula of ethene and state the meaning of the term unsaturated
hydrocarbon.
(2)
(b)
State an equation for the conversion of ethene to ethanol and identify the type of reaction.
(2)
(c)
Describe the complete oxidation of ethanol and name the product. Include the conditions,
reagents required and any colour changes.
(4)
(d)
State an equation for the reaction between ethanol and the product of complete oxidation
in (c). Include any other reagent required for this reaction. Name the organic product and
state one possible use of this product.
(4)
(Total 12 marks)
47.
Ethene is an unsaturated hydrocarbon used as a starting material for many organic chemicals.
(a)
State the meaning of the term unsaturated hydrocarbon.
(1)
(b)
State an equation for the conversion of ethene to ethanol and identify the type of reaction.
(2)
(c)
Describe the complete oxidation of ethanol. Include the conditions, reagents required and
any colour changes. Name the organic product X.
(4)
(d)
State an equation for the reaction between ethanol and compound X. Include any other
reagent required. Name the organic compound Z and state one use of this product.
(4)
(Total 11 marks)
15
48.
The equation for a reaction of ethane is
CH3 CH3 + Cl2 → CH3 CH2 Cl + HCl
The mechanism of this reaction involves initiation, propagation and termination steps. Describe
this reaction, including equations for each step and the role of ultraviolet light.
..............................................................................................................................................
..............................................................................................................................................
..............................................................................................................................................
..............................................................................................................................................
..............................................................................................................................................
..............................................................................................................................................
..............................................................................................................................................
..............................................................................................................................................
..............................................................................................................................................
..............................................................................................................................................
(Total 5 marks)
49.
(i)
Draw the structural formula of propan-2-ol.
(1)
(ii)
Identify the alcohol as primary, secondary, or tertiary.
....................................................................................................................................
(1)
(iii)
Identify the organic product formed by the oxidation of this alcohol using acidified
potassium dichromate(VI) solution.
..........................................................................................................................
(1)
(Total 3 marks)
16
50.
Secondary halogenoalkanes can undergo nucleophilic substitution reactions by both S N1 and
SN2 mechanisms. The mechanism showing the formation of the transition state in the reaction
between 2-bromobutane and potassium hydroxide can be represented as follows.
CH 2 CH 3
H
(a)
C
Br
CH3
:OH –
CH 2 CH 3
H
HO
C
Br
CH
Identify the type of mechanism shown.
....................................................................................................................................
(1)
(b)
State and explain how the following changes would alter the rate of the reaction by this
mechanism.
(i)
using water instead of potassium hydroxide.
..........................................................................................................................
..........................................................................................................................
..........................................................................................................................
..........................................................................................................................
(2)
(ii)
using bromoethane instead of 2-bromobuture.
..........................................................................................................................
..........................................................................................................................
..........................................................................................................................
..........................................................................................................................
(2)
(Total 5 marks)
17
51.
The following is a computer-generated representation of the molecule, methyl
2-hydroxy benzoate, better known as oil of wintergreen.
(i)
Deduce the empirical formula of methyl 2-hydroxy benzoate and draw the full structural
formula, including any multiple bonds that may be present.
The computer-generated representation shown does not distinguish between single and
multiple bonds.
(2)
(ii)
In this representation, two of the carbon-oxygen bond lengths shown are 0.1424 nm and
0.1373 nm. Explain why these are different and predict the carbon-oxygen bond length in
carbon dioxide.
(2)
(iii)
Name all the functional groups present in the molecule.
(2)
(Total 6 marks)
52.
(i)
State and explain the trend in the boiling points of the first six alkanes involving
straight-chains.
(2)
(ii)
Write an equation for the reaction between methane and chlorine to form chloromethane.
Explain this reaction in terms of a free-radical mechanism.
(5)
(Total 7 marks)
53.
(i)
Identify the formulas of the organic products, A–E, formed in the reactions, I–IV:
+
+
I.
H
H
⎯ A ⎯⎯→
⎯ B
CH3 (CH2 )8 OH + K2 Cr2O7 ⎯⎯→
II.
(CH3 )3 CBr + NaOH ⎯⎯→ C
III.
H
⎯ D
(CH3 )2 CHOH + K2 Cr2 O7 ⎯⎯→
+
18
IV.
H2 C=CH2 + Br2 ⎯⎯→ E
(5)
(ii)
H2 C=CH2 can react to form a polymer. Name this type of polymer and draw the
structural formula of a section of this polymer consisting of three repeating units.
(2)
(Total 7 marks)
54.
Ethene, propene and but-2-ene are members of the alkene homologous series.
(a)
Describe three features of members of a homologous series.
(3)
(b)
State and explain which compound has the highest boiling point.
(3)
(c)
Draw the structural formula and give the name of an alkene containing five carbon atoms.
(2)
(d)
Write an equation for the reaction between but-2-ene and hydrogen bromide, showing the
structure of the organic product. State the type of reaction occurring.
(3)
(e)
Propene can be converted to propanoic acid in three steps:
step1
propene
step 2
propan-1-ol
step 3
propanal
propanoic acid
State the type of reaction occurring in steps 2 and 3 and the reagents needed. Describe
how the conditions of the reaction can be altered to obtain the maximum amount of
propanal, and in a separate experiment, to obtain the maximum amount of propanoic acid.
