Exam Review for Grade 12 Chemistry
Unit 1—Atomic Theory
Theory:
 Contributions made by each of the following scientists: Dalton, Rutherford, Chadwick,
Thomson, Bohr, Heisenberg
 Electron configuration
 Explain what each of the following quantum numbers describes: principal quantum number (n),
secondary quantum number (l), magnetic quantum number (ml), spin quantum number (ms)
 VESPR theory-know the shapes [linear, trigonal planar, tetrahedral, pyramidal, bent, trigonal
bipyramidal, octahedral, square planar
 Orbitals: s, p, d, f and hybridized orbitals sp, sp2, sp3
 Formation of single, double and triple bonds
Problems:
1. Draw Lewis Structure for each of the following and predict its shape.
SiF4, NI3, H2S, XeF4, XeF6, PO43-, PCl5
Unit 2—Organic
Theory:
 Importance characteristic of the carbon atom for forming organic compounds
 Structural Isomers
 Properties, Nomenclature and drawing of Hydrocarbons (Alkanes, Alkenes, Alkynes, Aromatic
Hydrocarbons)
 Reactions of Alkanes: Combustion, Substitution
 Reactions of Alkenes and Alkynes: Halogenation, Hydrogenation, Hydrohalogenation,
hydration (with H2O), Elimination Reaction
 Properties, Nomenclature and drawing of: Organic Halides, Alcohols, Ethers, Aldehydes,
Ketones, Carboxylic Acids, and Esters
 Reactions Involving Alcohols—Hydration, Combustion, Elimination
 Reactions involving Ethers—Condensation
 Reactions involving Aldehydes and Ketones: Oxidation, Hydrogenation, making Carboxylic
Acid
 Reactions involving Carboxylic Acid and Esters: Esterification, Hydrolysis
 Difference between addition polymerization and condensation polymerization (give examples)
Problems:
1. Draw the organic compounds
a. 2-methyl-3-pentanone
b. p-ethylmethylbenzene
d. Cyclohexene
e. Trimethylamine
g. 2-iodo-3-ethyl-4-methylpentanoic acid
c. 3,4,5-trinitro-2-pentanol
f. Ethoxypropane
h. Heptanal
2. Predict the products and draw the organic structures. Include any catalysts.
a. Butanoic acid + NaOH
b. Ammonia + 1-bromopropane
c. Methyl propanoate + water
d. 1-hexanol + oxidizing agent
e. 1-propanol + ethanoic acid
f. 2-hexanol + oxidizing agent
g. 1-butene + water
h. Benzene + bromine
Unit 3—Energy Changes and Rates of Reaction
Theory:
 Define the following: thermal energy, chemical system, surrounding, heat, exothermic,
endothermic, temperature, open system, isolated systems, closed system, calorimetry, specific
heat capacity, enthalpy, molar enthalpies, entropy, Gibb’s Free Energy
 Hess’s Law
 Rates of Reaction r = ∆ c / ∆ t
 Factors Affection Reaction Rates: Nature of Reactants, Concentration, Temperature and
Catalyst, Surface Area
 Rate Laws and Order of Reaction
 Collision Theory and Rate of Reaction
 ∆Hf is equal to Zero (WHEN???)
 ∆ H = Σ n∆ Hof (products) - Σ n∆ Hof (reactants)
 Reaction Mechanism
Problems:
1. The standard enthalpy of formation of sulfur dioxide is -296.8 kJ/mol.
a. Write a thermochemical equation for the formation reaction
b. Sketch a potential energy diagram for the reaction, labeling axes, enthalpy of reactants,
enthalpy of products and ∆H.
c. If 9.63 g of sulfur dioxide is formed under standard conditions, what quantity of heat is either
absorbed or released?
2. For the combustion of hydrogen:
H2 (g) + ½ O2 (g)  H2O (g) + 243 kJ.
How many grams of hydrogen would be necessary for the release of 972 kJ?
3. Calculate ∆H for the reaction:
4 CH3NO2 (g) + 3 O2 (g)  4 CO2 (g) + 2 N2 (g) + 6 H2O (g)
Given the following data:
i)
ii)
iii)
C (s) + O2 (g)  CO2 (g)
2 H2 (g) + O2 (g)  2H2O (g)
2 C(s) + 3 H2 (g) + 2 O2 (g) + N2 (g)  2 CH3NO2 (g)
∆H = -393.5 kJ
∆H = -483.6 kJ
∆H = -226.2 kJ
4. A 2.0 kg copper kettle (c = 0.385 J/g oC) contains 0.500 kg of water at 20.0 oC. How much heat is
needed to raise the temperature of the kettle and its contents to 80.0 oC?
