1
Chemistry Final Exam Review
Introduction to Chemistry and the Metric System
Good things to know:
• scientific method - hypothesis, theory, data
• graphing - independent variable, dependent variable
• pure science vs. applied science
• chemistry, hazard symbols, common laboratory equipment
• accuracy vs. precision
• SI system of measurement
Problems:
1. How many significant figures in:
3 a. 345
3 b. 0.00540 1 c. 300,000
2. Convert the following to scientific notation:
a. 3,000,020,000
b. 0.000002546
3.00002x109
d. 45.00010
2.546 x 10-6
3. Convert the following to decimal notation:
a. 6.020 x 1010
b. 3.5582 x 10-9
60,2θ0,000,000
4. 3.55 x 10 = 35.5 40
7
0.0000000035582
6.
312
= 0.005
6 x 10 4
7.
456.0
= 228.0
2.0000
5. 2.0450 x 0.0044 = 0.008998 0.0090
8.
(40.5)(5.45 x 10 -5 ) = 1.08 x10 −15
(2.04 x 10 )
12
10. 35 + 40.02 + 450 = ?
nearest 10
530
9.
(3.045)(6.032 x 10 23 )
(1.2 x 10 )(7.00 x 10 )
5
= 2.2 x10 8
11. 60.320 - 20.4444 = ?
nearest 0.001
39.876
12. 3.44 L = ? mL
13. 430 cm = ? km
3
3.44 L ∗
10 mL
= 3440mL
1L
14. 3.50 x 104 feet = ? km
3.50 x10 4 ft ∗
10
430cm ∗
1km
= 0.00430km
10 5 cm
15. 40.5 cm3 = ? dm3
12in 2.54cm 1km
∗
∗ 5
= 10.7 km
1 ft
1in
10 cm
40.5cm 3 ∗
1dm 3
= 0.0405dm 3
103 cm 3
2
Matter & Energy
Good things to know:
• matter, conservation of matter, conservation of energy
• substances, elements, compounds, heterogeneous mixtures, homogeneous mixtures - solutions
• physical and chemical change, physical and chemical properties
• abundance of elements in the earth’s crust
• energy, work, forms of energy (mechanical, electrical, chemical, light, etc.)
• Temperature scales: Celsius, Kelvin, absolute zero
• calorimeter, specific heat, conductor, insulator
• potential energy, kinetic energy
• Kinetic Theory of Heat, temperature = average Kinetic Energy
• Kinetic Theory: solid, liquid gas
• electrostatic force, electric current, electrically neutral, electrical conductor
Problems:
1. What is the density of this rectangular solid with a mass of 234 g?
V = (4.8cm)(7.2cm)(10.5cm) = 362.88cm3 → (2sf )
D=
m
234 g
=
= 0.64 g cm3
3
V 362.88cm
4.8 cm
10.5 cm
7.2 cm
2. What is the density of a cylinder with a mass of 98.45g if you insert it in 25.0mL of water and it raises the
level of the water to 35.4 mL?
V = 35.4mL − 25.0mL = 10.4mL → D =
m 98.45 g
=
= 9.47 g mL
V 10.4mL
3. What mass of iron would displace 34.0 cm3 of water?
34.0cm 3 ∗
7.87 g
= 268 g
1cm 3
4. What is the volume of a 62.9g piece of magnesium?
62.9 g ∗
1mL
= 45.6mL
1.38 g
5. Using the following data regarding the density of 3 liquids:
Liquid Volume(mL) Mass(g) Density(g/mL)
A
48.5
37.2
0.767
B
12.8
174.1
13.6
C
64.7
71.2
1.10
Calculate the density of each liquid. Then assume they were placed in a cylinder, and show what order,
top to bottom, they would settle into.
Top Bottom: A, C, B
3
6. Perform the following conversions:
a. 41.2 kJ = ? J
b. 816 J = ? cal
3
10 J
41.2kJ ∗
= 41,200 J
1kJ
c. 0.0482 kJ = ? cal
1cal
816 J ∗
= 195cal
4.18 J
7. Perform the following conversions:
a. 71oF = ? oC
5(71 − 32)
= 22 o C
9
103 J 1cal
0.0482kJ ∗
∗
= 11.5cal
1kJ 4.18 J
b. 132oC = ? K
c. 255 K = ? oC
132 + 273 = 405 K
255 − 273 = −18o C
8. How much heat is required to raise the temperature of 23.5 g of copper from 250K to
298K?
∆H=(23.5g)(0.385)(48oC)=430J
9. What is the specific heat (in J/goC) of an unknown substance if adding 0.124 kJ of heat to 10θg of the substance
causes the temperature to increase by 12oC?
124 J = (100 g )c(12 o C ) → c =
124
= 0.10 J g •oC
(100)(12)
10. How much heat is absorbed by 100.0 mL of water if its temperature increased by 13.2oC?
∆H=(100.0g)(4.18)(13.2oC)=5520J
11. If 15.00g of water at an initial temperature of 19oC absorbs 2.03 kJ of heat, what will its final temperature
be?