(5)
(f)
Identify the strongest type of intermolecular force present in each of the compounds
propan-1-ol, propanal and propanoic acid. List these compounds in decreasing order of
boiling point.
(4)
(Total 20 marks)
55.
(a)
State two characteristics of a homologous series.
....................................................................................................................................
....................................................................................................................................
....................................................................................................................................
(2)
(b)
Describe a chemical test to distinguish between alkanes and alkenes, giving the result in
each case.
....................................................................................................................................
....................................................................................................................................
....................................................................................................................................
(3)
(Total 5 marks)
19
56.
The following transition state is formed during the reaction of a halogenoalkane with aqueous
sodium hydroxide:
CH3
H 3C
HO
C
Br
H
(a)
Deduce the structure of the halogenoalkane. Classify it as primary, secondary or tertiary,
giving a reason for your choice.
....................................................................................................................................
....................................................................................................................................
(2)
(b)
The mechanism of this reaction is described as SN2. Explain what is meant by the
symbols in SN2. Predict a rate expression for this reaction.
....................................................................................................................................
....................................................................................................................................
....................................................................................................................................
....................................................................................................................................
....................................................................................................................................
....................................................................................................................................
(3)
(c)
The same halogenoalkane reacts with sodium hydroxide by an SN1 mechanism. Deduce
the structure of the intermediate formed in this reaction.
....................................................................................................................................
....................................................................................................................................
....................................................................................................................................
....................................................................................................................................
(1)
(Total 6 marks)
57.
(a)
List two characteristics of a homologous series.
.....................................................................................................................................
.....................................................................................................................................
.....................................................................................................................................
(1)
20
(b)
Ethanol and ethanoic acid can be distinguished by their melting points. State and explain
which of the two compounds will have a higher melting point.
.....................................................................................................................................
.....................................................................................................................................
.....................................................................................................................................
.....................................................................................................................................
(2)
(c)
Draw the three isomers containing the alcohol functional group of formula C 4 H9 OH.
(2)
(Total 5 marks)
21
IB Chemistry – SL
Topic 10 Answers
1.
A
[1]
2.
D
[1]
3.
C
[1]
4.
C
[1]
5.
D
[1]
6.
B
[1]
7.
B
[1]
8.
B
[1]
9.
D
[1]
10.
B
[1]
11.
C
[1]
12.
B
[1]
13.
D
[1]
14.
D
[1]
15.
C
[1]
16.
A
[1]
17.
D
[1]
18.
D
[1]
19.
A
[1]
20.
A
[1]
21.
C
[1]
22
22.
A
[1]
23.
C
[1]
24.
A
[1]
25.
D
[1]
26.
B
[1]
27.
B
[1]
28.
B
[1]
29.
C
[1]
30.
A
[1]
31.
B
[1]
32.
A
[1]
33.
(i)
butane;
C4 H10 (g) + 132 O2 (g) → 4CO2 (g) + 5H2 O(l);
(ignore state symbols, accept balancing using 13O 2 )
[1] for all formulas and [1] for balancing equation.
CO produced;
CO is poisonous/combines with hemoglobin/OWTTE;
or
C;
which causes respiratory problems;
(ii)
add Br2 (water);
valid test needed to score further marks.
A – no effect;
B – would decolorise Br2 (do not accept discolour);
(iii)
5
CH3 CH==CHCH3 + HBr → CH3 CHBrCH2CH3;
[1] for HBr in balanced equation, [1] for structure of product.
3
3
addition;
[11]
23
34.
(i)
(ii)
II reacts with Br2
II is an alkene/has unsaturated R group/C
only saturated R groups;
C present, I contains
2
addition polymerization;
(
CH
CH 2
)
(
CH
O
O
C
CHO
CH 2
)
O
H
accept
2
[4]
35.
(a)
ethene;
2
H
H
C
H
(b)
/ CH2
C
CH2
H
A addition/hydration;
H2 O/water/steam;
B oxidation;
acidified K2 Cr2 O7
Accept acidified KMnO4 .
4
(c)
H
H
C
C
H
/ CH2
CHCl
Cl
Cl H
H
C
C
H
H
Cl / CH 2 ClCH 2 Cl ;
The compound formed directly may be circled or indicated by some
other means. Accept any other structure showing a Cl atom on each C atom.
3
(d)
H
H
Cl
C
C
H
H
Cl / CH3 CHCl 2;
addition across a double bond occurs at both C atoms/OWTTE;
If 1,1-dichloroethane is given in (c) accept 1,2-dichloroethane
as the isomer as ECF but Award [1] max;
2
24
(e)
(f)
addition polymer;
2
H
H
H
H
H
H
C
C
C
C
C
C
H
H
H
H
H
H
/
( CH2
allow
CH 2 )3 ;
( CH2 ) 3
condensation polymer;
polyesters;
polyamides;
3
[16]
H
36.
(i)
H
C
CH2 CHCl/CH2 = CHCl/
C
H
(ii)
(iii)
(iv)
;
1
Cl
addition (polymerization);
(carbon-carbon) double bond/unsaturation/OWTTE;
H
H
C
C
H
Cl
2
;
1
monomers have smaller molecules/surface area than polymers;
with weaker intermolecular/Van der Waals’ forces;
Accept opposite argument for polymers.
