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HONORS CHEMISTRY – FINAL EXAM REVIEW – SPRING 2014
STRATEGY: Start by reading through your notes (chapters 8 – 15 and 18) to refresh your memory on these topics. Then, use this
review sheet as a starting point to identify the areas on which you need to spend more study time. For those areas, go back to
homework assignments, and notes to practice more problems. These questions are only samples and do not include specific examples
of how vocabulary and other conceptual information might appear in a multiple-choice or other format.
FORMAT:
 Questions will include 65 multiple-choice and 2 Essay/Calculations.
 You will be given a periodic table and formula sheet.
Chapter 8: Balancing Equations
Balancing the following equations and label the reactions as either: Decomposition (D), Synthesis (S), Single Replacement
(SR), Double Replacement (DR), Combustion (C), Neutralization (N).
1. H2 + O2  H2O
2. H3PO4 + KOH  K3PO4 + H2O
3. K + B2O3  K2O + B
4. Cs + N2  Cs3N
5. C6H6 + O2  CO2 + H2O
6. BaCl2(aq) + K2(SO4)  KCl + BaSO4
7. N2 + H2  NH3
8. C10H22 + O2  CO2 + H2O
9. Al(OH)3 + HBr  AlBr3 + H2O
10. CaBr2 + Na2CO3  CaCO3(aq) + NaBr(aq)
Chapter 9: Stoichiometry
1. Ideal combustion, resulting in only carbon dioxide and water, rarely happens. In general, at the very least, some
carbon monoxide is also produced. The real combustion of methane is more closely represented, then, by the
unbalanced equation CH4 + O2  H2O + CO2 + CO. If 0.100 mol CH4 is allowed to react with 3.20 g O2, what is
the limiting reagent, how many grams of carbon monoxide will be produced, and how many liters of the excess
reactant will be left over when the reaction has gone to completion, assuming the reaction occurs at STP?
2. Elemental zinc, Zn, reacts violently with elemental sulfur, S8. If 1020. g of zinc react with 1.210 x 1024 atoms of
sulfur, what is the limiting reagent, how many moles of product are produced?
3. 25.0 mL of 0.350 M NaOH are added to 45.0 mL of 0.125 M copper (II) sulfate. How many grams of precipitate
will be formed? What are the molar concentrations of all the species remaining in the solution?
Chapter 10: States of Matter - Solids Liquids and Gases
1. Which of the following is NOT a characteristic of ionic substances?
a. Are usually gases at room temperature.
c. Have high melting points.
b. Conduct electricity in solution form.
d. Usually dissolve in water.
2. Which of the following is NOT a characteristic of metallic substances?
a. Malleable and ductile.
c. Have low melting points.
b. Conduct electricity.
d. Are usually solids at room temperature.
3. Which of the following is NOT a characteristic of covalent substances?
a. Have low melting points.
c. Form individual molecules.
b. Sometimes dissolve in water.
d. Conduct electricity.
4. What is the difference between an intermolecular force and an intramolecular force?
- What are the three types of intramolecular forces?
- What are the three types of intermolecular forces?
- Arrange the intermolecular and intramolecular forces from weakest to strongest.
5. What causes dipole-dipole interactions?
6. Why is hydrogen bonding only possible with hydrogen?
7. What are the three possible elements that hydrogen atom must be attached to, in a compound, in order for the
compound to contain hydrogen bonding?
8. Explain how London dispersion forces arise. Although London dispersion forces exist among all molecules, for
what type of molecules are they the only major intermolecular forces? Are London dispersion forces relatively
strong or relatively weak? Explain.
9. What causes surface tension?
10. Why is ice less dense than water?
11. What is the reasoning behind Chlorine being in a gaseous state, bromine being in a liquid state, and iodine being
in a solid state?