2030 J = (15.00 g )(4.18 J g •o C )∆T → ∆T =
2030
= 32.4 o C
(15.00)(4.18)
19 + 32.4 = 51o C
12. If you place 50.00g of a metal at a temperature of 98.0oC into 100.0mL of water at 22.3oC, and stir the mixture,
what is the specific heat of the metal if the final temperature of the mixture is 29.7oC? (answer in J/goC)
∆H water = (100.0 g )(4.18)(29.7 − 22.3) = 3093.2 J = − ∆H metal
− 3093.2 J = (50.00 g )c(29.7 − 98.0 o C )
− 3093.2 J
c=
= 0.91 J g •o C
(50.00 g )(−68.3o C )
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Atomic Theory
Good things to know:
• Dalton’ Atomic Theory, Law of Constant Composition, Law of Definite Proportions
• Cathode Ray Tube, Thomson, electron, “plum pudding” model
• Rutherford, Gold Foil Experiment, nucleus
• Bohr, ground state, excited state, bright line spectrum, flame tests
• Chadwick, neutron
• Heisenberg’s Uncertainty Principle, Charge Cloud Model, orbital (shapes of orbitals)
• atomic number, mass number, atomic mass, isotope, ion, charge
• proton, neutron, electron – know their relative mass and charge
Problems:
1. How many protons, neutrons, and electrons are there in:
238
22
a. 92 U
92p, 146n, 92e
+1
b. 11 Na
11p, 11n, 10e
32
−2
c. 16 S
16p, 16n, 18e
d. silver-108
47p, 61n, 47e
e. Argon-40
18p, 22n, 18e
2. Give the orbital diagram for:
1s
2s
2p
3s
a. carbon
b. magnesium
3. Give the electron configuration for:
a. aluminum
2
b. rhodium
1
c. tungsten
2
[Ne] 3s 3p
[Kr] 5s 4d
d. iodine
7
e. neon
2
10
5
[Kr]5s 4d 5p
2
2
14
f. dysprosium
2
1s 2s 2p
6
4. What is the atomic mass of an element with the following isotopes:
Isotope
Abundance
Magnesium-24
Magnesium-25
Magnesium-26
78.70%
10.13%
11.17%
(0.7870)(24) + (0.1013)(25) + (0.1117)(26) = 24.3247
4
[Xe] 6s 4f 5d
2
10
[Xe] 6s 4f
5
The Periodic Table
Good things to know:
• Pre-periodic table patterns: Newland’s Law of Octaves, Dobereiner’s Triads
• Mendeleev’s Periodic table (based on atomic mass), holes in his table
• Moseley, atomic number, protons, Modern Periodic Law
• basic characteristics and names of the major groups
• metals, nonmetals, metalloids – “staircase”
• ionization energy, electronegativity, atomic radius, trends shown in these properties on the periodic table
Problems:
1. Give the number of valence electrons, physical state (metal, nonmetal, or metalloid), and family name for the
following:
a. aluminum
b. rhodium
c. tungsten
3, M, boron group
X, M, metal
X, M, transition metal
d. iodine
e. neon
f. dysprosium
7, NM, halogens
8, NM, noble gases
X, M, rare earth
Nomenclature
Good things to know:
• ionic compound, molecular compound, acid
Problems:
1. Give the name for the following compounds:
a. NaOH
b. CO2
c. AgCl
sodium hydroxide
carbon dioxide
silver chloride
d. HCl
e. (NH4)2SO4
f. HgO
hydrochloric acid
ammonium sulfate
mercury(II) oxide
g. FeCl3
h. H2SO4
i. N2O5
iron(III) chloride
sulfuric acid
dinitrogen pentoxide
2. Give the formulas for the following compounds:
a. lithium oxide
Li2O
d. iron(III) oxide
Fe2O3
g. carbonic acid
H2CO3
j. sodium phosphate
Na3PO4
b. hydrosulfuric acid
H2 S
e. lead (II) sulfide
PbS
h. potassium oxalate
K2C2O4
c. calcium hydroxide
Ca(OH)2
f. copper(II) sulfate
CuSO4
i. nitrogen monoxide
NO
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Chemical Equations
Good things to know:
• 5 types of reactions: synthesis, decomposition, single replacement, double replacement, combustion
• monatomic and diatomic elements
• reactant, product, precipitate, know your solubility rules
• net ionic equation, spectator ion
Problems:
1. Name the reaction type and balance:
combustion
2
a.
15
C6H6(l) +
O2(g) 12
CO2(g) +
6
H2O(l)
double replacement b. 2 K3PO4(aq) + 3 Ba(NO3)2(aq) 6 KNO3(aq) + 2 Ba3(PO4)2(s)
decomposition
_____XXX_____ d.
c.
4
2
2
KClO3(s) FeS2(s) +
11
O2(g) KCl(s) +
2
3
O2(g)
Fe2O3(s) +
8
SO2(g)
2. Complete the following reactions by giving the correct formulas for the products (there is no need to balance):
a. P2O5 P + O2
b. Al + O2 Al2O3
c. AgNO3(aq) + NaOH(aq) AgOH(s) + NaNO3(aq)
d. LiNO3(aq) + Na3PO4(aq) NR
e. Ag + NaCl f. Al + CuSO4 NR
Al2(SO4)3 + Cu
g. Li2CO3(aq) + Pb(C2H3O2)2(aq) h.