2
[6]
37.
(CH3 )3 CI → (CH3 )3C+ + I−;
(CH3 )C+ + OH− → (CH3 )3 COH;
Do not allow S N 2 reaction.
2
[2]
38.
(a)
(b)
replacement of atom/group (in a molecule)/OWTTE;
Do not accept substitution.
by a species with a lone pair of electrons/species attracted to an
electron-deficient carbon atom;
correct structure of (CH3 )3 CBr;
curly arrow showing C⎯Br bond fission;
correct structure of (CH3 )3 C+;
curly arrow showing attack by OH − on correct C atom;
correct structure of (CH3 )3 COH;
Award [1] each for any four.
(c)
2
correct structure of CH3 CH2 CH2CH2Br;
curly arrow showing C⎯Br bond fission;
correct structure of transition state showing charge and all bonds;
curly arrow showing attack by OH − on correct C atom;
correct structure of CH3 CH2 CH2CH2OH;
Award [1] each for any four.
4
4
25
(d)
secondary
CH3 CHBrCH2CH3;
2-bromobutane;
other primary
(CH3 )2 CHCH2Br;
1-bromo-2-methylpropane;
4
[14]
39.
(i)
(SN2 mechanism)
3
Cl
CH 2
CH 2
CH 2
CH 3
;
–
HO
curly arrow must start from O or negative charge
Cl
HO
CH2
CH 2
CH 2
;
CH 3
Intermediate structure showing overall negative charge and
partial bonds.
Accept negative charge to be indicated as delocalised between
the HO-CH2 -Cl.
→ HO − CH2 CH2 CH2CH3 + Cl−;
(ii)
(SN1 mechanism)
2
CH 3
CH 3
CH3
C
CH 3
Cl
C
+ Cl -
;
CH 3
CH 3
formation of carbocation / loss of Cl -
CH 3
CH 3
CH 3
C
+ OH -
CH 3
CH 3
C
OH
;
CH 3
carbocation + OH [5]
26
40.
butan-1-ol:
butanal;
butanoic acid;
butan-2-ol:
butanone;
2 methylpropan-2-ol:
no oxidation;
Also accept correct structures. Where both name and structure
given structure must be correct and name largely correct.
[4]
41.
(a)
UV light/sunlight (present);
1
(b)
Throughout accept radical with or without •
initiation reaction(s):
Cl2 → 2Cl•;
1
propagation reactions:
Cl• + CH3 CH3 → CH3 CH2• + HCl;
CH3 CH2 • + Cl2 → CH3 CH2 Cl + Cl•;
2
termination reactions:
CH3 CH2 • + Cl• → CH3 CH2 Cl;
2Cl• → Cl2 ;
2CH3 CH2 • → CH3 CH2CH2CH3 ;
Award [1] for any termination reaction.
If initiation, propagation, termination not labelled or
incorrectly labelled award [3] max.
1
[5]
42.
H
higher;
H
H
C
C
C
C
H
H
O
H
H
H;
[2]
43.
(i)
(ii)
C4 H8 + Br2 → C4 H8 Br2 ;
Equation scores [1].
CH3 CHBrCHBrCH3 ;
Accept more detailed formula.
2
addition;
1
[3]
44.
propan-2-ol;
Accept 2-propanol.
oxidation/redox;
(potassium/sodium) dichromate(VI)/potassium manganate(VII);
27
Accept just dichromate, permanganate, KMnO4 , Mn O4− , K2 Cr2 O7 , Cr2 O72− .
(sulfuric) acid;
heat under reflux;
[5]
45.
(i)
CH 2 CH3
HO
C
H
(ii)
Br
CH3
all five groups around C correct;
negative charge and dotted lines to OH and Br correct;
Do not award 2 nd mark if bond from OH (i.e. OH-----).
2
CH3 CH2 CH(CH3)Br → CH3 CH2CH(CH3)+ + Br−;
CH3 CH2 CH(CH3)+ + OH− → CH3 CH2CH(CH3)OH;
Accept C4 H9 instead of CH3CH2CH(CH3) throughout.
2
[4]
46.
(a)
H
H
C
H
;
C
H
Allow CH2 =CH2 .
a hydrocarbon that contains at least one C=C (or CC)/carbon-carbon
double bond (or triple bond)/carbon to carbon multiple bond;
Do not accept just “double bond”.
(b)
(c)
C2 H4 + H2 O → C2 H5 OH;
addition/hydration reaction;
2
2
heat under reflux;
EITHER
potassium dichromate(VI)/K2 Cr2 O7 /Cr2O72–and acidified/H
orange to green;
+
;
OR
potassium permanganate/manganate(VII)/KMnO 4 /MnO4 – and
acidified/H +;
purple to colourless;
Penalize wrong oxidation state, but not missing oxidation state.
ethanoic acid;
4
28
(d)
CH3 COOH + C2 H5 OH → CH3 COOCH2CH3 + H2 O;
Accept CH3 COOC2H5
sulfuric acid/H2 SO4 /(ortho)phosphoric acid/H3 PO4 ;
ethyl ethanoate;
solvent/flavouring/perfumes/plasticizers;.
4
[12]
47.