12. Specify the predominant force of attraction that would affect the boiling point of the compound and list it
immediately following the substance. Then in the last column, indicate which member of the pair you would
expect to have the higher boiling point. Possible forces of attraction are: metallic bonding, ionic bonding,
network covalent bonding, hydrogen bonding, dipole-dipole and London dispersion forces
Substance #1
a. HCl(l)
Predominant
Intermolecular Force
Substance #2
I2
b. CH3F
CH3OH
c. H2O(l)
H2S
d. SiO2(s)
SO2
e. Fe(s)
f. CH3OH(l)
Substance with
Higher Boiling Point
Kr
CuO(s)
g. NH3
CH4
h. HCl(g)
NaCl
i.
Cu(s)
SiC
Predominant
Intermolecular Force
13. Draw a Heating Curve Diagram and label where each of the following would either be located on the
curve or label on what part of the curve it would be used to complete calculations.
a. Solid
e. Liquid
i. Vapor
b. Specific Heat of Solid
f. Specific Heat of Liquid
j. Specific Heat of Vapor
c. Melting/Freezing
g. Vaporizing/Condensing
d. Heat of Fusion
h. Heat of Vaporization
14. What formula is used to calculate the amount of energy needed to change the phase of a substance?
15. CCl2F2 has a boiling point of -30. C, and a heat of vaporization of 0.165 kJ/g. The vapor and the liquid have
specific heats of 0.61 J/gK and 0.97 J/gK respectively. How much heat must be evolved when 10.0 g of CCl2F2
is cooled from +40. C to -40. C?
16. For the following Phase Change Diagrams, identify what states are in sections A, B, and C and what occurs when
each line is crossed (in each direction). What is the point D called? What is the point E called? Which of the
diagrams is the Change of State Diagram fro water? How can you tell?
Chapter 11: Gas Laws
1. What does Boyle’s Law state, what is the formula, and what is held constant?
2. What does Charles’ Law state, what is the formula, and what is held constant?
3. What does Avogadro’s Law state, what is the formula, and what is held constant?
4. What is the formula for the Combined Gas Law and what is held constant?
5. What is the formula for the Ideal Gas Law and what is the value and units on R?
6. What is diffusion and effusion? What is the formula that is used for these calculations?
7. 1 mole of any gas at STP = ____________________
8. The conversion factors for Pressure are equal to:_______ atm = ______ mm = ______ torr = ______ kPa
9. A 2.00 L sample of a gas originally at 25 oC and a pressure of 700. torr is allowed to expand to a volume of
5.00 L. If the final pressure of the gas is 585 torr, what is its final temperature in Celsius?
10. A gas has a volume of 350. mL at 740. torr. How many milliliters will the gas occupy at 900. torr if the
temperature remains constant?
10. Calculate the number of liters occupied by the following at STP:
a. 0.200 moles of oxygen
b. 12.4 grams of chlorine gas
c. a mixture of 0.100 moles of nitrogen gas and 0.500 moles of oxygen gas
11. From the volume, temperature and pressure listed below, calculate the number of moles for the gas:
750.0 cm3 at 27.0 oC and 99.0 kPa.
12. A balloon has a volume of 2.0 L indoors at a temperature of 25.0 oC. If it is taken outdoors on a very cold day
when the temperature is –28.9 oC, what will its volume be in liters? Assume a constant air pressure within the
balloon.
13. Calculate the volume of 6.00 g of hydrogen gas at 24.0 oC and 1.50 atm.
14. Calculate the mass of 24.5 mL of SO2 gas at STP.
15. If excess hydrochloric acid is added to 13.5 grams of Al, what volume of hydrogen gas will be produced if
the gas is collected at a temperature of 80.0 oC and a pressure of 750. torr?
16. At a certain temperature, the velocity of chlorine molecules is 0.0410 m/s. What is the velocity of sulfur dioxide
molecules at the same temperature and pressure?
17. What is the ratio of the speed of carbon monoxide molecules to that of nitrogen monoxide molecules when both
gases are at the same temperature?
18. The rate of effusion of an unknown gas was determined to be 2.92 times faster than that of ammonia. What is the
approximate molecular weight of the unknown gas?
