Chemistry Study Guide
Answer the following questions on a separate sheet of paper.
Chapter 10 States of Matter
1. State the 5 basic assumptions of the kinetic molecular theory.
2. What is the relationship between the temperature, speed and kinetic energy of gas molecules?
3. Explain how the attractive forces between particles in a liquid are related to the equilibrium vapor pressure of that liquid.
4. What is the relationship between atmospheric pressure and the boiling point of a liquid?
5. Water has a high heat of fusion (6.009 KJ/mol) and a high heat of vaporization (40.79 KJ/mol). Explain what this means in
terms of attraction between particles.
6. Using the values given in #5, calculate the amount of energy needed to melt 7.95x10 5 g of ice. Is this an exothermic or
7. Explain why the water molecule is polar. In your explanation, include why the water molecule is bent?
8. From your knowledge of intermolecular forces, arrange the following in order of increasing surface tension (least to most):
Water, hexane, ethanol, ethanal
9. Describe how the intermolecular forces in water allow for each of the following properties of water:
a. low vapor pressure c. solid H2O is less dense than liquid H2O
b. high heat of vaporization d. high boiling point for a molecule of its mass
Chapter 11 Gases
1. Explain what happens to the pressure inside a balloon when you blow into it.
2. Why can’t an ideal gas be liquefied?
3. What would be the number of moles and grams of carbon dioxide gas contained in an 855 mL container at 35˚C and 1860
mmHg of pressure?
4. If a 2.75 L container of gas at 24.0˚C and 95.2 kPa was compressed to 1820 mL and warmed to 40.0˚C, what would be the new
5. A mixture of 3 gases, A,B and C, is at a total pressure of 6.11 atm. The partial pressure of A is 168 kPa and B is 3.89 atm. What
is the partial pressure of gas C?
6. Ammonia and ethanol are released at the same time across a room. Which will you smell first?
7. What is the volume of 8.00 grams of oxygen at STP?
Chapter 16 Thermochemistry
1. How much heat, in calories, does 32.0 g of water absorb when it is heated from 25ºC to 80ºC? How many joules is this?
2. The temperature of a piece of copper with a mass of 95.4 g changes from 25.0ºC to 48.0ºC when the metal absorbs 849 J of
heat. What is the specific heat of copper? Is this an exothermic or endothermic reaction?
3. Will the specific heat of 50.0 g of a substance be the same as or greater than the specific heat of 10.0 g of the same substance?
4. If 28.2 g of CaCl2 are dissolved in 125 mL of H2O at 25.0˚C and the final temp. reaches 42.0˚C, calculate the ΔH in kJ and
ΔHsoln in kJ/mol. Also diagram this process.
5. Use the following equation to calculate the amount of energy released when 725 g of CaO are dissolved in water:
CaO(s) + H2O(l) Ca(OH)2(s) + 65.2 kJ
Chapter 12 Solutions
1. Define the following terms:
a. solubility b. saturated solution c. unsaturated solution d. Henry’s law
2. How does each of the following affect the solubility of (a) a solid dissolved in a liquid, and (2) a gas dissolved in a liquid.
a. an increase in temperature c. an increase in pressure
b. shaking, agitation d. an increase in pressure with a decrease in temperature
3. Differentiate between the following:
a. a dilute unsaturated solution and a dilute saturated solution
b. a concentrated saturated solution and a concentrated unsaturated solution
4. Calculate the following (SHOW ALL WORK):
a. The molarity of a solution containing 42.6g of sodium hydroxide in 3.00L of water.
b. The number of moles of solute present in 680mL of a 0.25M Na2SO4 solution.
c. The number of grams of KBr present in 500mL of a 0.100 M solution.
5. Describe how to prepare the following solutions. Include calculations, a description of the procedure, and specific equipment.
a. 400.0 mL of 0.15 M solution of copper (II) sulfate from 0.75M stock solution.
b. 50.0 mL of a 0.20 M solution of potassium nitrate from a 4.0 M stock solution.
6. Calculate the following (SHOW ALL WORK):
a. The maximum amount of KCl that can be dissolved in 45.0g of water at 20oC (solubility of KCl is 34g/100g H2O).
b. A student dissolved 52.0g of sodium fluoride in 125g of water (solubility of NaF is 12g/100g H 2O at 25oC). Is the solution
saturated or unsaturated?
