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Chemistry M-TERM 1
Chapter 1 Chemical Equilibrium
WEEK 1
Basic Questions- Section 1.1 What is Equilibrium?
1. Which of the following systems constitute steady state situations, and which are at
equilibrium? For each, a constant property is indicated.
a) An unopened bottle of sparkling water (a bottle containing CO2 and water).
b) A Bunsen lit with a constant blue flame color.
Grid Question- Section 1.1 What is Equilibrium?
2.
a. What, specifically, is “equal” in a chemical reaction that has attained a state of equilibrium?
b. Why are chemical equilibria referred to as “dynamic”?
WEEK 2
Grid Questions -Section 1.6 Predicting new equilibrium concentrations
3.
 Effect of altering concentration of reactant or product on an equilibrium system
a. Predict what would happen to equilibrium concentrations of H2O(g) and H2(g) respectively if
the partial pressure of O2 is increased by injecting more O2(g) into the reaction chamber at 273°C.
H2O(g) ⇄ H2(g) + ½ O2(g)
H = + 242 kJ
b. 2SO3(g) ⇆ 2SO2(g) + O2(g)
After the equilibrium represented above is established, some pure SO2 (g) is injected into the
reaction vessel at constant temperature. Which of the following has a lower value compared to its
value at the original equilibrium?
a) The total pressure in the reaction vessel
b) The amount of SO3 (g) in the reaction vessel
c) The amount of O2 (g) in the reaction vessel
d) The amount of SO2(g) in the reaction vessel
c. Fe+3(aq) + SCN–(aq) ⇆ FeSCN+2(aq)
The reactants and products in the above reaction are at equilibrium. What would happen to the
[Fe+3(aq)] and [FeSCN+2(aq)] respectively if we add to the equilibrium mixture a few crystals of
KSCN?
d. Consider the following equilibrium system: Fe+3(aq) + SCN–(aq) ⇆ FeSCN+2(aq)
What would happen to the concentration of SCN–(aq) and FeSCN2+(aq), respectively, if Fe3+(aq)
is removed by adding phosphate ions?
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e. The chromate and dichromate ions are set up in equilibrium as follows:
2 CrO42- (aq) + 2 H+ (aq) ⇌ Cr2O7 2-(aq) + H2O (l)
Yellow
orange
What happens in each of the following cases? Complete the table below
Change
Increase [H+]
Increase in [CrO42-]
Decrease [H+]
Add NaOH
Add HCl
Equilibrium shifts to
Color change
Change to [Cr2O72-]
 Effect of adding a solid on a system at equilibrium
Predict what will happen to the equilibrium concentration of CO2 (g) if some CaCO3 (s) is added
to the system: CaCO3 (s) ⇆ CaO (s) + CO2 (g)

Effect of a change in volume/pressure on a gaseous system
a. Does Le Chatelier’s Principle predict a change of equilibrium concentrations for the following
reactions if the gas mixture is compressed? Give full explanation, showing the imposed change
and how the system reacts?
i. H2(g) + Cl2(g) ⇄ 2HCl(g)
ii. 2H2(g) + O2 (g) ⇄ 2H2O(g)
iii. 2NH3(g) ⇄ 3H2(g) + N2(g)
b. Predict what would happen to equilibrium concentrations of H2O(g) and H2(g) respectively if
the partial pressure of all gases is increased by decreasing the volume of the reaction chamber at
273°C.
H2O(g) ⇄ H2(g) + ½ O2(g)
H = + 242 kJ
c. Predict what would happen to equilibrium concentrations of CO(g) and NO(g) respectively if
the partial pressure of all gases is increased by decreasing the volume.
CO(g) + NO2(g) ⇆ CO2(g) + NO(g)
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d. Which of the systems in equilibrium represented below will exhibit a shift to the left when the
pressure on the system is increased by reducing the volume of the system? (Assume that
temperature is constant)
a) 2Mg(s) + O2(g) ⇄2MgO(s)
b) SF4(g) + F2(g) ⇄ SF6(g)
c) H2(g) + Br2(g) ⇄ 2HBr(g)
d) N2(g) + 3H2(g) ⇄2NH3(g)
e) SO2Cl2(g) ⇄SO2(g) + Cl2(g)
 Effect of adding helium gas into the reacting chamber
Predict what would happen to equilibrium concentrations of H2O(g) and H2(g) respectively if the
pressure in the reaction chamber at 273°C is increased by adding (unreactive) He(g).
H2O(g) ⇄ H2(g) + ½ O2(g)
H = + 242 kJ

Effect of a change of temperature on an equilibrium system
Given: 4HCl(g) + O2(g) ⇄ 2H2O (g) + 2Cl2 (g)
∆Ho = −113 kJ.mol-1
The equilibrium system represented above is contained in a sealed, rigid vessel. Which of the
following will decrease if the temperature of the mixture is raised?





[HCl (g)]
[O2 (g)]
[Cl2 (g)]
[HCl (g)] and [Cl2 (g)]
[O2 (g)], [HCl (g)] and [Cl2 (g)]
 Effect of catalyst on system at equilibrium
Consider the following system:
2 SO2 (g) + O2 (g) ⇆ 2 SO3 (g)
State the effect of adding a catalyst to the system on the position of equilibrium?
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4. The “yield” can be defined as the percentage of reactants that form products.
Ammonia is produced by combining nitrogen and hydrogen gases as follows:
N2(g) + 3H2(g) ⇄ 2NH3(g)
H = –92 kJ
How can we increase the yield of ammonia production?
Basic Questions- Section 1.7 Quantitative aspects of equilibrium
5. What is/ are the factor(s) that affect the value of the equilibrium constant?
6. Consider the following reaction whose Keq= 45.9 at 490 C: H2(g) + I2(g) ⇆ 2HI(g)
1.0 mole of each of H2 (g) and I2 (g) are injected into a 2.0 dm3 container at 490ºC. Find the
concentration of each species once equilibrium is reached.
7.
a) Given: AgCl(s) ⇄Ag+(aq) + Cl–(aq)
Calculate the [Ag+] at equilibrium,.
2+
–

 Pb (aq) + 2Cl (aq)
b) PbCl2(s) 

Calculate the [Pb2+] at equilibrium,.
Ksp = 1.7  10–10 at 25ºC
Ksp = 1.6  10–5 at 25ºC
Grid Question- Section 1.7 Quantitative aspects of equilibrium
8. Consider the following reaction: H2(g) + I2(g) ⇆ 2HI(g)
At 490ºC, its equilibrium constant value is, Keq = 45.9.
At a certain instant in a reaction the following concentrations were determined, while container is
maintained at 490ºC:
[HI] = 0.01 mole/dm3
[H2] = 0.005 mole/dm3
[I2] = 0.004 mole/dm3
Is the system at equilibrium?
If not, which concentrations are increasing and which are decreasing?
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Top Question
1. The graphs below show how the yield of product changes as pressure and temperature varies.
Which of the following reactions is represented in the graphs?
A. N2(g) + 3H2(g) ⇄ 2NH3(g)
B. 2SO3(g) ⇄ 2SO2(g) + O2(g)
C. H2CO3(aq) ⇄ H2O(l) + CO2(g)
D. N2(g) + O2(g) ⇄ 2NO(g)
E. 2H2O(l) ⇄ O2(g) + 2H2(g)
H < 0
H > 0
H > 0
H < 0
H > 0
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Chapter 2 Solubility: A case of equilibrium
WEEK 3
Basic Questions- Section 2.1 Solubility: A case of equilibrium
9. Propose two methods for increasing the rate at which salt dissolves in water.
10. Given:
a) What is the solubility of potassium nitrate at 20ºC? at 50ºC?
b) By how much will the solubility of potassium nitrate increase if the temperature is increased
from 20°C to 50°C?
c) What mass of potassium nitrate forms if a saturated solution in 100g of water is cooled from
40ºC to 10ºC? From 60ºC to 20ºC?
d) What mass of water is needed to dissolve 50g of potassium nitrate at 40ºC?
e) What mass of potassium sulfate will dissolve in 2 kg of water at 20ºC?
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Basic Questions- Section 2.2 Aqueous solutions
11. a) Glucose is (molecular / ionic) and it dissolves in water as such: C6H12O6(s) →
b) NaCl is (molecular / ionic) and it dissolves in water as such: NaCl(s) →
c) HCl gas is (molecular / ionic) and it dissolves in water as such: HCl(g) →
12. If we mix the following solutions:
 state what will be observed
 name of precipitate
 write a net ionic equation of the reaction occurring
a) Ba(NO3)2 and Na2SO4
b) AgNO3 and NaCl
c) Ba(NO3)2 and NaCl
d) Ca(NO3)2 and Na2CO3
13. A solution is known to contain Pb(NO3)2 and Mg(NO3)2. It is required to remove the lead(II)
ions from the solution without removing magnesium ions. This can be done by adding just
the right amount of
a) NaCl.
b) HNO3.
c) NH3.
14. What would be the weight of the precipitate formed when 100 mL of 0.50 M NaCl is added to
50.0 mL of 0.10 M AgNO3?
Grid Questions- Section 2.2 Aqueous solutions
15. Classify the following salts as soluble or slightly soluble:
a) all salts containing nitrates NO3-.
b) all salts of Na+ and K+ ions.
c) chlorides of silver and lead.