C5H12(l) + O2 PbCO3(s) + LiC2H3O2(aq)
CO2 + H2O
3) Give the net ionic equation for the following reactions from #2:
c)
Ag+ + OH- AgOH
f)
Al + Cu+2 Al+3 + Cu
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Stoichiometry
Good things to know:
• standard temperature and pressure (STP)
• Avogadro’s Number, mole
• atom, formula unit, molecule
Problems:
1. What is the formula mass of :
a. CH4
b. C6H12O6
2 + 4 = 16
c. Mg(ClO3)2
6(12)+12+6(16)=180
24.3+2(35.5)+6(16)=191
2. What is the mass of 7.03 moles of silver carbonate?
7.03mol Ag 2 CO 3 ∗
276 g Ag 2CO 3
= 1940 g Ag 2 CO 3
1mol Ag 2 CO 3
3. How many liters of space are occupied by 2.33 x 1022 molecules of methane (CH4) gas at STP?
2.33 x10 22 mcs CH 4 ∗
1mol CH 4
22.4 L CH 4
∗
= 0.867 L CH 4
23
6.02 x10 mcs CH 4 1mol CH 4
4. How many atoms are there in 10.3g of silver?
1mol Ag 6.02 x10 23 at Ag
10.3 g Ag ∗
∗
= 5.74 x10 22 at Ag
108 g Ag
1mol Ag
5. What is the percent composition of :
a. CO2 12 + 2(16) = 44
12
x100 = 27.3%
44
32
O:
x100 = 72.7%
44
C:
b. CuSO4 63.5 + 32.1 + 4(16) = 159.6
63.5
x100 = 39.7%
159.6
32.1
S:
x100 = 20.1%
159.6
64
O:
x100 = 40.0%
159.6
Cu :
6. If you have 20.0g of a compound which is composed of 16.6g of potassium and 3.40g of oxygen,
what is the empirical formula?
1mol K
= 0.425 K
2K
39.1g K
÷ .2125
→ ef : K 2 O
1mol O
1O
3.40 g O ∗
= 0.2125 O
16 g O
16.6 g K ∗
7. If you have a compound which is composed of 43.7% phosphorus and 56.3% oxygen, and has a formula
mass of 284 g/mol, what is the empirical formula? What is the molecular formula?
1mol P
= 1.41 P
1P
31g P
284
÷ 1.41
× 2 → ef : P2 O 5 →
= 2 → mf : P4O10
1mol O
2.5 O
142
56.3 g O ∗
= 3.52 O
16 g O
43.7 g P ∗
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8. Use the following equation:
C6H12O6(s) + 6 O2(g) 6 CO2(g) + 6 H2O(l)
a. How many moles of carbon dioxide will you get if you begin with 2.50 mol of glucose?
2.50mol C 6 H12O 6 ∗
6mol CO 2
= 15.0mol CO 2
1mol C 6 H12O 6
b. How many liters of oxygen gas are needed to produce 0.430 mol of water at STP?
0.430mol H 2 O ∗
6mol O 2 22.4 L O 2
∗
= 9.63L O 2
6mol H 2O 1mol O 2
c. How many grams of water would you expect if you began with 15.0g of glucose?
15.0 g C 6 H12 O 6 ∗
1mol C 6 H12O 6
6mol H 2O
18 g H 2 O
∗
∗
= 9.00 g H 2O
180 g C 6 H12 O 6 1mol C 6 H12 O 6 1mol H 2O
d. How many liters of carbon dioxide would you expect if you began with 15.0g of glucose at STP?
15.0 g C 6 H12 O 6 ∗
1mol C 6 H12 O 6
6mol CO 2
22.4 L CO 2
∗
∗
= 11.2 L CO 2
180 g C 6 H12 O 6 1mol C 6 H12 O 6 1mol CO 2
9. 3 Fe(s) + 4 H2O(l) Fe3O4(s) + 4 H2(g)
If you begin with 30.0g of iron and 25.0g of water:
a. What is the limiting reactant? What is the excess reactant?
1mol Fe 1mol Fe3O 4
∗
= 0.179mol Fe 3O 4 → LR : Fe
55.8 g Fe
3mol Fe
1mol H 2 O 1mol Fe3O 4
25.0 g H 2O ∗
∗
= 0.347 mol Fe3O 4
18 g H 2O 4mol H 2 O
30.0 g Fe ∗
b. How many grams of iron oxide and liters of hydrogen gas do you expect?
0.179mol Fe3O 4 ∗
231g Fe 3O 4
= 41.4 g Fe3O 4
1mol Fe3O 4
0.179mol Fe3O 4 ∗
4mol H 2
22.4 L H 2
∗
= 16.0 L H 2
1mol Fe3O 4 1mol H 2
c. How many grams of the excess reactant will be left over?
0.179mol Fe3O 4 ∗
4mol H 2 O 18 g H 2O
∗
= 12.9 g H 2O used
1mol Fe3O 4 1mol H 2O
25.0 - 12.9 = 12.1g H 2O left
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Bonding
Good things to know:
• Lewis Structures for atoms, ions, and molecular (covalent) compounds
• shared pair of electrons, unshared pair, single bond, double bond, triple bond
• VSEPR Theory, hybrid orbitals, shapes of molecules, sigma bonds, pi bonds, polarity
• Intermolecular Forces (in order from weakest to strongest): London Dispersion Forces, dipole-dipole
interactions, H-bonding, ionic bonding, covalent networks what substances fit theses groups
• metallic bonding: why are metals solids, but good conductors of heat and electricity?