(a)
(b)
(c)
a hydrocarbon that contains at least one C=C (or CC)/carbon-carbon
double bond (or triple bond)/carbon to carbon multiple bond;
Do not accept just “double bond”.
C2 H4 + H2 O → C2 H5 OH;
addition/hydration reaction;
1
2
heat under reflux;
EITHER
potassium dichromate(VI)/K2 Cr2 O7 / Cr2 O7 2– and acidified/H+;
orange to green;
OR
potassium permanganate/manganate(VII)/KMnO 4 / MnO4 – and acidified/H+;
purple to colourless;
Penalize wrong oxidation state, but not missing oxidation state.
ethanoic acid;
(d)
4
CH3 COOH + C2 H5 OH
CH3 COOCH2CH3 + H2 O;
accept equations including H +.
Reversible arrow not required for the mark.
sulfuric acid/H2 SO4 /(ortho)phosphoric acid/H3 PO4 ;
Z – ethyl ethanoate;
solvent/flavouring/perfumes/plasticizers;
4
[11]
48.
ultraviolet light causes Cl−Cl bond to split;
Cl2 → 2Cl•;
Cl• + CH3 CH3 → CH3 CH2• + HCl;
CH3 CH2 • + Cl2 → CH3 CH2 Cl + Cl•
CH3 CH2 • + Cl• → CH3 CH2 Cl/other correct termination step;
Penalize missing • symbol once only.
If different alkane used, then deduct [1].
No penalty for not labelling steps, but deduct [1] if any wrongly
labelled.
5
[5]
29
49.
(i)
H
CH 3
C
CH3
/ CH 3 CH(OH)CH 3 ;
OH
1
Allow bond to HO rather than OH or halfway between the two
(ii)
secondary;
1
(iii)
CH3 COCH3 /propanone/acetone;
Allow ECF from a different alcohol drawn in (i)
1
[3]
50.
(a)
SN2 / bimolecular;
(b)
(i)
(ii)
1
reaction slower;
neutral/uncharged/less polar/electrons donated less easily in H 2 O;
2
reaction faster;
less bulky group/reduced steric hindrance;
2
[5]
51.
(i)
(Empirical formula =) C8 H8 O3 ;
H
H
O
O
H C
C
H
O
H
;
H
H
H
Allow double bonds on arene in alternate positions, or allow
delocalized representation (of pi electrons).
(ii)
(iii)
the bond at 0.1373 nm is a double bond and the bond at 0.1424 nm is a
single bond;
in CO2 (g) both bonds are double bonds and would have a value
around 0.137 nm;
Ester;
Arene/benzene ring;
Alcohol;
Award 2 for any three correct, award [1] for any two correct.
Do not accept alkane as a type of functional group in this
molecule.
2
2
2
[6]
52.
(i)
(ii)
boiling point increases as the number of carbons increases/OWTTE;
Greater Mr and hence greater van der Waals’/London/dispersion forces present;
2
⎯/ UV
⎯light
⎯
⎯→ CH3 Cl + HCl;
CH4 + Cl2 ⎯hv
Do not award mark if hv/uv light is not given.
30
Initiation step:
Cl2 ⎯hv
⎯/ UV
⎯light
⎯
⎯→ 2Cl•;
Do not award mark if hv/uv light is not given.
Penalize once only.
Propagation step:
CH4 + Cl• → CH3 • + HCl;
CH3 • + Cl2 → CH3 Cl + Cl•;
Termination step:
Cl• + Cl• → Cl2 or Cl• + CH3 • → CH3 Cl or CH3 • + CH3 • → CH3 CH3 ;
Allow fish-hook half-arrow representations i.e. use of
.
Penalize use of full curly arrows once only.
Penalize missing dots on radicals once only.
5
[7]
53.
(i)
(ii)
A. = CH3 (CH2 )7CHO;
B. = CH3 (CH2 )7 COOH/CH3(CH2 )7 CO2 H;
C. = (CH3 )3 COH;
D. = (CH3 )2 CO;
E. = BrCH2 CH2 Br;
Allow correct structural formulas.
5
addition;
/-(CH2 -CH2 )3-/-(CH2)6-;
2
[7]
54.
(a)
(b)
same general formula/Cn H2n ;
formulas of successive members differ by CH 2 ;
similar chemical properties/same functional group;
gradation/gradual change in physical properties;
Award [1] each for any three.
3
but-2-ene;
Accept 2-butene.
strongest intermolecular/van der Waals’ forces;
largest (molecular) mass/size/surface area/area of contact;
(c)
CH2 CHCH2 CH2 CH3/CH3CHCHCH2 CH3/any correct branched structure;
Accept more detailed formula.
pent-1-ene/pent-2-ene;
Name must match formula.
Accept 1-pentene/2-pentene.
(d)
3
2
C4 H8 + HBr → CH3 CH2 CHBrCH3;
Award [1] for all molecular formulas correct and [1] for
correct product structure.
Award [1] for completely correct equation starting with
but-1-ene.
31
(e)
(f)
addition;
3
oxidation/redox;
(potassium) dichromate(VI)/ Cr2 O 72− ;
(sulfuric) acid;
distilling off propanal as it is formed;
heating under reflux (to obtain propanoic acid);
5
(propan-1-ol) hydrogen bonding;
(propanal) dipole-dipole attractions;
(propanoic acid) hydrogen bonding;
propanoic acid > propan-1-ol > propanal;
4
[20]
55.