16. List the observation upon adding a few milliliters of silver nitrate solution to the following
potassium halide solutions. Identify the precipitate if any.
a) KF
b) KCl
c) KBr
d) KI
17. Which mixture can be separated by adding water, stirring and filtering?
a)
b)
c)
d)
e)
BaCl2, and NaCl
CaCO3 and NaCl
Cu and Mg
C2H5OH and CH3COOH
Salt and water
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18. Fill the table below with the colors that the following cations impart on a blue flame.
Species tested
Result & Equation
Li+
Lithium ion
K+
Potassium ion
Na+
Sodium ion
Cu2+
Copper (II) ion
Ba2+
Barium ions
Ca2+
Calcium ions
19. Match the following ions to the below observations: Ca2+, Zn2+, Cu2+, Fe2+, Al3+, Fe3+
Observation with solution tested
Reddish-brown precipitate with sodium
hydroxide, does not dissolve in excess acid
Dirty-green precipitate with sodium
hydroxide, does not dissolve in excess base
Blue precipitate with sodium hydroxide,
does not dissolve in excess base
White precipitate with sodium hydroxide,
does not dissolve in ammonia
White gelatinous precipitate with ammonia,
does not re-dissolve in excess ammonia
White precipitate with sodium hydroxide,
re-dissolve in excess base
Cation present in solution
20. Fill in the table below by the following ions:
Sulfite ion SO32-, sulfate ion SO42-, carbonate ion CO32-, ammonium ion NH4+, aluminum ion
Al3+, nitrate ion NO3- in the presence of aluminum powder and heat.
Release a gas upon reacting
Release a gas upon reacting Do not release gases with
with a base
bases, nor with acids
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WEEK 4 (Chemical Analysis)
Basic Question- Section 2.3 The equilibrium law
21. Which of the following salts is the most soluble?
Salt
TlCl
SrCrO4
CaSO4
a)
b)
c)
d)
e)
Ksp
1.9 × 10-4
3.6 × 10-5
2.4 × 10-4
Salt
CuCl
BaCrO4
BaSO4
Ksp
3.2 × 10-7
8.5 × 10-11
1.5 × 10-9
Salt
AgCl
PbCrO4
PbSO4
Ksp
1.7 × 10-10
2 × 10-16
1.3 × 10-8
PbSO4
BaSO4
AgCl
CaSO4
PbCrO4
22. a) 0.0010 mol HCl and 1.0 × 10–6 mol of Pb(NO3)2 are mixed forming 2.0 L solution. Will a
precipitate form? For PbCl2 Ksp = 1.3 × 10–8.
b) Equal volumes of 0.020 M CaCl2 and 0.00040 M Na2SO4 are mixed. Will a precipitate form?
Grid Question- Section 2.3 The equilibrium law
23.
a. It is required to calculate the solubility of CuCl(s) in water in terms of the Ksp
b. It is required to calculate the solubility of PbCl2 (s) in water in terms of the Ksp
c. It is required to calculate the solubility of FeCl3 (s) in water in terms of the Ksp
d. Calculate the Ksp of PbCl2, knowing that it has a solubility of 1.1 × 10-2 M at T°C in water.
e. Calculate the solubility of CuCl in water, knowing that its Ksp is equal to 3.2 × 10-7 at T°C.
24. a) To 200 cm3 of 0.10 M HCl is added 300 cm3 of 0.20 M Pb(NO3)2. Will a precipitate form?
b) 300. cm3 of 0.200 M Mg(NO3)2 are added to 200. cm3 of 0.200 M NaOH. Will a precipitate
form? Ksp of Mg(OH)2 = 1.8 × 10-11
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SUMMARY of solubility rules to remember:
 Sodium, potassium, ammonium salts are soluble
 Nitrates and acetates are soluble
 Halides are soluble except lead and silver
 Sulfates are soluble except barium, calcium and lead
 Carbonates and hydroxides are insoluble, except sodium, potassium and ammonium
Colors of common compounds you need to remember:
 All group 1 salts are white and their solutions are colourless
 All carbonate salts are white except copper(II) carbonate which is green
 All sulfide salts are white except CuS, FeS and PbS which are black
 Iron (II) hydroxide, Fe(OH)2, is green
 Iron (III) hydroxide, Fe(OH)3, is red- brown
 Chromium (III) hydroxide, Cr(OH)3, is green
 Barium sulfate is white
 Silver chloride is white , silver bromide is creamy white and silver iodide is yellow
 Copper (II) hydroxide is blue
 lead (II) chloride is white, lead (II) bromide is creamy, lead(II) iodide is yellow
 Aluminium hydroxide is gelatinous white
 Zinc hydroxide is white
 Calcium hydroxide is white
Species tested
Test
Result & Equation
Test for Cations
+
Li
Lithium ion
K+
Potassium ion
Na+
Sodium ion
Cu2+
Copper (II)
ion
Ba2+
Barium ions
Ca2+
Calcium ions
Flame Test
red flame
Flame Test
Lilac flame
Flame test
yellow flame
Flame test
Blue-green flame
Flame test
Pale green
Flame test
Brick red
Flame test is used to show the presence of a certain metal ion in a compound:
 A platinum or nichrome wire is cleaned by dipping in concentrated
hydrochloric acid and then holding it in a hot Bunsen flame.
 This is repeated until the wire does not give any colour in the flame.
 The wire is dipped back into the concentrated acid, then into a tiny sample of
the solid you are testing and back into the flame.
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Gas
H2(g)
Hydrogen
O2(g)
Oxygen
CO2(g)
Carbon dioxide
Cl2(g)
Chlorine
SO2(g)
Sulphur dioxide
NH3(g)
Ammonia
Colour
Colourless
Colourless
Colourless
Pale green
Colourless
Colourless
Odour
Method of collection
Odourless
Less dense than air, collected
by:
 Upward delivery
 Over water
 Using a gas syringe
Odourless
Density close to that of air,
collected:
 Over water
 Using a gas syringe
Odourless
Denser than air, collected by:
 Downward delivery
 Using a gas syringe
 Over water
Denser than air, soluble,
collected by:
Has a pungent smell
 Downward delivery
 Using a gas syringe
Smells like burnt
matches
Denser than air, collected by:
 Downward delivery
 Using a gas syringe
Less dense than air, by:
Has a pungent smell
 Upward delivery
 Using a gas syringe
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Test for Gases
H2
Hydrogen gas
Lit splint
Explodes (burns with a squeaky popping sound)
O2
Oxygen gas
Glowing splint
Glowing splint relights
NH3
Damp red litmus paper
Ammonia gas
Turns blue.
Pass the gas through lime water for a short
time, lime water turns cloudy white (milky).
CO2
Carbon
dioxide gas
Lime water
Ca2+(aq) + 2OH-(aq) + CO2(g) → CaCO3(s)
+ H2O(l)
If the gas is passed through lime water for a long
time the precipitate dissolves and the solution
becomes clear again.
Cl2
Chlorine gas
SO2
Sulfur dioxide
Damp litmus paper
acidified potassium
manganate (VII)
blue litmus paper turns red then gets bleached
red litmus paper gets bleached
Decolorizes the purple solution
Turns blue cobalt chloride paper pink
anhydrous cobalt
chloride paper
H2O
Water vapour
OR liquid
CuSO4(s) + 5H2O(l/g) → CuSO4.5H2O(s)
white
blue
Turns white anhydrous copper(II) sulfate blue
anhydrous copper(II)
sulfate
CoCl2(s) + 6H2O(l) → CoCl2.6H2O(s)
blue
pink
Cobalt chloride paper is a filter paper that has been dipped into cobalt (II) chloride
solution then dried thoroughly in a desiccator.
A desiccator is a piece of glassware or a small cabinet which contains a tray of some
substance that absorbs water.
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Test for Cations
Species tested
NH4+
Ammonium
ion
2+
Ca
Calcium ion
Reagent used
NaOH(aq)
add to an aqueous
solution of the
specimen to be tested
then warm gently
NaOH(aq)
add to an aqueous
solution of the
specimen tested
NH3(aq)
add to an aqueous
solution of the
specimen tested
NaOH(aq)
add to an aqueous
solution of the
specimen tested
Al3+
Aluminium ion
NH3(aq)
add to an aqueous
solution of the
specimen tested
NaOH(aq)
Zn2+
Zinc ion
add to an aqueous
solution of the
specimen tested
NH3(aq)
add to an aqueous
solution of the
specimen tested
Result & Equation
No precipitate but a pungent smelling gas is
produced that turns damp red litmus paper blue.
∆
NH4+(aq) + OH-(aq) → NH3(g) + H2O(l)
A white precipitate is formed that does not
dissolve in excess NaOH(aq)
Ca2+(aq) + 2OH-(aq) → Ca(OH)2(s)
No reaction / no precipitate
A white precipitate is formed that dissolves in
excess NaOH(aq)
Al3+(aq) + 3OH-(aq) → Al(OH)3(s)
A white precipitate of Al(OH)3 is formed that
does not dissolve in excess NH3(aq).
