Dr. C. Rogers --- Chem 206 GENERAL CHEMISTRY II --- EXTRA PROBLEMS
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CHEMICAL EQUILIBRIA
# 1.
Decide if each of the following statements is true or false. If false, change the wording to
make it true.
a) The magnitude of the equilibrium constant is always independent of temperature.
b) When two chemical equations are added to give a net equation, the equilibrium constant for
the net equation is the product of the equilibrium constants of the summed equations.
c) The equilibrium constant for a reaction has the same value as K for its reverse reaction.
d) Only the concentration of CO2 appears in the equilibrium constant expression for the
reaction:
CaCO3(s)
CaO(s) + CO2(g)
e) For the above reaction, the value of K is numerically the same whether the amount of CO 2 is
expressed as moles per litre or as gas pressure.
# 2.
If the reaction quotient is smaller than the equilibrium constant for a reaction such as
A
B,
does this mean that reactant A continues to be consumed for form B, or
does B form A, as the system moves to equilibrium?
# 3.
The decomposition of calcium carbonate (equation shown in question #1) is an endothermic
process. Explain how increasing the temperature would affect the equilibrium (use Le Chatelier's
principle first to help you predict, and then use the more rigorous energetics descriptions used in
class to explain why in terms of kinetics). If more calcium carbonate to a flask in which this
equilibrium exists, how is the equilibrium affected? What if some additional carbon dioxide is
placed in the flask?
# 4.
Characterize each of the following as product- or reactant- favoured.
a) CO(g) + ½ O2(g)
CO2(g)
Kp=1.2x1045
b) H2O(g)
H2(g) + ½ O2(g)
Kp=9.1x10-41
c) CO(g) + Cl2(g)
COCl2g)
Kp=6.5x1011
# 5.
At 2300 K, the equilibrium constant for the formation of NO(g) is 1.7x10 -3
N2(g) + O2(g)
NO2(g).
a) Analysis shows that, in a particular experiment, the concentrations of N2 and O2 are both
0.25M, and that of NO is 0.0042M. Is the system at equilibrium? Why not?
b) If the system is not at equilibrium, in which direction does the reaction proceed?
c) What is the amount of free energy that dissipates when the system goes from the conditions
listed in part (a) to equilibrium?
d) When the system reaches this new state of equilibrium, what are the equilibrium
concentrations of all species?
Dr. C. Rogers --- Chem 205 GENERAL CHEMISTRY I --- SAMPLE TEST QUESTIONS
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# 6.
Sulfur dioxide is readily oxidized to sulfur trioxide: SO2(aq) + O2(g)
SO3(g)
At a specific temperature, K = 279. If we add 3.00 g of SO2 and 5.00 g of O2 to a 1.0L flask,
approximately what quantity of SO3 will be in the flask once the reactants and product reach
equilibrium?
 note: the full solution to this problem results in a cubic equation. Do not try to solve it
exactly. Decide only which of the following answers is most reasonable:
a) 2.21 g
b) 4.56 g
c) 3.61 g
d) 8.00 g
# 7.
The total pressure for an equilibrated mixture of N2O4 and NO2 is 1.5atm.
N2O4(g)
2NO2(g). If Kp=6.75 at 25C, calculate the partial pressure of each gas in
the mixture at equilibrium.
# 8.
An ice cube is placed in a beaker of water at 20C. The ice cube partially melts, and the
temperature of the water is lowered to 0C. At this point, both ice and water are at 0C, and no
further change is apparent. Is the system at equilibrium? How can you be sure? Remember that
chemical equilibrium is a dynamic state, where events are still occurring at a molecular level.
Suggest two different experiments you could try to test whether or not the system is really at
equilibrium (hint: consider using D2O in one of your experiments).
# 9.
Lanthanum oxalate decomposes when heated, to form lanthanum oxide, carbon monoxide
and carbon dioxide:
La2(C2O4)3(s)
La2O3(s) + 3 CO(g) + 3 CO2(g).
a) When the system has reached equilibrium, the total pressure in a 10.0L flask is 0.200 atm.
What is the value of Kp?
b) Suppose 0.100 mol of La2(C2O4)3 was originally placed in the 10.0L flask. What amount of
La2(C2O4)3 remains unreacted at equilibrium?
# 10.
Hemoglobin (Hb) can form a complex with both oxygen O2 and carbon monoxide CO. For
the reaction:
HbO2(aq) + CO(g)
HbCO(aq) + O2(g)
at body temperature, K is about 2.0x102. If the ratio of [HbCO]/[ HbO2] becomes close to 1, death
is probable. What partial pressure of CO in the air is likely to be fatal? Assume the partial
pressure of O2 is 0.20 atm.
Dr. C. Rogers --- Chem 205 GENERAL CHEMISTRY I --- SAMPLE TEST QUESTIONS
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# 11.
A reaction important in smog formation is O3(g) + NO(g)
O2(g) + NO2(g).
At typical air temperatures (whatever that is...the textbook writer clearly does not live in Montreal),
the equilibrium constant is K = 6.0x1034.
a) If the initial concentrations are [O3] = 1.0x10-6M, [NO] = 2.5x10-4M, [O2] = 18.2x10-3M, is the
system at equilibrium? If not, in which direction does the reaction proceed?
b) If the temperature is increased, as on a very warm day, will the concentrations of the
products increase or decrease? (Hint: you have to calculate the enthalpy change of the
reaction to find out if it is exothermic or endothermic).
# 12.
A solution of Co2+ ions in water containing hydrochloric acid is a system that involves the
following equilibrium:
Co(H2O)62+(aq) + 4 Cl-(aq)
CoCl42-(aq) + 6H2O(l)
The Co(H2O)62+ ion is pink, whereas the CoCl42- ion is blue (for your information, a solid-phase
version of this reaction is commonly used as a visual indicator of humidity levels in lab
desiccators). If this solution is placed in a beaker of hot water, it appears deep blue. When placed
in an ice bath, the solution changes to pink. Is the transformation of Co(H2O)62+ to CoCl42exothermic or endothermic?
# 13.
Iodine dissolves in water, but its solubility in a nonpolar solvent such as CCl4 is much
greater. The equilibrium constant is 85.0 for the process:
I2(aq)
I2(CCl4).
You place 0.0340 g of iodine in 100.0mL of water. After shaking it with 10.0 mL of CCl 4, then
letting the two immiscible solvents separate out into layers, how much I2 remains in the water
layer?