(a)
(b)
same general formula;
successive members differ by CH 2 ;
Do not allow elements or just “they”.
similar chemical properties;
Allow same/constant.
gradual change in physical properties;
Do not allow change periodically.
same functional group;
Award [1] each for any two.
2
add bromine (water);
alkanes − no change/stays or turns brown;
Allow red-brown or any combination of brown, orange or
yellow.
alkenes − bromine (water) decolorizes;
Do not allow clear or discoloured.
or
add (acidified) KMnO4 ;
alkanes − no change;
alkenes − KMnO4 decolorizes/brown/black;
3
[5]
56.
(a)
(CH3 )2 CHBr/more detailed formula;
secondary/2 because two alkyl groups attached to C with Br;
(b)
(c)
2
nucleophilic substitution;
bimolecular/molecularity of two/two species in rate-determining step;
Accept second order.
rate = k [(CH3 )2CHBr][OH−];
No penalty for incorrect halogenoalkane formula.
3
(CH3 )2 CH+/more detailed formula;
1
[6]
32
57.
(a)
(b)
one general formula/same general formula;
differ by CH2 ;
similar chemical properties;
gradual change in physical properties;
Award [1] for any two of the above characteristics.
1
ethanol lower/ethanoic acid higher;
due to larger mass of ethanoic acid/stronger van der Waals’/
London/dispersion forces;
due to stronger hydrogen bonding/2 hydrogen bonds per molecule;
Accept either answer
2
(c)
2
Allow condensed structural formulas such as
CH3 CH2 CH2CH2OH.
Award [2] for all three correct isomers, [1] for any two correct
isomers.
[5]
33
IB Chemistry – HL
Topic 10 Questions
1.
Which statement about the reactions of halogenoalkanes with aqueous sodium hydroxide is
correct?
A.
Primary halogenoalkanes react mainly by an SN1 mechanism.
B.
Chloroalkanes react faster than iodoalkanes.
C.
Tertiary halogenoalkanes react faster than primary halogenoalkanes.
D.
The rate of an SN1 reaction depends on the concentration of aqueous sodium hydroxide.
(Total 1 mark)
2.
Which compound can exist as optical isomers?
A.
H2 NCH2 COOH
B.
CH2 ClCH2 Cl
C.
CH3 CHBrI
D.
HCOOCH3
(Total 1 mark)
3.
Which compound reacts most rapidly by a SN1 mechanism?
A.
(CH3 )3 CCl
B.
CH3 CH2 CH2CH2Br
C.
(CH3 )3 CBr
D.
CH3 CH2 CH2CH2Cl
(Total 1 mark)
4.
Which compound can exist as optical isomers?
A.
CH3 CHBrCH3
B.
CH2 ClCH(OH)CH2 Cl
C.
CH3 CHBrCOOH
D.
CH3 CCl2 CH2 OH
(Total 1 mark)
1
5.
Which pair of compounds can be used to prepare CH 3 COOCH3 ?
A.
Ethanol and methanoic acid
B.
Methanol and ethanoic acid
C.
Ethanol and ethanoic acid
D.
Methanol and methanoic acid
(Total 1 mark)
6.
Which pair of compounds can be used to prepare CH 3 COOCH3 ?
A.
Ethanol and methanoic acid
B.
Methanol and ethanoic acid
C.
Ethanol and ethanoic acid
D.
Methanol and methanoic acid
(Total 1 mark)
7.
Nylon is a condensation polymer made up of hexanedioic acid and 1,6-diaminohexane.
Which type of linkage is present in nylon?
A.
Amide
B.
Ester
C.
Amine
D.
Carboxyl
(Total 1 mark)
8.
How many chiral carbon atoms are present in a molecule of glucose?
A.
1
B.
2
C.
3
D.
4
(Total 1 mark)
2
9.
Which amino acid can exist as optical isomers?
A.
B.
C.
D.
(Total 1 mark)
10.
What is the product of the following reaction?
Ni
CH3 CH2 CH2CN + H2 ⎯⎯→
A.
CH3 CH2 CH2NH2
B.
CH3 CH2 CH2CH3
C.
CH3 CH2 CH2CH2CH3
D.
CH3 CH2 CH2CH2NH2
(Total 1 mark)
3
11.
What is the correct IUPAC name for the following compound?
CH 3 CHCH 2 CH 2 CN
CH 3
A.
4-methylbutanenitrile
B.
4-methylpentanenitrile
C.
2-methylbutanenitrile
D.
2-methylpentanenitrile
(Total 1 mark)
12.
What is the organic product of the reaction between ethanol and ethanoic acid in the presence of
sulfuric acid?
A.
CH3 CHO
B.
CH3 COOCH3
C.
CH3 CH2 COOCH3
D.
CH3 COOCH2CH3
(Total 1 mark)
13.
Which compound can exist as optical isomers?
A.
H2 NCH2 COOH
B.
H3 CCONH2
C.
H3 CCHBrI
D.
HCOOCH3
(Total 1 mark)
14.
One of the alcohols containing four carbon atoms exists as optical isomers.
Give the structural formula and name of this alcohol and explain why it exists as optical
isomers. Outline how these two isomers could be distinguished.