Al3+(aq) + 3OH-(aq) → Al(OH)3(s)
In both cases:
A white precipitate of Zn(OH)2 is formed that
dissolves in excess reagent
Zn2+(aq) + 2OH-(aq) → Zn(OH)2(s)
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Species tested
Test
Result & Equation
NaOH(aq)
Fe2+
Iron (II) ion
add to an aqueous
solution of the specimen
tested
NH3(aq)
add to an aqueous
solution of the specimen
tested
NaOH(aq)
Fe3+
Iron (III) ion
add to an aqueous
solution of the specimen
tested
NH3(aq)
add to an aqueous
solution of the specimen
tested
NaOH(aq)
add to an aqueous
solution of the specimen
2+
tested
Cu
Copper (II) ion
NH3(aq)
add to an aqueous
solution of the specimen
tested
NaOH(aq)
add to an aqueous
solution of the specimen
Cr3+
tested
Chromium(III)
NH3(aq)
ion
add to an aqueous
solution of the specimen
tested
In both cases:
Green precipitate is formed that does not dissolve
in excess reagent
Fe2+(aq) + 2OH-(aq) → Fe(OH)2(s)
In both cases:
Red-brown precipitate is formed that does not
dissolve in excess reagent
Fe3+(aq) + 3OH-(aq) → Fe(OH)3(s)
Blue precipitate of is formed that does not dissolve
in excess NaOH(aq)
Cu2+(aq) + 2OH-(aq) → Cu(OH)2(s)
Blue precipitate of Cu(OH)2 is formed that
dissolves in excess NH3(aq) to form a deep blue
solution.
Cu2+(aq) + 2OH-(aq) → Cu(OH)2(s)
Grey green precipitate is formed that dissolves in
excess NaOH(aq)
Cr3+(aq) + 3OH-(aq) → Cr(OH)3(s)
Grey-green precipitate that does not dissolve in
excess NH3(aq)
Cr3+(aq) + 3OH-(aq) → Cr(OH)3(s)
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Species tested
Test
Result & Equation
Test for Anions
Effervescence/bubbles or fizzing is observed.
Dilute acid
(HCl(aq) or HNO3(aq))
A colorless, odorless gas is produced that turns
lime water milky when passed through it for a short
time.
CO32MCO3(s) + 2H+(aq) → M2+(aq) + H2O(l) +
Add dilute acid to a sample
Carbonate ion
CO2(g)
of the specimen to be
OR
tested
CO32-(aq) + 2H+(aq) → H2O(l) + CO2(g)
(solid or aqueous)
If the gas is passed through lime water for a long
time the milkiness dissolves and the solution
becomes clear again.
Ba(NO3)2(aq)
White precipitate is formed that does not dissolve
in acid.
SO42Sulphate ion
SO32Sulphite ion
add to an aqueous
solution of the specimen
to be tested followed by
Ba2+(aq) + SO42-(aq) → BaSO4(s)
a dilute acid (HCl(aq) or
HNO3(aq))
The solution is acidified to make sure there is no
OR
carbonate ion CO32- or sulphite ion SO32-.
Acidified Ba(NO3)2(aq)
If a carbonate ion CO32- or sulphite ion SO32- is
add to an aqueous
present, they will not precipitate in acidified
solution of the specimen reagent.
to be tested
A colorless, chocking gas is formed that will
decolorize purple acidified aqueous potassium
manganate(VII)
Add dilute hydrochloric
or nitric acid to a solid
MSO3(s) + 2H+(aq) → M2+(aq) + H2O(l) +
or aqueous sample then
SO2(g)
warm gently
OR
SO32-(aq) + 2H+(aq) → H2O(l) + SO2(g)
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Species tested
Test
Result & Equation
Aluminium + NaOH(aq)
NO3Nitrate ion
add to an aqueous
solution of the specimen
to be tested then warm
gently
A pungent smelling gas is produced that turns
damp red litmus paper blue.
(Ammonia gas)
AgNO3(aq)
ClChloride ion
add to an aqueous
solution of the specimen
to be tested followed by
dilute acid
(HCl(aq) or HNO3(aq))
OR
Acidified AgNO3(aq)
A white precipitate is formed, which upon exposure
to light turns violet then black.
Ag+(aq) + Cl-(aq) → AgCl(s)
add to an aqueous
solution of the specimen
to be tested
AgNO3(aq)
-
Br
Bromide ion
add to an aqueous
solution of the specimen
to be tested followed by
dilute acid
(HCl(aq) or HNO3(aq))
OR
Acidified AgNO3(aq)
add to an aqueous
solution of the specimen
to be tested
A pale cream precipitate is formed, which upon
exposure to light turns violet then black.
.
+
Ag (aq) + Br (aq) → AgBr(s)
17
Species tested
Test
Result & Equation
AgNO3(aq)
-
I
Iodide ion
add to an aqueous
solution of the specimen
to be tested followed by
dilute acid
(HCl(aq) or HNO3(aq))
OR
Acidified AgNO3(aq)
A yellow precipitate is formed, which upon
exposure to light turns violet then black.
Ag+(aq) + I-(aq) → AgI(s)
add to an aqueous
solution of the specimen
to be tested
The acid is added to exclude other substances that might also produce precipitates with
silver nitrate solution.
18
Top Question
1. Many insoluble salts can be prepared by precipitation reaction between soluble salts.
All potassium salts and all nitrate salts are soluble.
a) A student adds a potassium sulfate, K2SO4, solution to a barium nitrate, Ba(NO3)2 solution.
A white precipitate is formed.
i. What is the chemical formula of the white precipitate?
ii. Write the chemical equation for the reaction between potassium sulfate and barium
nitrate.
b) Another student adds potassium carbonate, K2CO3, solution to a silver nitrate, AgNO3,
solution. A white precipitate is also formed.
i. What is the chemical formula of the white precipitate? ii. Write the net ionic equation of
the reaction between potassium carbonate and silver nitrate.
c) Barium nitrate solution produces a white precipitate with both carbonate ions and sulfate
ions. Why acidified barium nitrate solution is used to distinguish between carbonate ions
and sulfate ions?
2.
a. Two white solids, A and B, were analyzed. Complete the table below for each solid.
i. Add a few drops of NaOH to each solid and heat:
Observation
Inference
Solid A: A colorless pungent gas with the
characteristic smell of ammonia gas evolves
that turns wet red litmus blue.
ii. Add a few drops of nitric acid to each solid:
Observation
Solid B: Fizzing/ effervescence. The gas
evolved is colourless and turns lime water
milky.
Inference
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b. A white solid, X, was analyzed. Fill in the inferences and identify X.
Test
A few drops of Ba(NO3)2 are
added to a colorless solution of X
followed by HCl(aq)
Add a few drops of NaOH(aq) to a
solution of X followed by excess
base
Add a few drops of NH3(aq) to a
solution of X followed by excess
base
Observation
Inference
A white precipitate forms
insoluble in acid
A white precipitate forms
soluble in excess base
A white precipitate forms
soluble in excess base
Solid X is:
c. A white solid, Y, was analyzed. Fill in the inferences and identify Y.
Test
A few drops of AgNO3 are added
to a colorless solution of Y
followed by HNO3 (aq)
Add a few drops of NaOH(aq) to
a solution of Y followed by
excess base
Add a few drops of H2SO4(aq) to
a solution of Y
Observation
Inference
A yellow precipitate
forms insoluble in acid
A white precipitate forms
insoluble in excess base
A white precipitate forms
Solid Y is:
d. A blue solid, B, was analyzed. Fill in the inferences and identify B.
Test
Add a few pieces of Al
foil and 2ml of NaOH(aq)
to a blue solution of B
Warm gently.
Observation
Al foil darkens and slowly
dissolves. Bubbles of a gas
evolve and this gas turns
wet red litmus paper into
blue
Add a few drops of
A blue precipitate forms
NaOH(aq) to a solution of
insoluble in excess base
Y followed by excess base
A blue precipitate forms
Add a few drops of
that dissolves in excess
NH3(aq) to a solution of Y
base to form a deep blue
followed by excess base
solution.
Solid B is:
Inference
20
e. A light grayish-green solution, C, was analyzed. Fill in the inferences and identify C.
Test
A few drops of
AgNO3 are added to
2ml of the solution
of followed by
HNO3 (aq)
Add a few drops of
NaOH(aq) to 2ml of
the solution
followed by excess
base
Add a few drops of
NH3(aq) to 2ml of
the solution
Observation
Inference
A white precipitate forms
insoluble in acid
A green precipitate forms
insoluble in excess base
A green precipitate forms
insoluble in excess base
Solid C is:
f. A yellowish solution, D, was analyzed. Fill in the inferences and identify D.
Test
A few drops of
AgNO3 are added to
2ml of the solution
of followed by
HNO3 (aq)
Add a few drops of
NaOH(aq) to 2ml of
the solution
followed by excess
base
Add a few drops of
NH3(aq) to 2ml of
the solution
Solid D is:
Observation
A creamy precipitate
forms insoluble in acid
A red- brown precipitate
forms insoluble in excess
base
A red- brown precipitate
forms insoluble in excess
base
Inference
21
Chapter 3 Acids, Bases and Salts
WEEK 5
Grid Question- Section 3.1 Electrolytes – strong or weak
25. Calculate the [H+] and the [OH–] in a 0.010 M solution of LiOH at 25℃.
Basic Questions- Section 3.2 Experimental introduction to acids and bases
26. Which of the following elements below is most likely to form an acidic oxide?
A. A only
B. B only
C. C only
D. D only
E. A, B, and C
27. Consider the following.
Aluminum oxide
Calcium oxide
Carbon dioxide
Carbon monoxide
Magnesium oxide
Sulfur dioxide
Al2O3
CaO
CO2
CO
MgO
SO2
Which of the oxides listed above
(i) can react with hydrochloric acid but not with aqueous sodium hydroxide?
(ii) can react with aqueous sodium hydroxide but not with hydrochloric acid?
(iii) can react with both hydrochloric acid and aqueous sodium hydroxide?
(iv) cannot react with hydrochloric acid or aqueous sodium hydroxide?