• alkane, alkene, alkyne, saturated, unsaturated hydrocarbons
• IUPAC nomenclature (know your prefixes), polymer
Problems:
1. For each of the following, give the Lewis Structure, shape, hybrid orbital type, sigma & pi bonds, and polarity:
e. NO3-1 trigonal planar, sp2, 3σ, 1π, P
a. CO2 linear, sp, 2σ, 2π, NP
-
O
O C
N
O
O
O
b. NH3 pyramid, sp3, 3σ, P
f. NH4+ tetrahedron, sp3, 4σ, P
+
H
H
H
N
H
H
N
H
H
c. CCl4 tetrahedron, sp3, 4σ, NP
g. F- ION
Cl
Cl
C
-
Cl
F
Cl
d. CH2O trigonal planar, sp2, 3σ, 1π, P
x
h. MgCl2 IONIC
O
C
-
H
H
-
+2
Mg
Cl
Cl
2. Give the Lewis structure and sigma & pi bonds for the following:
a. CH3CH2CH3 10σ
c. CH3COCH3 9σ, 1π
O
C
C
C
C
b. SeF4 4σ
C
C
d. C5H12 (3 isomers) 16σ
F
F
Se
F
C
C
F
C
C
C
C
C
C
C
C
C
C
C
C
C
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3. Using bond energies , estimate the total energy change (∆H) in the following reactions:
a. C2H2 + 2 H2 C2H6
(C=C) + 2 (H-H) – (C-C)- 4 (C-H) 2 C-H bonds don’t change
(837) + 2(436) – (347) – 4(414) = -294 kJ
b. Cl2 + C2H4 C2H4Cl2
(Cl-Cl) + (C=C) – (C-C) – 2(C-Cl) the 4 C-H bonds don’t change
(243) + (611) – (347) – 2(330) = -153 kJ
4. Give the structure for the following compounds:
a. 2-methyl butane
b. propene
c. 2,4 – dimethyl pentane
CH3
CH3
CH
CH3
CH2
CH3
CH2
CH3
CH
CH3
CH
CH3
CH2 CH
CH3
Phases of Matter
Good things to know;
• solid, liquid, gas - Kinetic Theory of Phases; fluid
• pressure - barometer, manometer, standard pressure
• boiling, melting, freezing, condensation, evaporation, sublimation, deposition
• vaporization - evaporation vs. boiling
• hydrate, anhydrous, dessicant, deliquescent substance
• crystal lattice, viscosity
o
• density of solids, liquids, and gases; density of water at 4 C
• heats of: vaporization, fusion, crystallization, condensation
• phase diagrams - triple point, critical temperature, unique properties of water
Problems:
1. Solve the following pressure unit conversions:
a. 355 kPa = ? atm
355kPa ∗
1atm
= 3.50atm
101.3kPa
b. 1022 Torr = ? kPa
1022Torr ∗
101.3kPa
= 136kPa
760Torr
2. What is the pressure of a 712 N object resting on a rectangular area 0.55 m x 0.025 m?
A = (0.55m)(0.025m) = 0.01365m 2 (2 sf )
F
712 N
P= =
= 52,000 Pa
A 0.01365m 2
3. What is the pressure exerted by the gas?
a.
P = 743 mm Hg
45 mm
743 + 45 = 788mmHg
b.
P = 103.2 kPa
203 mm
774+203=977 mmHg=13θ KPa
11
4. How much heat is required to melt 50.0g of ice at its freezing point?
∆H= qm =(333J/g)(50.0g) = 16,700J
5. How much heat is lost when 200.0g of steam condenses at its boiling point?
∆H = qm = (-2260J/g)(200.0g) = -452,000J
o
6. How much heat is required to change 50.0g of ice at -20.0 C to steam at 150oC?
∆H=mc∆T = (50.0g)(2.06J/g-oC)(20.0oC)
∆H = qm = (50.0g)(333 J/g)
∆H=mc∆T = (50.0g)(4.18J/g-oC)(100.0oC)
∆H = qm = (50.0g)(2260 J/g)
∆H=mc∆T = (50.0g)(2.02J/g-oC)(50.0oC)
=
2060J
= 16,650J
= 20,900J
= 113,000J
= 5050J
----------------
157,700J
The Gas Laws
Good things to know:
• Kinetic Theory of Gases
• directly proportional, inversely proportional
• 5 assumptions for Ideal Gases - conditions in which ideal behavior is not exhibited
• Standard Temperature and Pressure (STP)
Problems:
1. If you have 35.0 mL of a gas in a closed container under 1.23 atm of pressure, what will the volume be if the
pressure is increased to 3.00 atm with no change in temperature?
P1V1 = P2V2 → (1.23atm)(35.0mL) = (3.00atm)V2 → V2 = 14.4mL
2. If you have a 1.50L container at room temperature, what will the volume be if you place it in some boiling
water with no change in pressure?