Chapters 17-18 Rates of Reaction and Equilibrium
1. Define the following terms:
a. activation energy c. chemical equilibrium e. Le Chatelier’s principle
b. catalyst d. collision theory f. reaction mechanism
2. Draw energy diagrams and label the reactants, products, activated complex, activation energy, and energy change for:
a. an exothermic reaction b. and endothermic reaction c. a reaction with and without the presence of a catalyst
3. Use the collision theory to explain how and why each of the following affects the rate of a reaction:
a. an increase in temperature c. particle size
b. a decrease in the concentration of a reactant d. use of a catalyst
4. Predict how each of the following changes will affect the equilibrium position of the given reaction:
PCl5(g) PCl3(g) + Cl2(g) + energy
a. increase the pressure c. decrease the amount of Cl2 e. increase the volume of the container
b. decrease the temperature d. add more PCl5 to the reaction vessel
5. Write the equilibrium constant expression for each of the following reactions: (NOTE – you must first balance the equations!)
a. NH3 (g) + O2 (g) NO (g) + H2O (g) c. COCl2 (aq) CO (g) + Cl2 (g)
b. N2O4 (g) NO2 (g) d. CaO (s) + CO2 (g) CaCO3(s)
6. For each of the following reactions, determine whether the reactants or the products are favored:
a. H2Cr2O7 (aq) + 4 NaOH (aq) 2 Na2CrO4 (aq) + 3 H2O(l) Keq = 5.62 x 1049
b. N2 (g) + O2 (g) 2NO (g) Keq = 1.2 x 10-4
7. Calculate Keq for the following general equations, given the concentration provided:
a. 2 A + B 2C [A] = 0.6M [B] = 0.1M [C] = 0.4 M
b. A + 3B C [A] = 1.2M [B] = 0.4M [C] = 0.2 M
c. For the reactions above, are the products or the reactants favored?
Chapter 13-15 Acids and Bases, Neutralization and Titration
1. Differentiate between strong electrolytes, weak electrolytes, and nonelectrolytes. Give two examples of each and explain why
you placed them in each category.
2. Predict the solubility of each of the following substances in: (1)water and (2) heptane
a. sodium iodide c. hydrogen bromide e. ethanol g. tetrachloromethane
b. nickel (solid) d. calcium carbonate f. benzene
3. Why do you believe each of the above is or isn’t soluble in the solvents mentioned?
4. List at least three properties of acids and three properties of bases.
5. Differentiate between Arrhenius, Bronsted-Lowry, and Lewis acids and bases. Give an example of each.
6. For the following reaction, label the acid, base, conjugate acid, and conjugate base. List the conjugate acid-base pairs.
CH3COOH + H2O CH3COO- + H3O+
7. Calculate the following:
a. [OH] when [H+] of a solution is .0024 M c. pH of a .025M solution of HCl e. pH of a 2.20M solution of Ca(OH)2
b. [H+] when [KOH] is 1.6 x 10-12 d. pOH of a .025M solution of HCl
8. Show the balanced equation for the self-ionization of water. What is the pH at 25oC?
9. Differentiate between the following and give an example of each:
a. strong acids and weak acids b. strong bases and weak bases
10. Explain how it is possible to have:
a. a concentrated weak acid b. a dilute strong acid
11. Predict the products for the following acid-base reactions & write balanced equations.
a. hydrochloric acid + sodium hydroxide c. ammonia + sulfuric acid
b. carbonic acid + calcium hydroxide d. ethanoic acid + potassium hydroxide
12. Calculate the following:
a. The number of moles of NaOH needed to neutralize 3.5 moles of H 2CO3
b. The number of moles of HCl needed to neutralize 6.0 moles of Mg(OH) 2
13. Calculate the following:
a. [HCl ] if 2.0 mL was needed to neutralize 40.0 mL of a 0 .20M NaOH solution
b. Number of moles of Mg(OH)2 present in 50.0 mL of solution, if 75.0 mL of a 0.1 M HCl solution was needed to reach the
c. The molarity of a NaOH solution, if 20.0 mL of the solution was neutralized by 28.0 mL of a 1.0 M H 3PO4 solution