5. If you have 60 kg of water heated from 15oC to 700C, how much energy is required?
6. Using standard enthalpies of formation calculate the enthalpy change for:
4 NH3 (g) + 5 O2 (g)  4 NO(g) + 6 H2O (g)
Unit 4—Chemical Systems and Equilibrium
Theory:
 Define: chemical reaction equilibrium, common ion effect, dynamic equilibrium, Le Chatelier’s
principle, reaction quotient, reversible reaction, solubility, solubility equilibrium, Ksp, Ka, Kb, Kw,
endpoint, equivalence point, Lewis acid, Lewis base, pH, pOH, weak acid, weak base,
titration, conjugate acid/base pair
 When is a system considered at equilibrium?
 How does Le Chatelier’s Principle undergo equilibrium shift when concentration is changed,
temperature is changed, gas volume is changed?
 When does Le Chatelier’s Principle NOT go through a equilibrium shift? [adding catalyst,
adding inert gases]
 When using the ICE tables for acid/base reactions when can the “x” be ignored [HINT: [ ]/ Ka
or Kb > 100]
Problems:
1. How will the following system at equilibrium shift in each of the following causes?
2SO3 (g) ↔ 2 SO2 (g) + O2 (g)
∆ Ho = 197 kJ
a. SO2 (g) is added.
b. the pressure is decreased by increasing the volume of the container
c. the pressure is increased by adding Ne(g)
d. the temperature is decreased
e. O2 (g) is removed
2. Given the equilibrium system
PCl5 (g) ↔ PCl3 (g) + Cl2 (g)
K = 12.5 at 60oC
A 1.0 L reaction vessel is analyzed and found to contain 3.2 mol Cl2(g), 1.5 mol PCl3 (g) and 2.0 mol
PCl5(g). Show that the reaction mixture has not yet reached equilibrium.
3. If 1.00 mol each of carbon dioxide and hydrogen is initially injected into a 10.0 L reaction chamber
at 986oC, what would be the concentrations of each entity at equilibrium?
CO2 (g) + H2 (g) ↔ CO (g) + H2O (g)
K = 1.60
4. When 2.0 mol CS2 and 4.0 mol of Cl2 are initially placed in a flask they react according to the
following equation: CS2 (g) + 3 Cl2 (g) ↔ S2Cl2 (g) + CCl4 (g). At equilibrium, 0.30 mol CCl4 is found.
How much Cl2 is left?
5. Will a precipitate form in each of the following cases? Use Ksp tables to justify your answer.
a. 50 mL of 0.040 mol/L Ca(NO3)2 (aq) plus 150 mL of 0.080 mol/L (NH4)2SO4 (aq)
b. 10 mL of 0.005 mol/L AgNO3 (aq) plus 25 mL of 0.02 mol/L NaBr (aq)
6. If the [H+(aq)] in a solution is 3.0 x 10-4 mol/L, what is the hydroxide concentration?
7. A solution is made by dissolving 0.80 g Ca(OH)2 (s) in water to make 100 mL of final solution.
Calculate the pH of the solution.
8. Lactic acid, HC3H5O3 (aq) is a weak acid that gives yogurt its sour taste. Use HA to represent lactic
acid. Calculate the pH of a 0.0010 mol/L solution of lactic acid. The Ka for lactic acid is 1.4 x 10-4.
9. Strychnine C21H22N2O2 (aq) is a weak base but a powerful poison. Calculate the pH of a 0.0010
mol/L solution of strychnine. Use BOH to represent strychnine. The Kb of strychnine is 1.0 x 10-6.
10. 28 mL of 0.36 mol/L acetic acid is titrated with a standardized 0.43 mol/L KOH solution.
Calculate the pH of the solution after 21 mL of the KOH solution has been added. Assume the Ka of
acetic acid is 1.8 x 10-5.
Unit 5 –Electrochemistry
Theory:
 Define oxidation, reduction, oxidizing agent, reducing agent, REDOX reaction, Oxidation
number, spontaneity rule, protective coating to prevent rust, Cathodic protection, anode,
cathode, standard cell potential
 Describe the structure and function of the common dry cell. Give an example using MnO2,
NH4Cl and Zn electrode
 Predicting Spontaneity in galvanic and electrolytic cells
 Calculating ∆ EO
Problems:
1. Determine the oxidation number of
a. S in SO2
b. Cr in Cr2O72-
c. N in NH4Cl d. P in K4P2O7
2. Methanol reacts with acidic permanganate ions as shown below:
5 CH3OH(l) + 2 MnO4-(aq) + 6 H+(aq)  5 CH2O(l) + 2 Mn2+ + 8 H2O (l)
a. Assign oxidation numbers to all atoms/ions
b. Which atom/ion is oxidized (reducing agent)? Label the oxidation above the equation.
c. Which atom/ion is reduced (oxidizing agent)? Label the reduction below the equation.
3. Calculate EO for the following reactions:
a. Cd (s) │ Cd 2+(aq) ║ Cr2+(aq)│ Cr(s)
b. 2 Fe2+(aq) + Pb 2+(aq)  2 Fe3+(aq) + Pb (s)
4. For each of the above reactions predict if it will be spontaneous or not.