V1 V2
1.50 L
V
=
→
= 2 → V2 = 1.88 L
T1 T2
298 K 373K
3. If you have a 2.75dm3 balloon under 455kPa of pressure at 100oC, what will the volume be at STP?
P1V1 P2V2
(455kPa )(2.75 L) (101.3kPa)V2
=
→
=
→ V2 = 9.04 L
T1
T2
373K
273K
4. If you have a container with a mixture of nitrogen and oxygen gases showing 1023 Torr of pressure, and
the pressure of the nitrogen alone is 544 Torr, what is the pressure of oxygen gas?
Ptot = PN + PO → 1023Torr = 544Torr + PO → PO = 1023 − 544 = 479Torr
5. If I have 2.43 mol of gas at 35.0oC and a pressure of 789 Torr, what volume will it occupy?
atm • L
)
PV = nRT → (1.04atm)V = (2.43mol )(0.0821 mol
• K (308 K ) → V = 59.1L
12
6. What is the relative rate of helium (He) gas vs. chlorine (Cl2) gas? Which one diffuses faster?
rr =
FM Cl2
FM He
=
71
= 4.21 → He diffuses 4.21 times faster than Cl 2
4
7. What volume (in L) of oxygen gas would you collect over water at 755 mmHg of atmospheric pressure and a
room temperature of 22oC if you decomposed 5.00g of potassium chlorate via the following reaction:
*NOTE: The vapor pressure of water at 22oC is 2θ mmHg
2 KClO3(s) 2 KCl(s) + 3 O2(g)
5.00 g KClO3 ∗
1mol KClO3
3mol O 2
∗
= 0.0612mol O 2
122.5 g KClO3 2mol KClO3
755 − 20 = 735Torr = 0.967 atm
atm • L
)
PV = nRT → (0.967 atm)V = (0.0612mol )(0.0821 mol
• K ( 295 K ) → V = 1.53L
Solutions
Good things to know:
• mixtures (how are they different from substances?), solution, solvent, solute, 5 properties of solutions
• saturated, unsaturated, supersaturated, concentrated, dilute solutions
• examples of: gas, solid in liquid, liquid in liquid, gas in liquid, and solid solutions
• Kinetic Theory of solutions (“pirhana and cow theory”)
• factors affecting solution rate: temperature, surface area, stirring
• factors affecting solubility: temperature, nature of solute and solvent, pressure(gases)
• tincture, alloy, aqueous solution, colloid, suspension, emulsion
• “like dissolves like”, solubility curve
• colligative properties - antifreeze
Problems:
1. What is the molarity of a solution in which you dissolve 5.22 mol of NaCl into a 3.00L solution?
M=
mol solute 5.22mol
=
= 1.74 M
L solution
3.00 L
2. What is the molarity of a solution in which you dissolve 40.0g of NaCl into a 0.500L solution?
40.0 g NaCl ∗
1mol NaCl 0.684mol NaCl
=
= 1.37 M
0.500L
58.5 g NaCl
3. How many grams of silver sulfate (Ag2SO4) are needed to make 2.00L of a 0.0220M solution?
M=
mol
312 g Ag 2SO 4
→ mol = MV = (0.0220 M )(2.00 L) = 0.0440mol Ag 2SO 4 ∗
= 13.73 g Ag 2SO 4
L
1mol Ag 2SO 4
4. If you dilute 50.0 mL of a 6.00M HCl solution to 400.0 mL, what will the new concentration be?
M 1V1 = M 2V2 → (6.00M )(50.0mL) = M (400.0mL) → M = 0.750M
13
5. What is the concentration (in ppm) of a solution in which you dissolve 1.00 mg of NaCl in 300 mL of water?
30θmL = 0.300 L → ppm =
1.00mg NaCl
= 3.33 ppm
0.300 L
6. What is the percent by mass of a solution in which 30.0g of NaCl is dissolved in 250.0 mL of water?
msolution = 30.0 + 250.0 = 280.0 g
%=
g solute
30.0 g
=
x100% = 10.7%
g solution 280.0 g
7. Use the solubility chart below to answer the following:
a. How many grams of KNO3
can be dissolved in 100g of water at
20oC and 60oC?
20: 33g KNO3/100g H2O
60: 112g KNO3/100g H2O
b. Which of the following is
most soluble at 20oC: KNO3, NaCl,
KClO3, or NH3? Which is most soluble
at 50oC?
20: NH3 50: KNO3
c. How many grams of
potassium nitrate can be dissolved in
150.0g of water at 50oC?
87 g KNO 3
x
=
100 g H 2 O 150 g H 2 O
x=
(87)(150)
= 131g KNO 3
100
d. If you put 50.0g of potassium
nitrate in 100g of water at 40oC, is the
solution saturated, unsaturated,
or supersaturated? How much is left
over?
40: 67g KNO3 maxunsaturated
e. What is the molarity of a
saturated NaNO3 solution at 50oC?