..............................................................................................................................................
..............................................................................................................................................
..............................................................................................................................................
..............................................................................................................................................
..............................................................................................................................................
..............................................................................................................................................
..............................................................................................................................................
(Total 4 marks)
4
15.
Many organic compounds can exist as isomers. Draw and name an isomer of ethanoic acid,
CH3 COOH.
(Total 2 marks)
16.
Explain why butan-2-ol, CH3 CH(OH)CH2 CH3, exists as enantiomers, and describe how pure
samples of the enantiomers can be distinguished experimentally.
(Total 3 marks)
17.
Polyesters are formed in a condensation reaction. The structure of the repeat unit of a polyester
is
C
O
(i)
C H
C
O
CH
CH
O
O
Draw the structures of the two monomers that react to form this polyester.
(2)
(ii)
Identify the essential feature of the monomers in (i) that enable them to form a
condensation polymer.
(1)
(Total 3 marks)
18.
(i)
State the meaning of the term isomers.
(1)
(ii)
Draw the functional group isomers of C3 H6 O.
(2)
(iii)
State the meaning of the term optical isomers. Draw the alcohol with the molecular
formula C4 H10 O which exhibits optical isomerism and identify the chiral carbon atom.
(3)
(Total 6 marks)
19.
(i)
State the meaning of the term isomers.
(1)
(ii)
Draw the functional group isomers of C3 H6 O.
(2)
(iii)
State the meaning of the term optical isomers. Draw the alcohol with the molecular
formula C4 H10 O which exhibits optical isomerism and identify the chiral carbon atom.
(3)
(iv)
Other than the optical isomers in (iii), draw the other three alcohol isomers of molecular
formula C4 H10 O and identify the isomer that does not undergo oxidation.
(2)
(Total 8 marks)
5
20.
(i)
Draw the structural formula of the ester propyl ethanoate.
(1)
(ii)
Deduce the name and draw the structural formula of the alcohol and carboxylic acid that
react to form this ester.
Name of alcohol: .............................................................................................
Structural formula:
Name of carboxylic acid: .................................................................................
Structural formula:
(4)
(Total 5 marks)
21.
The compound, 2-bromobutane, CH3 CHBrCH2CH3, can react with sodium hydroxide to form
compounds F, G and H.
Compound F, C4 H10 O, exists as a pair of optical isomers. Compounds G and H, C4 H8 , are
structural isomers, and compound H exists as a pair of geometrical isomers.
(i)
Draw the structures of the two optical isomers of F.
(2)
(ii)
Outline the use of a polarimeter in distinguishing between the optical isomers.
(2)
6
(iii)
Draw diagrams to show the shapes of the two geometrical isomers of H.
(2)
(iv)
Draw the mechanism, using curly arrows to represent the movement of electron pairs, to
show the formation of G.
(3)
(Total 9 marks)
22.
A compound, J, has the molecular formula C2 H4 O2 and is obtained from a reaction between
methanoic acid and methanol. Write an equation for this reaction and state the name of
compound J.
(Total 3 marks)
23.
(a)
Compounds of formula C4 H7 Cl exhibits both geometrical and optical isomerism.
(i)
Explain why C4 H7 Cl shows geometrical isomerism.
........................................................................................................................
........................................................................................................................
(1)
(ii)
Draw the cis and trans isomers of C4 H7 Cl.
(2)
(iii)
Draw the structural formula of C4 H7 Cl that shows only optical isomerism. Show
the chiral carbon atom with “ * ”.
(2)
7
(b)
Explain why 1,2-dichlorocyclopropane has cis and trans isomers. Draw the structural
formulas of the two isomers.
...................................................................................................................................
...................................................................................................................................
(3)
(Total 8 marks)
24.
(a)
There are geometrical isomers of the cyclic compound C4 H6 Cl2 . Draw the structural
formula of two isomers and explain why these two isomers exist.
cis-isomer
trans-isomer
....................................................................................................................................
....................................................................................................................................
(3)
(b)
(i)
Draw the structural formulas of two isomers of but-2-ene-1,4-dioic acid.
(2)
(ii)
State and explain which isomer will have a lower melting point.
.........................................................................................................................
.........................................................................................................................
.........................................................................................................................
(2)
(iii)
Describe how the two isomers can be distinguished by a chemical test.
.........................................................................................................................
.........................................................................................................................
.........................................................................................................................
.........................................................................................................................
(2)
8
(c)
Consider the following compounds:
1-chloropentane, 2-chloropentane, 3-chloropentane
(i)
Identify the compound which exhibits optical isomerism and draw the structures of
the two isomers.
.........................................................................................................................
(3)
(ii)
Describe how these two isomers can be distinguished experimentally.
.........................................................................................................................
.........................................................................................................................
(1)
(Total 13 marks)
25.
Compounds with the molecular formula C3 H4 Cl2 exist as several structural isomers, some of
which are cyclic. Some of these structural isomers exist as geometric isomers.
(a)
Explain why geometrical isomerism is possible in the non-cyclic isomers.
....................................................................................................................................
....................................................................................................................................
(1)
(b)
Draw the structure of a non-cyclic structural isomer that does not exist as geometric
isomers, and explain why geometrical isomerism is not possible in this compound.