22
Grid Question Section 3.2 Experimental introduction to acids and bases
28. Fill the table with acidic, basic, neutral or amphoteric.
Oxide
SO2
ZnO
CaO
CO
N2O
Acid-base nature
Basic Question- Section 3.3 Salt
29. Three methods are used for salt preparation
Method A: using a burette and an indicator.
Method B: mixing two solutions to obtain a salt by precipitation.
Method C: addition of an excess base or metal to a dilute acid then removing the excess by
filtration.
Method D: addition of an excess metal oxide to a dilute acid then removing the excess by
filtration.
Method E: addition of an excess metal carbonate to a dilute acid then removing the excess by
filtration.
For the following salt preparations, choose the method name any reagent needed and complete
or write the reactions.
a) The soluble salt, zinc sulfate from the insoluble base zinc oxide.
Method:
Reagent:
Word equation:
b) The soluble salt, potassium chloride, from the soluble base, potassium hydroxide.
Method:
Reagent:
Equation ……+……→ KCl(aq) + H2O(l)
c) The insoluble salt, Lead(II) iodide from the soluble salt Lead(II) nitrate.
Method:
Reagent:
Equation: Pb2+(aq) +……→……
d) Name the chemical needed to be added to insoluble magnesium oxide to form magnesium
sulphate.
e) Which method can be used to prepare the soluble salt potassium chloride from potassium
hydroxide?
f) Name the chemical needed to be added to lithium hydroxide to form lithium sulphate.
g) Which method can be used to prepare the insoluble salt silver chloride from the soluble
salt silver nitrate?
h) Name the chemical added to silver nitrate to form silver chloride.
i) Which method can be used to prepare the soluble salt magnesium sulphate from the
insoluble base magnesium oxide?
23
j) Name the chemical needed to be added to insoluble calcium oxide to form calcium nitrate.
Which method can be used to prepare the soluble salt sodium iodide from sodium
hydroxide?
k) Name the chemical needed to be added to sodium hydroxide to form sodium chloride.
l) Which method can be used to prepare the insoluble salt lead (II) iodide from the soluble
salt lead (II) nitrate?
m) Name the chemical added to lead (II) nitrate to form lead (II) iodide.
WEEK 6
Salt Preparation
1. Which of the following reagents can be added to hydrochloric acid, HCl, to prepare zinc
chloride, ZnCl2?
a) Zn
b) ZnO
c) ZnCO3
d) ZnSO4
2. Which of the following reagents can be added to sulphuric acid, H2SO4, to prepare copper (II)
sulphate, CuSO4?
a) Cu
b) CuO
c) CuCO3
d) Cu(OH)2
3. Which of the following reactions produce ammonia gas?
a) acid + base
b) ammonium salt + base
c) acid + metal carbonate
d) acid + metal
e) acid + metal sulfide
4. Which of the following reactions produces a pungent smelling gas that turns damp red litmus
blue?
a) CaO + H2SO4
b) Ca(OH)2 + H2SO4
c) Ca + H2SO4
d) Ca(OH)2 + (NH4)2SO4
e) CaCO3 + H2SO4
24
5. Which of the following reactions produce (s) a gas?
a) acid + base
b) ammonium salt + base
c) acid + metal carbonate
d) acid + metal
e) acid + metal sulfite
f) acid + metal oxide
6. The reaction of a strong acid with a strong base to produce salt and water is called:
7. During preparation of silver iodide, AgI, by precipitation, what is the correct order of the
steps?
a) mixing two aqueous solutions
b) filtration
c) stirring with a glass rod
d) drying residue in an oven
8. The pieces of apparatus needed during the preparation of sodium chloride, NaCl, by titration
are:
a) stop clock, burette, dropper
b) balance, burette, beaker
c) flask, burette, pipette
d) pipette, balance, burette
e) pipette, stop clock, balance
9. During the titration of an acid with a base, the color change that occurs when phenolphthalein
is used as an indicator is ________ to ________
10. During the titration of an acid with an base, the color change that occurs when methyl orange
is used as an indicator is ________ to ________
11. Dilute sulphuric acid is added to a mixture that contains magnesium, zinc and copper in a
beaker. The beaker is left for about 15 minutes and its contents are then filtered.
What does the filtrate contain?
25
12. Some reactions of substance M are shown below
magnesium
substance M
reacts with
→
hydrogen
sodium carbonate →
carbon dioxide
copper (II) oxide →
copper (II) chloride
Identify substance M?
13. The information in the box is about preparation of magnesium nitrate crystals.
Step 1: Add spatula measures of magnesium oxide to some hot dilute nitric acid then stir.
Step 2: Continue adding magnesium oxide until it is in excess.
Step 3: Remove excess magnesium oxide so that colorless magnesium nitrate is left behind.
Step 4: Evaporate the magnesium nitrate solution until it is saturated.
Step 5: Leave the saturated solution to cool, white crystals are seen formed on cooling.
Step 6: Remove the crystals from the solution.
Step 7: Dry the crystals on filter paper.
a)
b)
c)
d)
e)
Explain why excess magnesium oxide is used in step 2.
Explain how excess magnesium oxide is removed from the solution in Step 3.
What is meant by the term saturated solution?
What is the practical method that could show that the solution is saturated?
Explain why the crystals in Step 7 are dried using filter paper and not by heating.
WEEK 7
Basic Questions- Section 3.5 Acid-base titrations
30. Suppose that 0.098 mol of solid NaOH is added to 0.100 L of 1.00 M HCl.
a) How many more moles of HCl are present in the solution than moles of NaOH?
b) From the excess number of moles and the volume, calculate the concentration of excess
H+(aq).
c) Calculate the concentration of OH–(aq) at equilibrium.
31. 200 mL of 1.00 M NaOH is added to 200 mL of 1.00 M HCl.
Calculate the final [H+] and [OH–].
26
Grid Question- Section 3.5 Acid-base titrations
32. a) Calculate the [H+] and the [OH–] in a 0.010 M solution of HCl.
b) Calculate the [H+] and the [OH–] in a solution of 0.40 g of NaOH dissolved in 10 cm3 of
solution.
Basic Question- Section 3.7 Some common acids and bases
33. Write the reaction between:
a) HCl(g) and water.
b) HNO3 and water.
c) H2SO4 and water. Show both steps of the dissociation.
d) CH3COOH(aq) and water.
e) H3PO4 and water. Show the three steps of the dissociation.
f) Where more than one proton may be released, in which step is the acid strongest? Weakest?
WEEK 8
Basic Questions- Section 3.9 Adding strong acid and bases to salts
34. Complete the equations:
a) CaO(s) + H2O(l)
b) SO3(g) + H2O(l)
c) CO2(g) + H2O(l)
d) SO2(g) + H2O(l)
e) K2O(s) + H2O(l)
f) Na2O(s) + H2O(l)
35. Write down the complete acid-base equations of the following:
a) Neutralization of an acid and a base
b) Carbonate ion in water
c) Sodium carbonate with a strong acid (net ionic reaction)
d) Ammonium ion in water
e) Strong base and ammonium salt
f) CaO(s) + 2H+(aq)
g) SO3(g) + NaOH(aq) (net ionic reaction)
36. If 23 g of formic acid, HCOOH, are dissolved in 10.0 L of water at 20°C, the [H+] is found to
be 3.0 × 10–3 M. Calculate Ka.
37. A chemist dissolved 25 g of CH3COOH in enough water to make 1 L of solution. What is the
concentration of this acetic acid solution? What is the concentration of H+(aq)? Assume a
negligible change in [CH3COOH] because of dissociation to H+(aq).
38. Potassium nitrate is a salt that dissolves in water in an endothermic process. What happens to
the temperature and pH of the water in which it is dissolved?
27
Grid Question- Section 3.9 Adding strong acid and bases to salts
39. Calculate the [H+] and [OH–] in a solution whose:
a. pH = 5.0
b. pH = 8.0
c. pH = 7.0
d. Are the above solutions acidic, basic or neutral?
40. Write down the net ionic reaction of sodium carbonate with a strong acid (net ionic reaction).
Top Questions
1. An oxide X reacts with both potassium hydroxide and sulfuric acid.
Which of the following is true about X?
A. X is an acidic oxide.
B. X is a basic oxide.
C. X is an amphoteric oxide.
D. X is a neutral oxide.
E. none of the above
2. Element X is on the left-hand side of the Periodic Table. Which of the following is true about
its oxide?
[-A-]
[-B-]
[-C-]
[-D-]
[-E-]
Type of oxide
metallic
metallic
non- metallic
non- metallic
metallic
Nature of oxide
acidic
basic
acidic
basic
neutral
3. Which of the following is the best method to prepare ZnSO4 from H2SO4 and ZnO?
a) titrate ZnO(s) against H2SO4.
b) Mix equal amounts of the two reagents and crystallize.
c) Add ZnO to excess acid solution.
d) Add excess ZnO to acid solution.
e) All of the above work.
4. A mixture of NH4Cl and Ca(OH)2 solution are heated, a gas is given off. The gas has a pungent
smell. It was tested using wet red litmus paper. Identify the gas and the final color of the litmus.
The gas is Ammonia NH3 and it turns red litmus into blue.
28
5. a) A student titrates an aqueous sulfuric acid solution using 25 cm3 of 0.050 M solution
calcium hydroxide and few drops of an acid-base indicator.
The volume of sulfuric acid needed to react completely with the calcium hydroxide solution
is 12.5 cm3.