1mol NaNO 3
= 1.32mol NaNO 3 1.32mol
85 g NaNO 3
= 13.2 M
0.100 L
100 g H 2 O = 100mL H 2 O ≈ 0.100 L H 2 O
112 g NaNO 3 ∗
14
Chemical Kinetics & Thermodynamics
Good things to know:
• collision theory - collision effectiveness, collision frequency
• 4 factors that affect reaction rate
• intermediate, activated complex, activation energy, heat of reaction (∆H) - exothermic vs. endothermic
• reaction mechanism, overall reaction, rate-determining step
• rate laws, k, general rate law
• enthalpy (∆H) - exothermic vs. endothermic
• entropy (∆S) - disordered vs. ordered
• Gibbs’ Free Energy (∆G) - spontaneous vs. nonspontaneous
Problems:
1. For the following reaction mechanism:
H2O2 + I-1
HOI + OH-1
-1
+1
OH + H
H 2O
HOI + H+1 + I-1 I2 + H2O
(slow)
(fast)
(fast)
H2O2 + 2 H+ + 2 I-1 2 H2O + I2
b. What is the rate-determining step? step 1 (slow step)
-1
c. Name any intermediates. Are there any catalysts? HOI, OH ; no catalysts
a. What is the overall reaction?
d. What is the effect of :
No effect (H+ in fast step)
rate decreases (I-1 in slow step)
i. increasing the concentration of hydrogen ions?
ii. decreasing the iodide concentration?
e. Write the rate law for this reaction. What order is it?
rate = k(H2O2)(I-1) 2nd order overall
A+B C+D
2. Observe the following data from a rate study for the reaction:
Trial
1
2
3
4
5
6
7
8
Initial [A]
0.166
0.332
0.498
0.664
0.166
0.332
0.166
0.332
Initial [B]
0.100
0.100
0.100
0.100
0.200
0.200
0.400
0.400
Initial Rate (M/s)
4.28 x 10-6
8.56 x 10-6
1.28 x 10-5
1.71 x 10-5
3.42 x 10-5
6.85 x 10-5
2.74 x 10-4
5.48 x 10-4
a. Find the rate law for this reaction. What is the overall order of this reaction? in A? in B?
general rate law: rate = k(A)x(B)y
x
 (0.332) 
8.56 x10 −6
 (0.166)  = 4.28 x10 −6


y
 (0.200) 
3.42 x10 −5
=
 (0.100) 
4.28 x10 −6


 ( A2 )
Rate2

 =
Rate1
 ( A1 ) 
 (B2 ) 
Rate2
 (B )  = Rate
 1 
1
* This reaction is
x
2x = 2
x =1
y
4
th
order overall,
1
st
in A and
2y = 8
y=3
3
rd
in B
b. Find k for this reaction.
4.28 x 10-6 M/s = k(0.166M)(0.100M)3
k = 2.58 x 10-2 M-3s-1
15
c. If [A] = 0.500 M, and [B] = 0.500M, what would the initial rate be for the reaction?
rate=(0.0258)(0.500)(0.500)3 = 1.61x10-3 M/s
d. What concentration of B would you need if the initial rate were 3.06 x 10-3 M/s, and [A] = 0.650 M?
3.06 x10 −3 = (0.0258)(0.650) x 3 → x 3 = 0.182 → x = 3 0.182 = 0.567 M
3. Energy Diagram. Use the diagram shown below: (6 pts)
#2
70
#1
c
Energy (kJ)
60
50
a
40
b
d
30
#3
20
e
10
course of reaction
B
Exothermic
C
#3
It would decrease
Activated complex
a. What letter represents the enthalpy or heat of this reaction?
b. Is the reaction endothermic or exothermic?
c. What letter indicates the activation energy for the forward reaction?
d. Which number indicates the products?
e. What would happen to the value of “c” when a catalyst is added?
f. What is the term for the “stuff” described at point #2?
4. For each of the following find ∆Ho, ∆So, ∆Go and state if the reaction is exothermic or endothermic;
favors more or less disorder; and is spontaneous or nonspontaneous:
* Assume that all of these reactions occur at room temperature (298K)
a. 4 NH3(g) + 7 O2(g) 4 NO2(g) + 6 H2O(g)
∆H = 4(NO2) + 6(H2O(g)) - 4(NH3) - 7(O2)
∆H = 4(+33.1) + 6(-242) – 4(-46) – 7(0) = -1136kJ, exothermic
∆S = 4(240.5) + 6(188.7) – 4(192.5) – 7(205) = -111J/K, less disorder
∆G = -1136 kJ – (298K)(-0.111 kJ/K) = -1103 kJ, spontaneous
b. 4 NO2(g) + O2(g) 2 N2O5(g)
∆H = 2(N2O5) - 4(NO2) - (O2)
∆H = 2(+11.3) + 4(33.1) - (0) = -110kJ, exothermic
∆S = 2(356) + 4(240.5) – (205) = -455J/K, less disorder
∆G = -110 kJ – (298K)(-0.455 kJ/K) = +26 kJ, nonspontaneous
16
Equilibrium
Good things to know:
• characteristics of chemical equilibrium
• equilibrium constants (Keq, Ksp, Ka), mass-action expression
• dissociation equation, solubility product expression, molar solubility
• Le Chatelier’s Principle
Problems:
1. For the following equilibrium: N2(g) + 3 H2(g)
a. give the mass-action expression
2 NH3(g) + E ∆H = -92.6 kJ 4 mol gas 2 mol gas
2
NH 3 ]
[
K eq =
[ H2 ]3 [ N 2 ]
b. If you have 1.50 mol of nitrogen, 2.00 mol of hydrogen and 0.230 mol of ammonia gases in a 3.50 L
container, what are the concentrations of the gases?