....................................................................................................................................
....................................................................................................................................
(2)
(c)
1,3-Dichloropropene exists as geometric isomers. Draw and label the structures of its cis
and trans isomers.
(2)
9
(d)
Draw structures to show the two geometric isomers of 1,2-dichlorocyclopropane.
(2)
(Total 7 marks)
26.
This question is about structural isomers and stereoisomers with the molecular formula
C4 H6 Cl2 .
(a)
The compound 1,3-dichlorobut-1-ene can be used to illustrate two types of
stereoisomerism. For each type of stereoisomerism, draw two structures to show the
relationship between the two.
two geometrical isomers
two optical isomers
(4)
(b)
Explain the term racemic mixture.
....................................................................................................................................
....................................................................................................................................
(1)
(c)
Outline how the two optical isomers of 1,3-dichlorobut-1-ene can be distinguished from
each other and from a racemic mixture.
....................................................................................................................................
....................................................................................................................................
....................................................................................................................................
....................................................................................................................................
(3)
10
(d)
1,3-dichlorocyclobutane exists as geometrical isomers. Draw the 3-dimensional structures
of these isomers.
(2)
(Total 10 marks)
27.
The molecular formula, C3 H4 Cl2 represents several isomeric compounds. Some isomers are
cyclic and some are unsaturated.
(a)
Draw the structures of two cyclic compounds that are structural isomers and state the
names of both isomers.
(2)
(b)
Two of the non-cyclic compounds have geometrical isomers. Draw the structures of these
compounds and their geometrical isomers.
(2)
(Total 4 marks)
11
IB Chemistry – HL
Topic 10 Answers
1.
C
2.
C
3.
C
4.
C
5.
B
6.
B
7.
A
8.
D
9.
B
10.
D
11.
B
12.
D
13.
C
14.
CH3 CH(OH)CH2 CH3 ;
Accept more detailed formula.
butan-2-ol;
Accept 2-butanol.
ECF for correct name of another C 4 alcohol.
contains a chiral/asymmetric carbon atom/four different groups around one
carbon atom;
(plane of) plane-polarized light rotated in opposite directions;
4
[4]
15.
methyl methanoate;
HCOOCH3 ;
Accept other correct alternative.
[2]
16.
(the molecule contains a) chiral/asymmetric carbon atom/carbon atom with four
different groups;
polarized light passed through;
(plane of polarization) rotated in opposite/different directions;
[3]
17.
(i)
HOCH2 CH2 OH;
HOOCC6H4COOH;
(ii)
reactants have two functional groups/OWTTE;
2
1
[3]
12
18.
(i)
same molecular formula but different structural formula/
arrangement of atom within a molecule/OWTTE;
1
(ii)
2
H
CH 3 COCH 3
H
O
C
C
H
C
H
H
CH 3CH 2CHO
H
;
H
H
C
C
H
H
C
;
O
H
H
Accept unsaturated alcohol and cyclic alcohol as alternative
answers. If more than two correct isomers given − no penalty −
but a third incorrect structure cancels a correct one. i.e. two
correct, one incorrect equals [1].
(iii)
isomers that can rotate plane polarized light in opposite directions;
Do not accept bend, reflect plane-polarized light.
H
H
H
H
H
C
H
C
C
C* H
H
H
O
H
H
H
H
H
C
C
C
H
H
OH H
* C
H
H
H
C 2H 5
C*
CH 3
other correct structure;
OH
Penalize missing bonds/hydrogens.
correct identification of chiral carbon (*);
3
[6]
19.
(i)
same molecular formula but different structural formulae/arrangement
of atoms within a molecule/OWTTE;
1
13
(ii)
H
CH 3 COCH 3
H
O
C
C
H;
C
H
H
CH 3 CH 2CHO
H
H
H
C
C
H
H
H
H
C
;
O
2
Accept unsaturated alcohol and cyclic alcohol as alternative
answers.
If more than two correct isomers given − no penalty − but a
third incorrect structure cancels a correct one. i.e. two correct,
one incorrect equals [1].
(iii)
isomers that can rotate plane polarized light in opposite directions;
Accept two molecules/compounds, which are mirror images of
each other.
Do not accept bend, reflect plane-polarized light.
H
H
H
H
H
C
H
C
C
C* H
H
H
O
H
H
H
H
H
C
C
C* C
H
H
OH H
H
H
H
C2H5
C*
CH 3
other correct structure;
OH
correct identification of chiral carbon (*);
3
14
(iv)
H
H
H
H
H
C
C
C
C
H
H
H
H
H
OH
H
CH 3 H
C
C
C
H
H
H
OH
H
C
H
H
H
H
H
C
C
C
H
OH H
H ;
Award [1] mark for 3 structures
H
H
C
H
2-methylpropan -2-ol /
H
H
H
C
C
C
H
OH H
H ;
2
Penalize [1] mark for the omission of H in (i) to (iv).
[8]
20.
(i)
O
CH 3
C
O
CH 2CH 2 CH 3 / CH 3COOCH 2CH 2CH 3 ;
1
(ii)
propan-1-ol/1-propanol;
CH3 CH2 CH2OH;
Accept full structural formula showing all bonds and atoms but
no mark if H atoms missing
ethanoic acid/acetic acid;
CH3 COOH;
Accept full structural formula penalize missing Hs only once
ECF from incorrect ester in (i) above
4
[5]
15
21.