The equation of the reaction is shown:
H2SO4(aq) + Ca(OH)2(aq)  CaSO4(s) + 2H2O(l)
What is the concentration of sulfuric acid solution?
 Calculate the number of moles of calcium hydroxide.
 Calculate the number of moles of sulfuric acid to react completely with calcium
hydroxide.
 Calculate the concentration of the sulfuric acid solution.
b) After titration is done, the final solution in the beaker contains solid calcium sulfate, water,
and some dissolved acid-base indicator.
Describe how to extract dry sample of solid calcium sulfate from the solution. Write the
process including all key steps.
c) The student conducted the following tests on the products of the reaction between sulfuric
acid and sodium carbonate.
Describe the student’s observation in each test.
i. Flame test on the solution.
ii. Passing gas released in lime water.
iii. Addition of aqueous solution of barium nitrate.
29
d) Dilute sulfuric acid reacts with metals, metal oxides and bases.
Write chemical equations for the reaction of dilute hydrochloric acid with each of the
following:
i. magnesium.
ii. sodium oxide
iii. potassium hydroxide
6. A student conducts an acid-base titration to determine the concentration of a sodium hydroxide
solution using a 0.1 M hydrochloric solution. He is provided with the following apparatus.
a) Name the items that the student needs to perform the titration.
b) The student adds few drops of an acid-base indicator to 20 cm3 of sodium hydroxide and
titrates the mixture with the hydrochloric acid solution.
i. Why is it necessary to add an acid-base indicator before starting the titration?
ii. The student makes two reading to determine the volume of the hydrochloric acid needed
to neutralize completely the sodium hydroxide solution. The initial reading he made
before starting the titration is 10.00 cm3 and the final reading at the end of the titration is
23.00 cm3.
What is the volume of hydrochloric acid used to completely neutralize 20 cm3 of sodium
hydroxide solution?
c) The student repeats the same experiment using the same materials. He discards the content
of the final solution from the flask and refills it directly with another 20.00 cm3 of the same
sodium hydroxide.
What did the student do wrong in this procedure? Describe the correct procedure.
7. Plan an experiment to find the concentration of an aqueous solution of sulfuric acid provided
with dilute solution of sodium hydroxide of known concentration and common laboratory
apparatus.
30
8. Copper sulphate solution was prepared by reacting excess copper oxide with dilute sulphuric
acid. The diagram below shows the method used.
a) Fill in the empty boxes.
b) Explain what is meant by excess.
c) Draw a diagram to show how the mixture is filtered.
31
Chapter 4 Oxidation-Reduction Reactions
WEEK 9
Basic Question- Section 4.3 Electrochemical cells
41. These elements are listed in decreasing tendency to lose electrons. Which of these reactions
below takes place spontaneously?
Zn(s) → Zn2+(aq) + 2e–
H2(g) → 2H+(aq) + e–
Cu(s) → Cu2+(aq) + 2e–
Ag(s) → Ag+(aq) + e–
a) Cu(s) + Zn2+(aq) → Zn(s) + Cu2+(aq)
b) Zn(s) + Cu2+(aq) → Zn2+(aq) + Cu(s)
c) Cu(s) + H+(aq) → Cu2+(aq) + H2(g)
d) H2(g) + Cu2+(aq) → Cu(s)+ 2H+(aq)
e) 2Ag(s) + 2H+(aq) → 2Ag+(aq) + H2(g)
f) Zn(s) + 2Ag(s) → Zn2+(aq) + 2Ag+(aq)
WEEK 10
Grid Question- Section 4.3 Electrochemical cells
42. A zinc rod is placed in a 1.0 M zinc nitrate solution. A copper strip is placed in a 1.0 M copper
nitrate solution. The two half-cells are connected by a salt bridge and, externally, by a
voltmeter. The voltmeter reads 1.10 V, and it shows that copper is the positive terminal of the
cell.
a) Which electrode is the cathode?
b) Write the equation of the net reaction that takes place at the cathode.
c) Which electrode is the anode?
d) Write the equation of the half-reaction that takes place at the anode.
e) In what direction does the current pass in the solution?
f) How much energy is released by the cell when 2.0 C pass by any point in the circuit?
32
Basic Questions- Section 4.5 Electromotive force of a cell
43. Using the table of standard reduction half-cell potentials, answer the following:
a) What would happen if an aluminum spoon is used to stir Fe(NO3)2 solution?
b) What would happen if an iron spoon is used to stir an AlCl3 solution?
c) Can 1 M Fe2(SO4)3 solution be stored in a container made of nickel metal? Explain your
answer.
44. A half-cell consisting of a cadmium rod dipping into a 1.0 M Cd(NO3)2 solution is connected
with a standard hydrogen half-cell. The cell voltage is 0.40 volt and the platinum electrode in
the hydrogen half-cell is the cathode. Determine E° for the reaction:
Cd → Cd2+ + 2e–
45. a) Predict the Eºnet for the cell made up of the standard half-cells of zinc and silver.
b) How does the current flow in the outside circuit?
c) Which electrode is the cathode?
d) Which electrode gains weight?
e) When this electrode gains 0.0010 mol, how many grams will the other electrode lose in
weight?
WEEK 11
Grid Question- Section 4.5 Electromotive force of a cell
46.
a. Predict the E°net for the cell made up of the standard half-cells of copper and bromine. Given:
Cu → Cu2+ + 2e- Eº = -0.34 V
2Br- → Br2 + 2e- Eº = -1.06 V
b. Which electrode is the positive terminal of the cell?
c. How do the positive and negative ions flow in the solutions?
d. Do you expect a copper metal to react with bromide ions?
e. Do you expect copper ions to react with bromine?
WEEK 12
Basic Questions- Section 4.6 Oxidation numbers-an electron bookkeeping device
47. Which of the following is true? The oxidation number of nitrogen increases as:
a) NO3– becomes NO
b) N2O4 becomes NI3
c) CN– becomes HCN
d) NH3 becomes NH4+
e) NO2 becomes N2O5
48. When copper metal is placed in a dilute solution of nitric acid bubbles of NO gas are
produced. The solution turns blue, indicating that Cu2+ is forming. Write a balanced equation,
showing all steps.
33
Grid Question- Section 4.6 Oxidation numbers-an electron bookkeeping device
49. Find the oxidation number of each element in
H3PO4, HSO4–, C2H3O2–, Pt(H2O)42+, Fe2(CO3)3
50. Which of the following equation is an oxidation reaction?
a. CuCO3 → CuO + CO2
b. ZnO + 2HCl → ZnCl2 + H2O
c. 2Cu + O2 → 2CuO
d. Na2CO3 + 2HCl → 2NaCl + H2O + CO2
Basic Question- Section 4.7 Electrolysis
51. Write the two half-reactions occurring when molten ZnCl2 is electrolyzed.
Grid Question- Section 4.7 Electrolysis
52. The following are results for the experiments on electrolysis using inert electrodes.
Complete the table below, the first one has been done for you.
Electrolyte
Molten lead(II)
chloride
Change at the
negative electrode
(cathode)
Change at the
positive electrode
(anode)
Change to
electrolyte
Lead formed
Chlorine formed
Used up
Potassium formed
Iodine formed
Used up
Hydrogen formed
Bromine formed
Potassium hydroxide
formed
Dilute aqueous
sodium chloride
Aqueous copper (II)
sulphate
34
WEEK 13
Basic Questions- Section 4.8 Applications of electrochemical cells and electrolysis
53. Concentrated hydrochloric acid is electrolyzed using the apparatus below.
a) Label on the diagram the position of the electrodes.
b) What are the two observations when the circuit is switched on?
i) The product at the positive electrode is called:
ii) The test and the result of this product are:
Test:
Result:
54. Which of the following aqueous solutions produces H2(g) and O2(g) during electrolysis?
a) KI(aq)
b) CuI2(aq)
c) molten NaOH
d) CuSO4(aq)
e) K2SO4(aq)
55. An iron spoon is to be plated with copper metal. Write the two half-reactions occurring at each
electrode.
35
56. Which of the following apparatus could be used to electroplate an iron nail with nickel?
Grid Question- Section 4.8 Applications of electrochemical cells and electrolysis
57. Give the chemical reaction that produces electricity in the fuel cell.
58.
a. A current of one ampere is passed through a voltammeter containing concentrated aqueous
sodium chloride with inert electrodes for 32 minutes. Find the volume of each gas liberated at
room conditions.
b. Calculate the mass of lead deposited and the volume of iodine gas released at STP when 2
faradays (2 F) of electricity flow through molten lead (II) iodide, PbI2.
c. A current of 0.020 A passes for 5.00 hours through a solution of gold (III) nitrate. Calculate the
mass of metal deposited.
d. A metal has a relative atomic mass of 48.0 was deposited by electrolysis. If 0.239 g of the metal
was deposited when 0.100 A flowed for 4.00 hours, what is the charge on the ion of this metal?
e. A current of 2.00 A was passed through a solution of nitric acid for 6.00 hours. Calculate the
volume of hydrogen and oxygen gases produced measured at STP conditions.
36
Top Questions
1. Different metals are used as electrodes in a galvanic cell.
Which of the following pairs of metals would produce the largest voltage?