19. In the reaction, N2 + H2  NH3, how many mL of nitrogen, measured at STP, are required to produce 400 mL of
NH3, measured at STP? How many mL of hydrogen are required at STP?
Chapter 12: Solutions
1. What is molarity (M), molality (m), normality (N), and mass percent?
2. What term is used to describe a substance dissolving?
3. What is a solid – solid solution of two or more metals called?
4. A mixture in which the particles are so small that they will not reflect the “light” from a laser are called ________.
5. A solution that contains a large amount of solute per amount of solvent is called a _____________ solution.
6. What is the term that is used to describe a solid being formed when two or more liquids are mixed together?
7. When a solution is made, what is the smaller portion of the solution called?
8. A solution that contains as much solute as it can hold at a given temperature is called a ____________ solution.
9. When a solution is made, what is the larger portion of the solution called?
10. A mixture in which the particles settle out upon stand is called a ____________.
11. If a substance will NOT dissolve it is said to be _____________.
12. A solution that will hold more solute is said to be ____________.
13. A solution that contains a small amount of solute per amount of solvent is called a _____________ solution.
14. A mixture in which the particles can only be seen with a laser and do NOT settle out upon standing is called a
_____________.
15. The amount of solute that can be dissolved in a solvent at a given temperature is described as _____________.
16. Due to the fact that so many substances dissolve in water, water has been called the _____________________.
17. Calculate the molarity of a solution prepared by dissolving 15.6 g of solid KBr in enough water to make 1.25 L of
solution.
18. Calculate the number of moles of Ca(OH)2 in 150.0 mL of a 0.500 M solution. How many moles of each ion are
in solution? a. 0.0750 moles of Ca(OH)2 b. 0.0750 moles of Ca+2 and 0.150 M of (OH)-1
19. A solution is prepared by adding 2.0 L of 6.0 M HCl to 500 mL of a 9.0 M HCl solution. What is the molarity of
the new solution? (Remember, the volumes are additive)
20. Convert the following Molarities to Normalities. (a. 2.5 M HCl, b. 1.4 M H2SO4, c. 1.0 M NaOH, d. 0.5 M Ca(OH)2
21. 25.0 mL of 0.339 M Magnesium nitrate and 35.0 mL of 0.238 M Sodium hydroxide are mixed, what is the total
concentration of all ions left over in the solution?
22. If 63.8 grams of Aluminum sulfide are precipitated when aluminum nitrate and sodium sulfide are mixed, how
many milliliters of 2.50 M Aluminum nitrate is needed to form that mass of precipitate? (Assume that there is an
excess of sodium sulfide dissolved in solution.)
23. A solution is made by dissolving 25 g of NaCl in enough water to make 1.0 L of solution. Assume the density of
the solution is 1.0 g/cm3. Calculate the molarity, and molality of the solution.
24. How many grams of KBr are contained in 250. grams of a 6.25 % KBr solution?
25. Answer the following questions based on the solubility graph below.
- What is the temperature (oC) at which HCl and KNO3 have the same solubility?
- What is the maximum amount of KClO 3 that can be dissolved in 25 g of water at 100 oC?
- What is the solubility of KNO3 in 100 g of water at 60 oC?
- What is the order of decreasing solubility for the substances NaCl, NaNO 3, HCl, NH4Cl in 100 g of
water at 40 oC?
Chapter 13: Colligative Properties and Solubility Rules (Net Ionic Equations)
1. What is a Colligative property?
2. What are freezing point depression, boiling point elevation, and osmotic pressure? List the formulas that are used
to calculate each.
3. What is the molal freezing-point constant (Kf) for water?
4. What is the molal boiling point constant (Kb) for water?
5. a. What is the molality of a solution of 1.25 moles of sugar dissolved in 0.750 kg of water?
b. What is the boiling point of the resulting solution?
c. What is the freezing point of the resulting solution?
