5.2 & 5.3 Equilibrium and
Chemical Equilibrium
5.1.Can a battery that has a voltage be considered a system
at equilibrium? How about a dead battery? Justify each con-
clusion.
5.2.What is the difference between a static equilibrium and
a dynamic equilibrium? Give examples different from the ex-
amples in the text. What is similar for the two types of equi-
libria?
5.3.Which system in each pair best represents equilibrium
species under standard conditions of temperature and pres-
sure? Be prepared to justify your choice.
(a)Rb & H 2 O or Rb& OH& H 2
(b)Na & Cl 2 or NaCl (crystal)
(c)HCl & H 2 O or H(aq) & Cl(aq)
(d)C (diamond) or C (graphite)
5.4.Supersaturatedsolutions can be made in which more
solute is dissolved in solution than would normally dissolve.
These solutions are inherently unstable, however. A seed crys-
tal of calcium acetate, Ca(C 2 H 3 O 2 ) 2 , precipitates the excess
solute from a supersaturated solution of calcium acetate. When
the excess solute has finished precipitating, a chemical equi-
librium is established. Write the chemical equations for that
equilibrium, and write the net chemical reaction that occurs
overall.
5.5.Following is a chemical reaction between zinc metal and
hydrochloric acid in a closed system:
Zn (s) 2HCl (aq) →H 2 (g) ZnCl 2 (aq)If the initial amounts present are 100.0 g of zinc and 150.0 mL
of 2.25 M HCl, determine maximum and minimum possible
values of for this reaction.
5.6.The following is a reaction with its initial conditions
(amounts of each substance):
6H 2 P 4 → 4PH 3
10.0 mol 3.0 mol 3.5 mol(a)Determine if 1.5 mol of P 4 reacts to make products.
(b)Is it possible for to equal 3 in this case? Why or why not?
5.7.The hemoglobin in blood establishes an equilibrium with
oxygen gas very quickly. The equilibrium can be represented as
heme O 2 heme O 2where “heme” stands for hemoglobin and “heme O 2 ” stands
for the hemoglobin-oxygen complex. The value for the equi-
librium constant for this reaction is about 9.2 1018. Carbon
monoxide also binds with hemoglobin by the following re-
action:
heme CO heme COThis reaction has an equilibrium constant of 2.3 1023. Which
reaction’s equilibrium lies farther toward products? Does your
answer justify the toxicity of CO?
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5.8.1.00 g of sucrose, C 12 H 22 O 11 , dissolves completely in
100.0 mL of water. However, if 200.0 g of sucrose were added
to the same amount of water, only 164.0 g would dissolve.
Write the equilibria reactions for both systems and comment
on their differences.
5.9.If N 2 , H 2 , and NH 3 gases were contained in a system
such that the total pressure were 100.0 bar, then the p° terms
in equation 5.9 would be equal to 100.0 bar. True or false?
Explain your answer.5.10.Determine (^) rxnG° and (^) rxnGfor the following reaction
at 25°C, using data in Appendix 2. The partial pressures of the
products and reactants are given in the chemical equation.
2CO (g, 0.650 bar) O 2 (g, 34.0 bar)
2CO 2 (g, 0.0250 bar)
5.11.In atmospheric chemistry, the following chemical reac-
tion converts SO 2 , the predominant oxide of sulfur that comes
from combustion of S-containing materials, to SO 3 , which can
combine with H 2 O to make sulfuric acid (and thus contribute
to acid rain):
SO 2 (g) ^12 O 2 (g) SO 3 ()
(a)Write the expression for Kfor this equilibrium. (b)Calculate
the value of G° for this equilibrium using the (^) fG° values in
Appendix 2. (c)Calculate the value of Kfor this equilibrium.
(d)If 1.00 bar of SO 2 and 1.00 bar of O 2 are enclosed in a
system in the presence of some SO 3 liquid, in which direction
would the equilibrium move?
5.12.Assume that a reaction exists such that equilibrium oc-
curs when the partial pressures of the reactants and products
are all 1 bar. If the volume of the system were doubled, all of
the partial pressures would be 0.5 bar. Would the system still
be at equilibrium? Why or why not?
5.13.Show that KK1/2if the coefficients of a balanced
chemical reaction are all divided by two. Give an example.
5.14.The balanced chemical reaction for the formation of
ammonia from its elements is
N 2 3H 2 (g) 2NH 3 (g)
(a)What is (^) rxnG° for this reaction? (b)What is (^) rxnGfor this
reaction if all species have a partial pressure of 0.500 bar at
25°C? Assume that the fugacities are equal to the partial pres-
sures.
5.15.The answers in exercise 5.14 should show that chang-
ing the partial pressure changes the instantaneous (^) rxnGeven
though the ratio of partial pressures stays the same (that is,
1 1 1 for standard pressure conditions is equal to 0.50.50.5
for the given conditions). This suggests the interesting possi-
bility that at some equal partial pressure pof all components,
the reaction reverses; that is, the instantaneous (^) rxnGbecomes
negative. Determine pfor this equilibrium. (You will have to
use the properties of logarithms as mentioned in the chapter
to find the answer.) Is your answer of value to those who work
with gases at high pressures, or at low pressures? What is your
reasoning?
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138 Exercises for Chapter 5
EXERCISES FOR CHAPTER 5