A silver and copper
B. platinum and zinc
C. zinc and copper
D. zinc and magnesium
E. platinum and copper
2. Which of the following electrolytic cells forms gases at both of its inert electrodes?
A. Cell 1: aqueous sodium iodide
B. Cell 2: aqueous copper(II) sulfate
C. Cell 3: aqueous sodium chloride
D. Cell 4: molten sodium chloride
E. Cell 5: potassium iodide
3. Which of the following represent(s) a reduction change?
1. Fe  Fe2
2. H 2 O2  H 2O
3. Cr2 O 24  Cr2O 72
A. 1 and 2 only
B. 1 and 3 only
C. 1 only
D. 2 only
E. 2 and 3 only
4. a) A student assembled the following electrochemical cell.
i. What is the direction of electron flow in this electrochemical cell?
ii. The student replaced the iron electrode with a silver electrode. The voltmeter reading
changed. Describe the change and explain why this change occurred.
37
b) A student conducts electrolysis of concentrated aqueous solution of sodium chloride using
the following apparatus.
i. Which particles allows the flow of the electric current in the sodium chloride solution?
ii. Give two reasons for the use of graphite as the electrodes.
iii. What is the gas produced at the anode?
iv. Write the ionic half-equation for the reaction occurring at the anode. Include state
symbols.
v. A blue litmus paper is dipped in the remaining solution after electrolysis is over.
Does the blue litmus paper change color? Explain why.
5. Electrolysis is used to break down ionic compounds.
a) A student conducts an electrolysis for a concentrated solution of copper(II) chloride.
i. Why is it recommended to use inert electrodes in electrolysis?
ii. A solid deposit is observed at the negative electrode. Identify this solid.
iii. Write the ionic half equation for the reaction occurring at the negative electrode.
iv. A fizzing is observed at the positive electrode. Identify the gas released and explain how
it forms.
v. Identify the species that is reduced in the solution undergoing electrolysis.
b) A student wants to electroplate an iron rod with chromium.
i. At which electrode the student should place the chromium rod?
ii. Suggest a suitable electrolyte for this electrolysis?
iii. Write the ionic half equation occurring at the negative electrode.
iv. At which electrode oxidation occurs?
v. Give a reason why chromium is used in electroplating.
38
6. A student conducts electrolysis on dilute solution of sodium chloride using the apparatus shown
below.
a) Gas X makes a pop sound when tested with a lighted splint.
Identify gas X.
b) Suggest a test for oxygen gas. Include result of the test in your answer.
c) Suggest a material to be used as an electrode and give one reason why it is suitable.
7. The below apparatus is used to break down molten lead bromide.
a) What is the process shown in the figure called?
b) What are the expected observations on the positive electrode (anode) and on the negative
electrode (cathode)?
Observation at the positive electrode:
Observation at the negative electrode:
c) A Bunsen burner is used to melt solid lead bromide. How can the student adjust the Bunsen
burner to give the hottest flame?
d) Why is it recommended to use graphite electrodes in the above apparatus?
e) When replacing molten lead bromide with an aqueous solution of lead bromide in the above
apparatus, hydrogen gas is produced at the negative electrode and bromine gas is produced at
the positive electrode.
i. Suggest a test for hydrogen gas. Include observation.
ii. How can you identify bromine gas?
39
f) Lead bromide is classified as hazardous chemical. Suggest a safety precaution when dealing
with lead bromide.
8. A student conducted two titration experiments to investigate the reaction between aqueous
solution of potassium manganate(VII) and two aqueous solutions of iron(II) sulfate, solution R
and solution T, of different concentrations.
Experiment 1
A 50 cm3 burette was filled with potassium manganate(VII) solution to the 0.0 cm3 mark.
25 cm3 of solution R were measured using a pipette and poured into a conical flask.
Potassium manganate(VII) solution was added to the flask while shaking it until a pale pink
color appears.
a) The burette showed the following measurement.
i. Record the final burette reading.
ii. Calculate the volume of potassium manganate(VII) solution used in cm3.
Experiment 2
25 cm3 of solution T was measured using a measuring cylinder and experiment 1 was
repeated. But this time the burette was not filled to the 0.0 cm3 mark.
b) The initial and final readings of the burette are shown.
i. Record the final reading in cm3.
ii. Calculate the volume of potassium manganate(VII) solution used in cm3.
40
c) Which solution is more concentrated solution R or solution T? Explain
d) i. If in experiment 2, the student used 50 cm3 of solution T, instead of 25 cm3, calculate the
volume of potassium manganate(VII) needed to reach the end point.
ii. What issue will the student face when titrating 50 cm3 of solution T?
e) Another student conducted the same experiments using a measuring cylinder to measure
25 cm3 of solutions R and T instead of a pipette.
How would his results differ from that of the student who used the measuring cylinder?
f) How can the student reach more reliable results in this experiment?
g) Indicators are used in most titrations to determine when the end point is reached. Why no
indicator was used in this titration?
41
Topics – Chapter 1
Experimental Techniques
Basic Questions - Section 1.2 Separating techniques
59. Define an element, a compound, and a mixture.
60. Classify the following as elements, compounds, or mixtures.
Chemical
Blood
Oxygen
Ammonia
Orange juice
Element
Compound
Mixture
61. Briefly describe how to separate a mixture of a salt and sand into its constituents.
62. What process can be used to separate the dyes in food coloring?
63. a. The following techniques are used to separate mixtures:
Filtration, fractional distillation, crystallization, simple distillation, diffusion
From the above list, choose the most suitable technique that can be used to separate the following:
Water from aqueous copper (II) sulphate.
Helium form a mixture of argon and helium
Copper (II) sulphate crystals from aqueous copper (II)
sulphate
Ethanol from aqueous ethanol
Barium sulphate from a mixture of barium sulphate and
water
Liquefied air into its components
b. Give the name of an apparatus or the name of a technique you can use to separate the following
mixtures into their constituents:
 A mixture of sulfur and steel.
 A mixture of cooking oil and water.
42
Grid Questions- Section 1.2 Separating techniques
64. Give the name of the most suitable technique you can use in order to obtain crystals of
potassium nitrate from its aqueous solution.
65. Which of the following mixture(s) is/are best separated by simple distillation?
a) sodium bromide aqueous solution
b) sand, salt, and water mixture
c) water-alcohol solution
d) sugar aqueous solution
66. Fractional distillation is used to separate a mixture of water and ethanol.
The diagram below is used to carry out the separation in the laboratory.
a. Name apparatus A, B and C.
b. State the function of D.
c. How can you check the purity of ethanol collected?
d. Which of the following mixtures is best separated by fractional distillation?
a) Ethanol-water solution
b) Oil-water mixture
c) Salt water solution
d) Sugar alcohol solution
43
67. The following shows the chromatogram of a dye that was separated into two components A
and B using paper chromatography in an eluent X:
Explain the difference between the distances travelled by the two spots.
68. Describe how the chromatography procedure should be adapted to analyze colorless mixtures.
Basic Question- Section 1.3 Measuring physical constants
69. You are given a sample of colorless liquid in a test tube and told that this is water.
How can you assess the purity of the given sample?
Section 1.4 Working with gases
Basic Questions- Section 1.3 Measuring physical constants
70. a) Calculate the molecular masses of H2S and NH3. Compare their densities under identical
conditions. Report your answer as a simple fraction.
b) Which of the above two gases is collected by upward delivery, and which is collected by
downward delivery?
71. Name two common drying agents that are used to dry gases.
44
Top Questions
1. The table below provides the Rf values of three amino acids.
Amino acid
lysine
glycine
proline
valine
Rf value
0.14
0.26
0.43
0.60
The following diagram shows a chromatogram for a mixture of two amino acids R and T.
What are the identities of the amino acids separated?
A
B
C
D
E
R
proline
lysine
valine
valine
glycine
S
valine
proline
glycine
proline
proline
45
2. The apparatus shown is used to collect gases in the lab.
a) Name the apparatus A, B and C.
b) A student used this apparatus to collect sulfur dioxide gas.
i. Explain why the student made a mistake using this apparatus to collect sulfur dioxide gas.
ii. Name one gas that can be collected using this apparatus.
c) Suggest a modification on the apparatus to collect sulfur dioxide gas. You may include a
diagram in your answer.
3. A student investigated the rate of reaction between dilute hydrochloric acid and excess calcium
carbonate.
The student used a gas syringe to collect and measure the amount of gas released as shown in
the diagram.
As the calcium carbonate sample was dropped into the hydrochloric acid solution, the volume
of carbon dioxide gas was monitored every 30 seconds 25°C.
46
a) Use the syringe diagram to record the volume of carbon dioxide gas released.
Time
/ sec
0
30
60
90
120
150
180
210
Syringe diagram
Total volume of
gas collected /
cm3
47
b) Plot the results on the grid and draw a smooth line curve.
c) Use the graph and deduce the volume of carbon dioxide gas released 80 seconds. Show clearly
on the graph how you deduced your answer.
d) If instead of using a gas syringe, another student collected the gas using an inverted 50 cm3
measuring cylinder filled with water.
i. Suggest an error in this experiment.
ii. Suggest an improvement to fix the error (other than using a gas syringe).
e) The following expression is used to calculate the average rate of the reaction.
volume of gas collected / cm3
average rate of reaction =
time / sec
i. Calculate the volume of hydrogen gas collected between 30 seconds and 120 seconds.
ii. Calculate the average rate between 30 seconds and 120 seconds. Give your answer to two
significant figures and include the unit.
f) Sketch on the same graph the expected curve if the concentration of hydrochloric acid was
increased.
48
g) If the gas to be collected was hydrogen rather than carbon dioxide would it have been possible
to collect it over water using a 250 cm3 measuring cylinder? Explain.
4.
a) A student was asked to measure the time taken for 3.00 g of calcium carbonate to react
completely with excess of dilute sulphuric acid.