6. An aqueous solution contains 3.00 % phenylalanine (C9H11NO2) (MM = 165.21 g/mole) by mass at 25.0 oC.
Assume the phenylalanine is non-ionic and volatile and that the density of the solution is 1.00 g/mL. Calculate the
following:
a. The freezing point of the solution (T f)
b. The boiling point of the solution (T b)
c. The osmotic pressure of the solution ()
7. What is the molarity of a solution in which the osmotic pressure is 80.7 atm at 27 oC? Assume the solute is a
nonelectrolyte.
8. At what temperature, in Celsius, would the osmotic pressure have to be 347.5 kPa in order to stop osmosis if
there was 24.37 g of glucose (MM = 180.18 g/mole) dissolve in enough water to make a 1.00 L solution?
9. When 5.0 g of CaCl2 dissolves in 50.0 g of water, what is the boiling point of the solution?
10. Seawater is about 3.5% (by weight) dissolved solids, almost all of which is NaCl. Calculate the normal boiling
point of seawater.
11. Calculate the freezing point of a solution of 5.00 g of diphenyl C12H10 and 7.50 g of naphthalene, C10H8 dissolved
in 200.0 g of benzene (fp = 5.5 °C and Kf = 5.12 °C/m).
12. Salt is often used to remove ice from roads and sidewalks. Explain how this process works in terms of
Colligative properties.
13. Which salt, NaCl or CaCl2, has a greater effect on freezing point? Explain.
14. Determine the products (indicating the states) and write the net ionic equation for each of the following.
a. AgNO3(aq) + KI(aq) →
b. Pb(NO3)2(aq)+ Na2S(aq) →
c. (NH4)3PO4(aq) + Ca(NO3)2(aq) →
d. Fe(NO3)3(aq) + NaOH(aq) →
15. Write a balance equation and label the states of the reactants and products for each the reactions. After writing
the balanced equation, write the net ionic equations below the balanced equation. If there is no driving force for
a reaction, write no reaction in place of the net ionic equation.
a. Hydrogen sulfide is bubbled through a solution of strontium hydroxide.
b. Aluminum metal is added to a solution of copper (II) chloride.
c. A manganese (II) nitrate solution is mixed with a sodium hydroxide solution.
d. Solutions of sodium carbonate and lead (II) nitrate are mixed.
e. Aqueous solutions of sodium phosphate and barium chloride are mixed.
f. Hydrogen sulfide gas is added to a solution of cadmium (II) nitrate.
g. A solution of sulfuric acid is added to a solution of barium hydroxide.
h. Dilute sulfuric acid is added to a solution of calcium acetate.
i. Ammonium chloride solution is added to a solution of sodium hydroxide.
j. A precipitate is formed when solutions of sodium sulfite and calcium chloride are mixed.
Chapter 14/15: Acids and Bases
1. What is the Arrhenius definition of an acid and a base?
2. What is the Brønsted-Lowry definition of an acid and a base?
3. List the 6 strong acids and the 8 strong bases.
4. What is the difference between a strong acid or base and a weak acid or base?
5. What are diprotic and polyprotic acids?
6. What are conjugate acids and a conjugate bases?
7. In each of the following chemical equations, identify the conjugate acid/base pairs and label them as weak or
strong. (HSO4-1 + H2O
SO4-2 + H3O+1) (NH3 + H2O
NH4+1 + OH-1)
8. Label each of the following as either strong conjugate acid, weak conjugate acid, strong conjugate base, or weak
conjugate base. (Na+1, CH3COO-1, Cl-1, Al+3, Mg+2)
