Lecture #26
What’s on the Final?
Chemistry 142 B
Autumn Quarter, 2004
J. B. Callis, Instructor
General
• Date as Announced in Syllabus
• 44 Multiple-Choice Questions - Chapters 1-8
• 22 questions from chapters 1-5
• 22 questions from chapters 6-8
• Practice exam available on the class web site – key on
class bulletin board.
Chapter 2
Dalton’s Atomic Hypothesis
and Beyond
• Law of multiple proportions as explained by
Dalton
• Naming compounds
Problem 26-1: Multiple Proportions
A chemist studies three oxides of iodine, and
finds their % oxygen as follows: 20.14 %,
23.97% and 22.10 %. (a) Calculate the mass of
oxygen per gram of iodine in each compound.
Express the result as a ratio. Then form the
ratio of ratios by dividing by the ratio of the
first compound. (b) How do the numbers in
part (a) support Dalton’s atomic theory?
Problem 26-1: Multiple Proportions
Cmp
Ratio of
(symbolic Ratios
rep)
IO2
(IiOn)
Amt O, Amt I, mO/mI
g
g
(from Dalton)
n  i 2 1

1
i  n 1 2
20.14
79.86
I2O5
(IjOo)
o  i 5 1 5


j n 22 4
23.97
76.03
I4O9
(NkOp)
p  i 9 1 9


k n 42 8
22.10
77.90
Ratio of Ratios
(from masses)
Chapter 3 - Stoichiometry
• Convert between moles, mass and number of
molecules.
• Determine empirical formula from elemental
composition.
• Balance a chemical reaction using the atom
balance equations.
• Solve stoichiometric problems involving a
limiting reagent using the reactant ratio method.
Problem 26-2: Bromoform is 94.85% bromine, 0.40% hydrogen
and 4.75% carbon by mass. Determine its empirical formula.
Problem 26-3: Balance the following chemical equation:
x NaCl + y SO2 + z H2O + w O2
=> u Na2SO4 + v HCl
Problem 26-4: Sulfuric acid (H2SO4) forms in the following reaction:
2 SO2 + O2 + 2 H2O => 2 H2SO4
Suppose 400. g SO2, 175. g O2 and 125. g H2O are mixed and the reaction proceeds
until one of the reactants is used up. Identify the limiting reactant and determine
what masses of the other reactants remain.
Chapter 4
Chemical Reactions
• Classify a given reaction and identify products
such as precipitates.
• Write the net ionic reaction.
• Determine the oxidation number of an element in
a compound.
• Balance a redox equation using the atom balance
and charge balance equations.
• Perform stoichiometric calculations on solutions
Problem 26-5: Balance the following chemical equation:
x H+ + y H2O2 + z Fe2+ => u Fe3+ + v H2O
Chapter 5 – The Gas Phase
• Use the ideal gas law to compute density from
other independent variables, e.g. pressure,
temperature.
• Work with partial pressures.
• Calculate the amount of gas collected over water.
• Use the ideal gas law in stoichiometric problems.
Problem 26-6: A 20.6 L sample of air is collected in Greenland at –20.
oC at a pressure of 1.01 atm and forced into a 1.05 L bottle for shipment
to Europe for analysis. Compute the pressure inside the bottle as it is
opened in the laboratory at 21. oC.
Chapter 6 – Chemical
Equilibrium
• Write the equilibrium constant for a
balanced chemical reaction.
• Solve equilibrium problems given starting
concentrations and K.
• Apply Le Chatelier’s principle to determine
direction of shift of reaction in response to a
change in reaction conditions.
Problem 26-7: At 5000 K, even the nitrogen molecule (N2)
breaks down (into 2 atoms of nitrogen). At this temperature,
when the total pressure of nitrogen is 1.00 atm, N2(g) is a =
0.65% dissociated at equilibrium: N2(g) = 2 N(g). Compute the
equilibrium constant at 5000K.
Pressure, atm
Initial P
Change in P
Equilibrium P
N2(g) =
2 N(g)
Problem 26-7 (cont.)
Chapter 7 – Acid/Base Equilibria
• Understand the pH scale
• Calculate the pH of strong, weak and very
weak acids and bases
Problem 26-8: What is the pH of 0.15 M methylammonium
bromide, CH3NH3Br? (Kb of CH3NH2 = 4.4 x 10-4
Chapter 8 – Applications of
Aqueous Equilibria
• Calculate the pH of buffer solutions
• Give recipes for making up buffers
• Calculate pH at different points in a titration
curve
• Perform solubility calculations
• Understand the basic ideas of complex ion
equilibria
Problem 9: You have at your disposal an ample quantity of a
solution of 0.0500 M NaOH and 500 mL of a solution of 0.100 M
formic acid (HCOOH, Ka = 1.77 x 10-4). How much of the NaOH
solution should be added to the acid solution to produce a buffer of
pH 4.00?
Ans: Use the base to produce a sufficient amount of the formate ion to
provide a buffer of the desired pH. Do the problem at the HendersonHasselbalch level, ignoring the ionization of water. Allow for dilution
of the original acid solution with the base, just as in a titration.
The relevant reaction is:
Problem 9 (cont.):
Problem 9 (cont.):
Problem 10: A saturated solution of Mg(OH)2 at 25oC is
prepared by equilibrating solid Mg(OH)2 with water.
Concentrated NaOH is then added until the solubility of
Mg(OH)2 is 0.001 times that in H2O alone. (Ignore the change in
volume resulting from the addition of NaOH.) The solubility
product Ksp of Mg(OH)2 is 1.2 x 10-11 at 25oC. Calculate the
concentration of hydroxide ion in the solution after the addition
of the NaOH.
Ans: First calculate the solubility of Mg(OH)2 in water. Then
calculate the concentration of [Mg2+] after addition of OH- put this
into the mass action expression and solve for [OH-].
Problem 10 (cont.):
Addition of OH- shifts the equilibrium to the left and the
concentration of Mg2+ must diminish to maintain the
mass-action expression at a constant value. After the
addition of base, the new concentration of Mg2+ is
Answers to Questions from Lecture 26
1. 1, 1.25, 1.125
2. CHBr3
3. 4 NaCl + 2 SO2 + 2 H2O + O2
=> 2 Na2SO4 + 4 HCl
4. SO2 is the limiting reactant; 75 g O2 remaining; 13 g H 2 O remaining
5.
2 H+ + H2O2 + 2 Fe2+ =>
6. 23.0 atm
7. K = 1.67 x 10-4
8. pH = 5.73
9. 639 mL
10. 9.1 x 10-3
2 Fe3+ + 2 H2O