(N 2 ) =
1.50mol
2.00mol
0.230mol
= 0.429 M , (H 2 ) =
= 0.571M , (NH 3 ) =
= 0.0657 M
3.50 L
3.50 L
3.50 L
c. Using your results from (b), what is the equilibrium constant? Does this favor the products or the
reactants?
2
NH 3 ]
[
(0.0657) 2
K eq =
=
= 0.0540, favors reactants
[H 2 ]3 [N 2 ] (0.571) 3 (0.429)
d. If, at equilibrium, the constant is 0.0243, and the concentrations of nitrogen and hydrogen are
1.02M and 2.31M, respectively, what is the concentration of ammonia?
2
NH 3 ]
[
x2
K eq =
→
0
.
0243
=
→ x 2 = 0.3055 → x =
3
3
(2.31) (1.02)
[H 2 ] [N 2 ]
0.3055 = 0.553M
e. What happens to [H2] and [NH3] if you increase the concentration of nitrogen?
(H2) decreases, (NH3) increases
f. What happens to [H2] and [N2] if you decrease the concentration of ammonia (NH3)?
(H2) and (N2) decrease
g. What happens to the concentrations of all the species if you raise the temperature?
(H2) and (N2) increase, (NH3) decreases
h. What happens to the concentrations of all the species if you increase the pressure?
as P increases, the container gets smaller, so the side with less moles of gas is favored
(H2) and (N2) decrease, (NH3) increases
i. What happens to the concentrations of all the species if a platinum catalyst is added?
No effect
17
Acids and Bases
Good things to know:
• Arrhenius definitions of acid and base
• Lowry-Bronsted definitions of acid and base
• 5 properties of acids and bases
• salts, electrolyte, nonelectrolyte
• ionization vs. dissociation
• self-ionization of water, pH, pOH, what is acidic and basic on each scale
• buffers, common-ion effect
• titration
Problems:
1. Identify each of the following as Arrhenius (A)cid, Arrhenius (B)ase, (S)alt, or (N)either:
a. NaCl S
b. H2SO4 A
c. KOH B
d. NH4Cl S
e. Ba(OH)2
f. K3PO4 S
g. CH4 N
h. Al2(SO4)3 S i. HBr A
j. CO2 N
2. Complete and balance the following dissociation equations:
+
b. Ba(OH)2(s) a. NaOH(s) Na + OH
3. Complete and balance the following ionization equations:
a. HCl(aq) H+ + Cl-
b. H3PO4(l) B
Ba+2 + 2 OH-
H+ + H2PO4-
4. Complete and balance the following neutralization equations:
a. HCl(aq) + NaOH(aq) NaCl + HOH
b. H2SO4(aq) + 2 KOH(aq) K2SO4 + 2 HOH
+
5. What is the concentration of hydrogen ions [H ] in a:
a. solution with a pH of 3.54
b. 0.500M HF solution
3.54 = − log( H + )
x2
= 6.3 x10 − 4
0.500
x 2 = 3.15 x10 − 4
( H + ) = 10 −3.54 = 2.88 x10 − 4 M
x = 3.15 x10 − 4 = 0.018M
+
pH = − log( H )
c. solution in which concentration of hydroxide = 0.00120M
( H + )(OH − ) = 1.00 x10 −14 → ( H + ) =
1.00 x10 −14
= 8.3 x10 −12 M
0.00120
6. For each of the following Lowry-Bronsted reactions:
- name the acids and bases
- list the conjugate pairs
- name any amphoteric substances
a. HCl(aq) + H2O(l)
A
B
b. NH3(aq) + H2O(l)
B
A
c. H2SO4(aq) + CO3-2(aq)
A
Cl-1(aq) + H3O+1(aq)
B
A
NH4+1(aq) + OH-1(aq)
A
B
HSO4-1(aq) + HCO3-1(aq)
HCl, Cl- ; H2O, H3O+
NH3, NH4+; H2O, H3O+
B
B
A
H2SO4, HSO4-; CO3-2, HCO3H2O is amphoteric (base in (a), acid in (b))
18
7. What is the pH for a solution in which:
a. [HCl] = 2.00 M strong acid
pH = − log( H + ) =
− log(2.00) = −0.301
b. [HC2H3O2] = 2.00 M
x2
= 1.8 x10 −5 → x = 3.6 x10 −5 = 0.0060 M
2.00
pH = − log(0.0060) = 2.22
c. pOH = 2.05
pH + pOH = 14 → pH = 14 − 2.05 = 11.95
d. you dissolve 3.54L of HCl gas in 3.00 L of water
1mol HCl 0.158mol HCl
=
3.00 L
22.4 L HCl
+
= 0.0527 M = ( H ) → pH = − log(0.0527) = 1.28
3.54 L HCl ∗
e. you dissolve 0.0325g of NaOH in 0.500 L of water
1mol NaOH 8.125 x10 − 4 mol NaOH
=
= 0.001625M = (OH − )
0
.