(i)
OH
OH
C
CH3 CH2
H
H
CH3
CH3
C
CH 2 CH 3
2
Award [2] for both tetrahedral structures, or [1] if tetrahedral
structure is not clear.
(ii)
plane polarized light;
rotation in opposite/different directions;
2
(iii)
CH 3
CH 3
C
;
H
H
CH 3
H
C
H
(iv)
C
C
;
CH 3
curly arrow showing attack by – OH on end H;
curly arrow showing C–Br bond fission;
curly arrow showing formation of double bond;
H2 O and Br– shown as products;
max
Award [1] each for any three.
If but-2-ene formed, award [2 max].
2
3
[9]
22.
CH3 OH + HCOOH → HCOOCH3 + H2 O
Award [1] for both reactants and [1] for both products (accept
C2 H4 O2 ).
methyl methanoate;
3
[3]
16
23.
(a)
(i)
no rotation possible due to double bond/ bond;
Accept restricted or hindered rotation.
1
(ii)
2
H
H
C
Cis
;
C
CH 2Cl
CH 3
H
CH 3
C
trans
C
;
CH 2Cl
H
OR
H
H
C
;
C
CH2CH 3
Cl
Cis
H
Cl
CH 2CH 3
H
trans
;
C
C
OR
CH 3
CH3
C
Cis
C
Cl
H
CH 3
H
C
trans
CH 3
C
Cl
17
(iii)
H
H
C
H
C
;
*C
H
CH 3
Cl
2
Award [1] for the structure and [1] for showing * on the correct
carbon atom.
(b)
restricted rotation because C⎯C bond is now part of a cyclic system;
3
CH 2
H
Cl
C
;
C
H
Cl
trans
CH 2
H
H
C
;
C
Cl
Cl
Cis
[8]
24.
(a)
restricted rotation because C−C bond is now part of a cyclic system;
H
H
Cl
H
H
H
Cl
H
H
;
H
Cl
H
3
H
Cl
H
;
H
Award [1] for each correct 3D structure.
If correct structure, but not 3D, or wrongly labelled award [1]
only.
Accept 1,3-disubstituted cyclo compound, or any other correct
isomer.
18
(b)
(i)
2
COOH
HOOC
C
C
(ii)
(iii)
(c)
(i)
C
;
C
;
H
HOOC
H
H
COOH
H
cis isomer (has lower melting point than the trans-isomer);
intramolecular hydrogen bonds/weaker intermolecular forces
/less close packing;
2
(gentle) heating of a sample of each isomer;
cis isomer readily releases water vapour (forming a cyclic anhydride);
2
2-chloropentane;
C3H7
C3H7
C
C
;
C1
H
H
C1
CH3
;
CH3
3
Award [1] for each correct 3D structure.
If correct structures, but not 3D, award [1] only.
(ii)
rotation of the plane polarized light in opposite directions;
1
[13]
25.
(a)
restricted/no rotation around double bond/pi bond;
(b)
Accept either of these structures
Cl
H
C
Cl
H
C
Cl
C
CH 3
H
1
C
;
CH 2Cl
two identical atoms on one side of double bond/interchanging CH 3
and H/Cl and CH2 Cl makes no difference;
2
19
(c)
H
H
C
C
Cl
CH 2Cl
H
CH 2 Cl
C
Cl
cis;
C
trans;
H
Award [0] if structure of another isomer drawn.
Award [1] if both structures correct but unlabelled or wrongly
labelled.
Award [1] for each correctly drawn and labelled structure.
2
(d)
Cl
C
H
Cl
H
C
C
;
H
H
H
Cl
C
H
H
C
C
;
Cl
H
2
Award [0] if structure of another isomer drawn.
[7]
20
26.
(a)
H
Cl
C
CHClCH3
Cl
C
C
H
C
H
H
CHClCH 3
one correct structural formula;
two distinct isomeric structures shown;
ClHCHC
(b)
(c)
CH 3
CH 3
C
C
H
CHCHCl
H
Cl
Cl
one correct structural formula shown as 3-D;
two distinct isomeric structures shown;
4
it contains equal amounts/moles of the two optical isomers/enantiomers/
d (dextro/dextrorotatory) and l(levo/levorotatory);
1
polarimeter/plane-polarized light;
plane (of polarization) rotated in opposite directions;
no rotation for racemic mixture;
3
(d)
H
H
Cl
H
H
H
H
Cl
H
H
Cl
one correct structural formula;
two distinct isomeric structures shown;
H
Cl
H
H
H
2
[10]
21
27.
(a)
Cl
Cl
and 1,1 dichlorocyclopropane;
Cl
Cl
Cl
Cl
Cl
Cl
Cl
Cl
(cis- or trans-) 1,2 dichlorocyclopropane;
Award point for the correct name corresponding to the related
isomer.
Accept diagrams that do not display 3 dimensional structure.
Award [1 max] for correct structures only, without the
corresponding names.
2
(b)
Cl
Cl
CH 3
Cl
and
H
Cl
Cl
H
CH 2 Cl
Cl
CH 3
H
and
H
H
H
CH 2 Cl
2
[4]
22