9. Which of the following is a conjugate acid/base pair? (HCl/OCl-1, H3O+1/OH-1, H2SO4/SO4-2, HPO4-2/PO4-3)
10. Indicate whether the following will produce an acidic, basic, or neutral solution when dissolved in water.
NaNO3, BeO, NH4ClO4, NO2, CaCO3, Cs2O, CuHSO4, and Cl2O3
11. Calculate the pH, pOH and [OH-1] of a solution in which the [H+1] = 1.2 x 10-3. Calculate each significantly!
12. Calculate the pH, pOH, and [H+1] of a solution in which the [OH-1] = 2.34 x 10-5. Calculate each significantly!
13. Calculate the pH of a 0.10 M solution of Ca(OH)2.
14. Using the following Ka values, place the following acids in order of increasing acid strength.
HClO4 (ka = 1 x 107), HCN (Ka = 4.93 x 10-10), CH3OOH (ka = 1.76 x 10-5 ), HF (Ka = 3.53 x 10-4)
15. Using the Ka values from above, calculate the Kb values and place the following bases in
order of increasing strength. CN-1, F-1, CH3OO-1, ClO4-1 (Kw = Ka x Kb)
16. Draw titration curves for the following titrations and label the x-axis, y-axis, and equivalence point :
a. 0.100 M NaOH is titrated into a solution a solution of 0.200 M HNO3.
b. 1.0 M HCl is titrated into a solution of 0.50 M NaOH.
17. What volume of 0.050 M NaOH(aq) will completely neutralize 100.0 mL of 0.075 M HBr solution? 150 mL
18. What is the conjugate base for each of the following acids? (HSO3 -1, H2S, HF, CH3COOH)
19. Which of the following ions can react as a buffer? For each buffer, write a reaction for how it would react with a
strong acid and a reaction for how it would react with a strong base.(HSO4 -1, HS -1, HCO3 -1, C2H3O2 -1)
20. Which of the following would be an example of a buffered solution? (0.30 M HF / 0.30 M HCl,
0.30 M HC2H3O2 / 0.30 M NaC2H3O2, 0.30 M NaCl / 0.30 M NaOH, 0.30 M HClO4 / 0.30 M NaClO4)
21. What is an electrolytic solution? What is the difference between a strong electrolyte and a weak electrolyte?
22. Which of the following solutions would be the strongest electrolytic solution?
(3.0 M HCN, 2.8 M HBr, 2.75 M HNO3, 3.1 M Mg(OH)2, 3.0 M RbOH)
Chapter 18: Equilibrium
1. What is Activation Energy? How does Activation Energy effect a reaction?
2. Draw an Activation Energy diagram for an endothermic and an exothermic reaction and label the x-axis, y-axis,
reactants, products, activation energy, and total energy change.
3. What is a Catalyst and how does a catalyst effect a reaction?
4. What occurs in a reaction that causes a Chemical equilibrium?
5. Write an Equilibrium expression for the following reactions:
a. CaCO3(s)
CaO(s) + CO2(g)
b. Ni(s) + 4 CO(g)
Ni(CO)4(g)
c. 5 CO(g) + I2O5(s)
I2(g) + 5 CO2(g)
d. Ca(HCO3)2(aq)
CaCO3(s) + H2O(l) + CO2(g)
6. What is an equilibrium constant and what does the equilibrium constant indicate about the reaction?
7. Calculate the value of the equilibrium constant for the following system, given the data shown: H2(g) + CO2(g)
H2O(g) + CO(g) Concentrations at equilibrium: [H2] = 1.5 M, [CO2] = 2.5 M, [H2O] = 0.5 M, [CO] = 3.0 M K = 2.0
For the reaction 2 H2O(g)
2 H2(g) + O2(g) K = 2.4 x 10-3 at a given temperature. At equilibrium it is found
-1
that [H2O] = 1.1 x 10 M and [H2] = 1.9 x 10-2 M. What is the concentration of O2(g) under these conditions?
9. What is the difference between a homogeneous equilibrium and a heterogeneous equilibrium?
10. What is LeChâtelier’s Principle?
11. Consider the equilibrium PCl3(g) + Cl2(g)
PCl5(g). How would the following changes affect each gas
at equilibrium?
PCl3(g) + Cl2(g)
PCl5(g)
a. addition of PCl3
____
____
____
b. removal of Cl2
____
____
____
c. removal of PCl5
____
____
____
d. decrease in the volume of the container
____
____
____
e. addition of He without change in volume
____
____
____
8.