What pieces of apparatus did the student use?
[-A-]
[-B-]
[-C-]
[-D-]
[-E-]
Pipette, balance and thermometer.
Pipette, clock and balance.
Clock, pipette and thermometer.
Clock, thermometer and balance.
Clock, balance, thermometer and pipette.
b) A student carried out an experiment to find out how fast 4 cm pieces of magnesium ribbon
dissolve in
25 cm3 samples of hydrochloric acid at different temperatures.
Which piece of apparatus did the student not use?
[-A-]
[-B-]
[-C-]
[-D-]
[-E-]
Measuring cylinder
Thermometer
Balance
Clock
thermometer
5. The boiling point of liquid A is less than that of water. To test a student, a teacher covers the
numbers on the thermometer. The student then places the thermometer in boiling liquid X.
What is the temperature reading on the thermometer?
[-A-]
[-B-]
[-C-]
[-D-]
[-E-]
105.5◦C
65.5◦C
83.5◦C
104.5◦C
54.5◦C
49
6. The diagram shows the diffusion of hydrogen bromide and ammonia in a glass tube.
The gases are given off by the solutions at each end of the tube. When hydrogen bromide and
ammonia mix they produce a white solid, ammonium bromide. Which line shows where the
white solid is formed?
A
B
C
Cotton wool soaked
in concentrated ammonia
solution
D
Cotton wool soaked
in concentrated hydrobromic acid
solution
7. The method that can be used to obtain crystals of aqueous copper (II) sulphate is:
[-A-]
[-B-]
[-C-]
[-D-]
[-E-]
electrolysis
fractional distillation
neutralization
chromatography
crystallization
8. What is the correct order of steps required to separate salt from a mixture of sand?
[-A-]
[-B-]
[-C-]
[-D-]
[-E-]
shake with water → filter → evaporate
shake with water → evaporate → filter
shake with water → decant → boil
filter → evaporate → shake with water
filter → shake with water → evaporate
9. Fractional distillation and chromatography are techniques used for separation of compounds.
Which method requires the use of thermometer to check that complete separation is done?
[-A-]
[-B-]
[-C-]
[-D-]
[-E-]
Chromatographic separation of two colored liquids.
Chromatographic separation of two colorless solids.
Fractional distillation of two colored solids.
Fractional distillation of two colorless liquids.
Fractional distillation of two colorless solids.
50
10. An experiment was done in the laboratory to separate a mixture of three dyes. On the
chromatography paper, a line was drawn and a spot of the mixture of dyes was placed on it. The
paper was then dipped in a suitable solvent for several minutes.
Which of the following statements is true about the above experiment?
[-A-]
[-B-]
[-C-]
[-D-]
[-E-]
The line is drawn in ink not in pencil.
The line is placed below the level of the solvent.
The dyes have different solubilities in the solvent.
The dyes have the same solubilities in the solvent.
The dyes have different boiling points.
11. A sample of protein contains amino acids that can be identified by chromatography.
The diagram below shows the chromatogram obtained when four samples of amino acids were
analyzed. The paper was then sprayed with ninhydrin.
a. Why is the original line drawn in pencil?
b. Which amino acids are likely to be the same?
c. Which sample contains more than one amino acid and why?
d. Why is it important to spray the chromatogram with ninhydrin?
51
12. The colors present in blackcurrant sweets can be separated by chromatography.
The colors in the sweets are said to be water-soluble dyes.
The diagrams below represent how the colors can be extracted from the sweets.
a.
The diagram below shows how chromatography is done.
b.
i. Name the solvent used.
ii. Label the origin on the diagram.
c. Sketch the chromatogram you would get if three different colours were found in the sweets.
52
13. Sulphur dioxide is soluble in water and denser than air.
A sample of this gas can be prepared by adding dilute hydrochloric acid to sodium sulphite and
warming the mixture. The diagram below represents the apparatus used.
a.
i. Name the chemicals used.
ii. show on the diagram where heat is supplied
b. Identify and explain two mistakes in the diagram.
14. The diagram below shows results on a paper chromatogram of four substances V, X, Y and Z.
Which two substances are pure?
[-A-]
[-B-]
[-C-]
[-D-]
[-E-]
V and Y
X and V
X and Z
Y and Z
V and Z
53
15. A teacher placed a small amount of liquid bromine in the bottom of a sealed gas jar of air. After
three minutes red-brown fumes were seen just above the liquid surface. After one hour the redbrown colour had spread completely throughout the gas jar.
Use the kinetic particle model of matter to explain these observations.
54
Chapter 2 Hydrogen, oxygen, and air
Basic Questions- Section 2.1 Hydrogen
72. What is the test for hydrogen gas.
73. a) Write the equation of the reaction of lithium with liquid water.
b) Write the equation of the reaction of magnesium with sulfuric acid.
c) Write the equation of the reaction of zinc with steam.
d) Consider the process of electrolysis.
i) Name the two gases that are produced when brine is electrolyzed.
ii) Write equations for the reactions that take place at each electrode.
74. a) Write the equation of the reaction of hydrogen with chlorine.
b) Hydrogen reduces lead(II) oxide, PbO, into lead. Water is produced in the process. Write
the equation of this reaction. Do not include state symbols.
75. List three common uses of hydrogen.
76. Which of the following is an industrial process for the production of hydrogen?
i.
Thermal catalytic cracking of long chain alkanes
ii.
Electrolysis of acidified water
iii. Reaction between methane (natural gas) and steam.
Basic Questions- Section 2.2 Oxygen
77. a) To which group in the periodic table does oxygen belong? Deduce the number of valence
electrons of oxygen.
b) How do you test for oxygen gas?
78. Write an equation to represent the decomposition of hydrogen peroxide. Name a catalyst that
can be used in the process.
79. a) Write an equation to represent the reaction of oxygen with the following elements: Na, Ca,
Mg, S.
b) Write two equations to show the complete and the incomplete combustion of carbon.
55
80. Complete the following table to identify the properties of each oxide.
Oxide
Acidic
Basic
Class
Neutral
Amphoteric
Reacts
with acid
Reacts
with base
CO2
CaO
PbO
CO
H2O
Li2O
Al2O3
81. List two common uses of oxygen.
Basic Questions- Section 2.3 Air
82. Is air considered a compound or a mixture? Give one reason to justify your answer.
83. a) List three common uses of air.
b) How are the different raw materials extracted from air?
c) Which two gases are the main contributors to acid rain?
Grid Questions- Section 2.3 Air
84. Give the composition of air under normal conditions.
85. List some air pollutants, with their sources.
56
Top Questions
1. The below apparatus includes a glass tube containing copper powder connected to syringe A
holding 80 cm3 of air and empty syringe B.
The tube containing copper is heated while pushing A’s and B’s plungers for several times until
the copper turns black. At the end of the experiment, the remaining gas mixture is collected in
syringe A.
Which diagram shows the amount of gas remaining in syringe A at the end of the experiment,
knowing that copper is in excess?
A.
B.
C.
D.
E.
2. Clean air is a mixture of gases containing O2, N2, CO2 and Ar gases.
a) Which gas makes up 78% of clean air?
b) Liquids and gases are separated from the mixture based on the difference in their boiling
points. Name the technique used to separate components of air?
c) Many pollutants like oxides of sulfur and nitrogen are formed from burning fossil fuels.
i. Write the chemical formula for the sulfur containing air pollutants.
ii. Why the gas containing sulfur is considered an air pollutant?
iii. Suggest a method to reduce the gas containing sulfur from emissions of power stations.
Name the chemical used.
d) Catalytic converters are installed in exhausts of modern cars to absorb pollutants like oxides
57
of nitrogen, carbon monoxide and unburnt hydrocarbons.
i. Name two oxides of nitrogen formed in car engines?
ii. Carbon monoxide is formed from incomplete combustion of hydrocarbons like butane,
C4H10.
Write chemical equation for the incomplete combustion of butane.
iii. What is the role of transition metals found in catalytic converters?
e) Carbon dioxide gas is absorbed from the atmosphere by dissolving in oceans and by another
natural process X, involving plants.
i. Name the process X.
ii. Write chemical equation for process X.
iii. Name two processes, one natural and another one artificial, responsible of the emission of
carbon dioxide into the atmosphere.
f) Carbohydrates are polymers that break down by hydrolysis.
i. Name the simplest product of the hydrolysis of the carbohydrate, starch.
ii. State two ways to break down carbohydrates.
iii. Partial hydrolysis of a given carbohydrate gives a mixture of sugars. Suggest a method to
identify these sugars?
3. Rain water is said to be slightly acidic due to the presence of carbon dioxide dissolved in it.
The word equation showing the acidity of acid rain is: carbon dioxide + water → carbonic acid
The acidity of acid rain is due to other pollutant gases in the atmosphere that also dissolve in water
a) What is the acid formed when sulfur dioxide reacts with water?
b) Name another pollutant gas that leads to the formation of acid rain.
c) Acid rain leads to two problems, these are ……
4. The diagram shows a limestone column in an industrial town. Limestone is largely calcium
carbonate.
iron pins
column when first built
the same column after 40 years
Describe and explain the changes to the column over 40 years. In your answer refer to
 the change to the limestone,
 the name of a pollutant causing this change,
 the chemistry involved in this change.
58
5. Nowadays, hydrogen is being used as a fuel for cars to reduce air pollution. Petrol that is used
as fuel for cars is the major source of air pollution. This pollution can be reduced using the
reaction below:
2CO + 2NO → N2 + 2CO2
a) Where in the car does this reaction take place?
b) How is carbon monoxide formed in the car engine?
c) Why does the hydrogen powered car produce less pollution?