500
L
40 g NaOH
pOH = − log(0.001625) = 2.79 → pH = 14 − 2.79 = 11.21
0.0325 g NaOH ∗
8. What is the pOH for a solution in which:
a. [OH-1] = 0.012M
pOH = − log(0.012) = 1.92
b. concentration of hydrogen = 0.00435M
pH = − log(0.00435) = 2.36
pOH = 14 − 2.36 = 11.64
c. pH = 7.93
pOH = 14 − 7.93 = 6.07
9. If you have 500 mL of a 6.00M HCl solution, how much water must be added to make it 1.50M HCl ?
M 1V1 = M 2V2 → (6.00 M )(500mL) = (1.50 M )V → V = 2000mL solution
2000 - 500 = 1500mL water
10. If you titrate 50.0mL of an unknown solution of HCl with 35.4mL of 0.600M NaOH, what is the concentration of the acid?
M aVa = M bVb → M (50.0mL) = (0.600 M )(35.4mL) → M = 0.425M
11. If you titrate 35 drops of vinegar to the equivalence point with 43 drops of 0.600M NaOH, what is the
concentration of the vinegar?
M aVa = M bVb → M (35dr ) = (0.600 M )(43dr ) → M = 0.74 M
12. Rank the following acids from weakest to strongest: HCN, H2SO4, HC2H3O2, HNO2, HF, HCl
HCN, HC2H3O2, HNO2, HF, H2SO4 , HCl
19
Redox Reactions
Good things to know:
• what types of reactions are redox reactions
• oxidation, reduction, oxidizing agent, reducing agent
• oxidation number
• galvanic cell, cathode, anode, direction of electron flow
• types of batteries and how they work
Problems:
1. Give the oxidation number for all the elements in the following:
+2
-2
+4
+1 - 2
b. Na 2 S
0
a. O 2
+2
+6
+1
-8
+1 + 6- 2
e. K 2 S O 4
+5
+1
-4
-6
+6
+ 1+ 5 - 2
f. K Cl O 3
+3
+12 −14
−8
+6−2
h. Cr2 O 7 −2
+ 6− 2
g. Cr O 4 − 2
2. Balance the following reactions. Assume they are in acidic solution. Name the oxidizing agent and reducing
agent:
a)
+2 +2
+5 −2
0
4 H+ + Cu + 2 N O 3 −1 Cu
+4 −2
+ 2 N O2
+ 2 H2O
(lose 2)
2(gain 1)
OA- NO3
b)
-1
RA- Cu
+4 −2
+1 −2 −1
H2O + 2 N O 2 + Cl O
+5 −2
−1 -1
2 N O 3 −1 + C1 + 2 H+
2(lose 1)
(gain 2)
OA – ClO-1 RA – NO2
c)
+4 −2
+5 −2
+1
+6 -2
0
2 H+ + 14 H2O + 5 Te O 2 + 2 Br O 3 −1 5 H 6 Te O 6 + Br2
5(lose 2)
(gain 10)
OA – BrO3
-1
RA – TeO2
3. Determine the standard potential for the following reactions. Also give the overall reaction, and identify the
cathode and anode:
a. Ni – Ni+2 / Cu – Cu+2
cathode:
anode:
Cu+2 + 2 e- Cu
Ni Ni+2 + 2 e-
Cu+2 + Ni Ni+2 + Cu
+0.34
+0.26
+0.60V
-4
+1 + 3- 2
d. Na N O 2
+ 4- 2
c. C O 2
20
b. Ag – Ag+1 / Zn – Zn+2
2Ag+1 + 2 e- Ag
Zn Zn+2 + 2 e-
cathode:
anode:
2Ag+1 + Zn Zn+2 + 2Ag
+0.80
+0.76
+1.56V
c. Al – Al+3 / Mg – Mg+2
2Al+3 + 6 e- 2Al -1.66
3Mg 3Mg+2 + 6 e- +2.37
cathode:
anode:
2Al+3 + 3Mg 3Mg+2 + 2Al
+0.71V
Nuclear Chemistry
Good things to know:
• types of radiation: alpha particle, beta particle, neutron, proton, positron
• stability of nuclei; deviation from 1:1 ratio
• radioactive decay, nuclear disintegration, fission, fusion
• half-life
• radioactive dating, nuclear weapons, nuclear power plant
• nuclear energy - pros and cons
• rem, Curie, Bequerel
• dosimeter, Geiger Counter
Problems:
1) Complete the following nuclear equations:
218
a)
222
86
Rn 84
b)
209
84
Po 209
85
Po
+
4
2
0
At + −1
α
c)
207
82
β
d)
1
0
211
Pb + 42 α 84 Po
n +
235
92
U 142
56
91
Ba + 36
2. How much of a 3.0 g sample of 82Br will remain after 4.66 hours?
t
A = Ao ( 2 )
1
t1
2
= (3.0 g )( 1 2 )
4.66 hrs
35.5 hrs
= 2.7 g
3. How long will it take for a sample of 45Ca to decay to 10.0% of its original mass?
0.693
0.693
→k =
= 0.0042da −1
k
165da
 A
1
ln  = − kt → ln  = −(0.0042da −1 )t
 10 
 Ao 
t 12 =
− 2.30 = −(0.0042)t → t =
− 2.30
= 548da
− 0.0042
Kr
+3
1
0
n