Practice Exam 3: Thermodynamics, Kinetics, Equilibrium and Acid-Base
Name
Directions: Show all work.
1. Select all true statements below regarding the effects of temperature on the parameters of a reaction (3)
a.) The rate constant of a reaction changes with temperature.
b.) The value of the equilibrium constant of a reaction increases with temperature.
c.) The fraction of molecules having thermal energy greater than Ea increases with temperature.
e.) Entropy has a dominant impact on the value of the equilibrium constant at high temperature.
2. Select all true statement(s) below regarding a system at equilibrium. (3)
a.) Once equilibrium is reached, the reaction stops.
d.) The system has attained maximum order
b.) ΔG for a reaction approaching equilibrium is negative
e.) At equilibrium, ΔG = 0 kJ/mol
c.) K< 1 suggests that the reaction is favored to the left.
3. For the spontaneous combustion of sodium metal: 2Na(s) + O2 (g)  Na2O (s)
ΔH = -416 kJ/mol
Select the true statement(s) below. (3)
a.) The reaction is enthalpically (energetic) and entropically driven, and will be spontaneous at all temperatures.
b.) The reaction is entropically driven, but not enthalpically driven.
c.) The reaction is enthalpically driven, but not entropically driven.
d.) | ΔH | > | TΔS |
e.) | TΔS | > | ΔH |
4. The HA- anion has a Ka value of 4.7 x 10-5 M. Select all true statements below. (3)
a.) A2- (aq) must be neutral
d.) HA- (aq) is a stronger base than H2A (aq)
b.) HA- (aq) is a weaker acid than H2A
e.) A2-(aq) is a weak, Bronsted base
c.) H2A (aq) must be a strong acid
5. The non-elementary reaction A(aq) + B(aq)  AB (s) is found to be second order with respect to A and second
order overall. Select the true statement below: (3)
a.) B(aq) reacts to form an intermediate.
b.) A(aq) reacts to form an intermediate.
c.) The reaction rate increases with increasing concentrations of B(aq).
d.) The reaction rate increases with increasing concentrations of A(aq).
6. The combustion of iron metal leads to the formation of rust, Fe2O3(s). This is a spontaneous process. What are
the signs of ΔG, ΔH and ΔS? What drives the reaction? Under what circumstance would the reaction not be
spontaneous? (7)
ΔG (negative, spontaneous), ΔH (negative, combustion), ΔS (negative, produces order)
The reaction must be driven by the enthalpy, because the entropy change of the system is
unfavorable. Therefore, ΔH > TΔS. The reaction would reverse direction if T were very large.
7. In terms of G and S, explain why systems approach equilibrium (4 sentences or less). You may include a figure
if you deem this helpful. (5)
Reactions approach equilibrium because it is the state of maximum disorder, which means that
the change in Gibbs free energy is also maximized.
8. Based on the definition of “irreversible” from thermodynamics, explain why spontaneous reactions are
irreversible (3 sentences or less). (5)
Though physically reversible, spontaneous processes are not thermodynamically
reversible because work must be done on the system to reverse the process, which
will therefore change the surroundings permanently.
9. Use the data shown to the right for the reaction H2O 2H + O
to answer the following questions.
1
Write the rate law of the reaction, including k. Show units. (6)
0.8
*Reaction is clearly 2nd order based on units of y-axis and sign of
slope. Value of k approximated by taking slope of line.
1/[H2O]
Rate = 0.03 M-1s-1 [H2O]2
0.6
0.4
0.2
What is the rate of the reaction at t= 25 sec? (7)
*You must know [H2O] at 25 sec. Use 2nd order
equation!
(
)
*Use rate law!! Rate = 0.03 M-1s-1 (1.05 M)2 = 0.033 M/s
0
0
5
10
Time (sec)
15
20
10. The pKa values of oxalic acid, H2C2O4, are 1.25 and 3.81. Write the forms of the molecule, including charge,
that you would expect to find at the following pH values. Specify relative abundances.
pH = 9
Mainly C2O42-
pH = 2.4
Mostly HC2O4- with some H2C2O4
pH = 1
Mainly H2C2O4
pH = 4
Mostly C2O42- with some HC2O4-
11. Consider the reaction between hydrofluoric acid (aq) and ammonia (aq)
T = 25oC
Initially, [hydrofluoric acid]=1M and [ammonia] = 0.74M. The reaction is allowed to proceed to equilibrium, at
which point, [hydrofluoric acid]=0.4M and [ammonia] = 0.14M.
a.) Write out the full, balanced reaction. Determine the value of K. (5)
HF (aq) + NH3 (aq) ⇌ F-(aq) + NH4+ (aq)
(
)(
)
(
)(
)
b.) Now, the system is purged, causing the ammonia concentration to drop by 31%. Determine ΔG and the new
equilibrium concentrations. (x values are not required. 1.5% bonus credit will be given for a correct value) (10)
New INITIAL concentration of NH3 = 0.0966 M
)(
(
(
(
)(
)(
)
(
)
(
)
)
)
5.42x2+4.388x-0.1122 = 0
Solve quadratic formula: x = 0.0248M
New concentrations!! [HF]= 0.4248M, [NH3] = 0.1214M, [F-] = [NH4+] =0.5752M
12. Determine the pH and [H3O+] of a 0.038M solution of methylamine, CH3NH2(aq) given that the Kb of
methylamine is 4.6 x 10-4M.
pOH = 2.38, pH = 11.62, [H3O+] = 10-11.62 = 2.4 x 10-12 M
13. Sketch a reaction curve for a reaction having an activation energy of 500 KJ/mol and ΔH0rxn = +200 KJ/mol.
The reaction progress should go on the x-axis. Also, label the transition state. (4)Equations and constants
Enthalpy
Transition state (intermediate)
ΔH =200 kJ/mol
Ea =500 kJ/mol
Reaction Progress
ΔGo = ΔHo -TΔSo
ΔG0 = -RT ln K (R = 8.314 J/mol K)
pH = -log [H3O+]
ΔG = RT ln (Q/K)
pH + pOH = 14
ln (K2/K1) = (ΔH0/R)(
–
pKa = -log Ka
)
1st order: ln [A]t = -kt + ln [A]o
2nd order: 1/[A]t = kt + 1/[A]o
k = Ae-Ea/RT
ln K = -ΔH/(RT) + ΔS/R
eln(x) = x
10log(x) = x
10-pH = [H3O+]