59
Chapter 3 Water
Basic Question- Section 3.2 Water pollution
86. Name one natural phenomenon that can cause water pollution.
Basic Question- Section 3.3 Water treatment
87. Briefly state how water is treated.
Grid Question- Section 3.3 Water treatment
88. Study the diagram below.
a) At which stage or stages is chlorine added?
b) At which stage screening/ filtration is performed?
c) What do we call the last stage?
Grid Question- Section 3.4 Test for water
89. When few drops of water are added to anhydrous copper(II) sulfate, the color change observed
is from ____a_____ to _____b______
When few drops of water are added to anhydrous cobalt(II) chloride, the color change observed
is from ____c_____ to _____d______
Basic Question- Section 3.6 Chemistry of water
90. a) Write an equation for the reaction of sodium with liquid water.
b) Write an equation for the reaction of calcium with liquid water.
c) Write an equation for the reaction of magnesium with steam.
d) i) What two acids are produced when chlorine dissolves in water?
ii) Write an equation to represent the above reaction.
Basic Question- Section 3.7 Solution
91. a) Define the following terms: solution and aqueous solution.
b) Explain what is meant by the “Solubility of a solute in a particular solvent.”
60
Top Question
1. Polluted water undergoes many steps, in a water treatment, before it becomes suitable for
drinking.
Which of the following steps is used to remove insoluble impurities from polluted water?
A. filtration
B. coagulation
C. addition of chlorine
D. aeration
E. none of the above
61
Chapter 9 The halogens
Basic Questions- Section 9.2 Chlorine
92. a) What is the color of chlorine gas? Is it lighter or denser than air?
b) How do you test for chlorine gas?
93. a) Write the anode and cathode half reactions that take place in the electrolysis of molten
sodium chloride.
b) Chlorine gas is produced industrially by electrolysis in a diaphragm cell. Answer the
following questions:
i) What material is used as the electrolyte?
ii) What type of energy is used in electrolysis?
iii) State the materials from which the anode and cathode are usually made.
iv) Write the anode and cathode half reactions that take place during the electrolysis process.
v) Besides chlorine and hydrogen gases, what important byproduct is produced?
94. How do the halogens react with hydrogen gas?
95. List the three major uses of chlorine gas.
Basic Questions- Section 9.3 Other halogens
96.
a) Are the halogens monatomic or polyatomic molecules?
b) Give the physical state and color of each of them at room temperature.
c)
Halogen
Physical state at room temperature
Color
Fluorine
Chlorine
Bromine
Iodine
Astatine
97. Astatine is an element in group VII of the Periodic Table. It has been produced in very small
amounts. Which of the following is the best description for the properties of astatine?
Colour
State
Reaction with aqueous potassium iodide
[-A-]
colorless
solid
no reaction
[-B-]
black
solid
no reaction
gas
brown color
[-C-] dark brown
[-D-]
green
solid
no reaction
[-E-]
yellow
liquid
brown colour
62
98. Write the net ionic equation for the reaction that takes place when a few drops of silver nitrate
solution are added separately to the aqueous solutions listed below.
State the observations you expect to make in each case.
a) Sodium bromide
b) Potassium iodide
c) Sodium chloride
d) Potassium fluoride
99. a) Complete and balance the following ionic equations. Where no reaction is expected, write
NR.
i) Cl2(g) + Br-(aq) →
ii) I2(aq) + Br-(aq) →
iii) F-(aq) + Br2(aq) →
iv) I-(aq) + Cl2(aq) →
b) What are the colors of bromine and iodine solutions in cyclohexane (an organic solvent
similar to CCl4)?
Top Question
1. Which sequence best describes the relative reactivity of halogens from most reactive to least
reactive?
A. chlorine, bromine, iodine
B. chlorine, iodine, bromine
C. bromine, chlorine, iodine
D. bromine, iodine, chlorine
E. iodine, bromine, chlorine
63
Chapter 11 Qualitative chemical analysis
Basic Questions- Section 11.2 Identifying cations
100. a) How can you verify the presence of H+ ions in a solution using a pH meter?
b) How can you verify the presence of H+ ions in a solution using the materials listed below?
Write the net ionic equation for each reaction and report any expected observations.
i) Sodium carbonate
ii) A piece of magnesium ribbon
c) How are the ions of lithium, sodium, and potassium identified?
d) A flame test is conducted on a sample of calcium chloride. What color will the flame have?
e) How can you test for calcium ions using a concentrated solution of sodium hydroxide? Write
the net ionic equation for the reaction and report the expected observation.
101. How can you test for the following ions using sodium hydroxide solution? Write the net ionic
equation for the reaction, and report any expected observations.
a) Al3+
b) Zn2+
102. a) How can you test for the following ions using sodium hydroxide solution? Write the net
ionic equation for the reaction, and report the expected observations.
i) Fe2+
ii) Fe3+
iii) Cu2+
b) How can you test for copper (II) ions using ammonia solution? Report the expected
observations. (No equations are necessary.)
103. How can you identify ammonium ions using a solution of sodium hydroxide?
Write the net ionic equation for the reaction, and report any expected observations.
Basic Questions- Section 11.3 Identifying anions
104. a) How can you verify the presence of OH- ions in solution using a pH meter?
b) How can you test for the following ions using silver(I) nitrate solution? Write the net ionic
equations for the reactions and report the expected observations.
i) Clii) Briii) I105. How can you test for the following ions using hydrochloric acid solution? Write the net ionic
equation for the reaction, and report the expected observations.
i) CO32ii) NO3iii) SO32106. How can you test for SO42- ions in a solution? Write the net ionic equation for the reaction,
and report the expected observations.
64
Basic Questions- Section 11.4 Identifying gases
107. Describe how each of the following gases are tested. Write an equation for the expected
reaction.
Test for Gases
H2
Hydrogen gas
O2
Oxygen gas
NH3
Ammonia gas
CO2
Carbon dioxide
gas
Cl2
Chlorine gas
SO2
Sulfur dioxide
H2O
Water vapour
OR liquid
65
Grid Question- Section 11.4 Identifying gases
108. Complete the below table with the correct observation.
Gas
Color
Odor
Effect on litmus
paper
Cl2
H2
O2
NH3
Basic Question- Section 11.5 Identifying water
109. How can you test for the presence of water in a sample of colorless liquid? Write an equation
for the expected reaction.
Basic Question- Section 11.6 Distinguishing between alkanes and alkenes
110. How can you distinguish between hexane and hexene both of which are colorless liquids are
room temperature? Do not write equations for any expected reactions.
Top Questions
1. A student is asked to identify two salts T and W.
a) The student noted the following observations while testing salt T.
Observation 1: solution of T forms a white precipitate with aqueous sodium hydroxide and
the precipitate does not dissolve upon adding excess of sodium hydroxide.
Observation 2: solution of T forms a yellow precipitate with acidified aqueous silver nitrate.
i. What is the cation of salt T?
ii. What is the anion of salt T?
iii. Deduce the formula of salt T.
b) The student noted the following observations while testing salt W.
Observation 1: salt W gives a red color in a flame test
Observation 2: solution of W upon addition of dilute hydrochloric acid and gentle heating,
releases a gas that turns blue wet litmus red.
i. What is the cation of salt W?
ii. What is the anion of salt W?
iii. Deduce the formula of salt W.
iv. Describe how a flame test is done on a sample of salt W.
66
2. A student was analyzing two solids A and B.
Solid A is zinc iodide but solid B is unknown.
a) Tests on solid A.
Solid A was dissolved in water and the resulting solution was divided into three test tubes in equal
portions.
Complete the expected observations:
Portion
First portion
Test performed
A few drops of aqueous
sodium hydroxide were added
Observations
An excess of aqueous sodium
hydroxide was then added to
this mixture
Second portion
A few drops of aqueous
ammonia were added
An excess of aqueous
ammonia was then added to
this mixture
Third portion
dilute nitric acid and aqueous
silver nitrate were added to the
solution.
b) Tests on solid B
After dissolving solid B in water, the following observations were recorded while testing the
solution B.
Test on solution of B
dilute solution of sodium
hydroxide was added
Aluminum foil and aqueous
sodium hydroxide were
added, then the mixture was
heated and the gas produced
was tested with litmus paper
Observations
Pale blue precipitate formed
Bubbles released that turns wet litmus blue
Write the chemical formula for solid B
67
3. A student analyzed soluble carbonate salt R.
a) No observations were recorded when the carbonate salt R was slightly heated in a test tube.
What does this test indicate about salt R?
b) Salt R was dissolved in water and divided equally into four test tubes.
The following tests were conducted on the test tubes. Complete the observations for each of
the following tests.
test 1: a small amount of dilute of hydrochloric acid was added to the first test tube. The
mixture effervesces producing a gas.
i. test 2: The gas released was bubbled in lime water solution.
Observation:
ii. Identify the gas released.
c) Two more tests were conducted and the following observations were recorded.
tests on solution of salt R
observations
test 3:
To a new test-tube of solution, dilute sodium
hydroxide was added.
pungent gas given off
gas turned damp litmus paper blue
The mixture was then warmed carefully.
i. Identify the ion present in the solution from the observation of test 3?
ii. Why did the damp litmus turn blue?
iii. Identify the gas released.
d) Write the chemical formula for salt R.
e) Why it’s necessary to warm carefully the mixture in test 3? Explain your answer.