Final Exam, Spring 2004

1. Complete and balance the following reactions:

a. H2SO4(aq) + NaOH(aq)

b. Pb(NO3)2(aq) + NaCl(aq)

c. Zn2+(aq) + NH3(aq)

d. C6H5NH2(aq) + H2O(l)

e. Cr3+(aq) + H2O(l)

f. NiCl2(aq) + K3PO4(aq)

g. HF(aq) + CH3NH2(aq)

h. Fe(NO3)2(aq) + KCl(aq)

i. VO2+(aq) + S(s) → VO2+(aq) + H2SO3(aq)

j. [Cr(OH)4]–(aq) + ClO3–(aq) → CrO42–(aq) + Cl –(aq)

Answer

a. H2SO4(aq) + 2 NaOH(aq) → 2 H2O(l) + 2 Na+(aq) + SO42–(aq)

b. Pb(NO3)2(aq) + 2 NaCl(aq) → PbCl2(s) + 2 Na+(aq) + 2 NO3–(aq)

c. Zn2+(aq) + 4 NH3(aq) → ← [Zn(NH3)4]2+(aq)

d. C6H5NH2(aq) + H2O(l) → ← C6H5NH3+(aq) + OH–(aq)

e. Cr3+(aq) + 2 H2O(l) → ← CrOH2+(aq) + H3O+(aq)

f. 3 NiCl2(aq) + 2 K3PO4(aq) → Ni3(PO4)2(s) + 6 K+(aq) + 6 Cl–(aq)

g. HF(aq) + CH3NH2(aq) → CH3NH3+(aq) + F–(aq) (Kc = 6.6×10–4/2.4×10–11 = 2.8×107 > 100)

h. Fe(NO3)2(aq) + KCl(aq) → NR

i. VO2+(aq) + S(s) → VO2+(aq) + H2SO3(aq)

Oxidation: S(s) + 3 H2O(l) → H2SO3(aq) + 4 H+(aq) + 4 e–

Reduction: VO2+(aq) + 2 H+(aq) + e– → VO2+(aq) + H2O(l)

Net: 4 VO2+(aq) + S(s) + 4 H+(aq) → 4 VO2+(aq) + H2SO3(aq) + H2O(l)

j. [Cr(OH)4]–(aq) + ClO3–(aq) → CrO42–(aq) + Cl–(aq)

Oxidation: [Cr(OH)4]–(aq) → CrO42–(aq) + 4 H+(aq) + 3 e–

Reduction: ClO3–(aq) + 6 H+(aq) + 6 e– → Cl–(aq) + 3 H2O(l)

Net in acid: 2 [Cr(OH)4]–(aq) + ClO3–(aq) → 2 CrO42–(aq) + Cl–(aq) + 2 H+(aq) + 3 H2O(l)

Neutralization: 2 H+(aq) + 2 OH–(aq) → 2 H2O(l)

Net in base: 2 [Cr(OH)4]–(aq) + ClO3–(aq) + 2 OH–(aq) → 2 CrO42–(aq) + Cl–(aq) + 5 H2O(l)

2. Consider the following reaction:

HgCl2(aq) + C2O42–(aq) → Cl–(aq) + CO2(g) + Hg2Cl2(s)

a. Balance the reaction.

b. The following rate data was collected:

[HgCl2] (M) [C2O42–] (M) Initial rate (mol L–1 min–1)

 

0.105 0.15 1.8×10–5

 

0.105 0.30 7.1×10–5

 

0.052 0.30 3.5×10–5

 

0.052 0.15 8.9×10–6

 

Determine the order of reaction with respect to HgCl2, with respect to C2O42–, and the overall order.

c. Write the rate law for the reaction.

Answer

a. 2 HgCl2(aq) + C2O42–(aq) → 2 Cl–(aq) + 2 CO2(g) + Hg2Cl2(s)

b. Doubling the oxalate concentration quadruples the rate (experiments 1 and 2) so the order in oxalate ion is 2. Doubling the mercury(II) chloride concentration doubles the rate (experiments 2 and 3) so the order in mercury(II) chloride is 1. The total order = 2 + 1 = 3.

c. Rate = –k[HgCl2][C2O42–]2

3. Al(OH)3 is a sparingly soluble salt. Determine if the solubility will increase, decrease, or stay the same in the following solutions. Briefly (15 words or less) explain your answer.

a. 1 M HCl solution.

b. 1 M NaClO4 solution.

c. 1 M NH3 solution.

d. 1 M NaOH solution.

Answer

a. Solubility increases: acid base reaction.

b. Solubility is unchanged: no reactivity

c. Solubility is decreased: the base creates a small amount of a common ion

d. Solubility increases: complex ion formation.

4. You need to prepare a buffer with pH = 6.00 and a buffer capacity of about 0.1 M.

a. Which substances should you use to prepare the buffer?

b. 0.010 mol of NaOH is added to 1.0 L of your buffer; estimate the new pH.

c. 1.00 mol of NaOH is added to 1.0 L of your buffer; estimate the new pH.

Answer

a. For a buffer capacity of 0.10 M requires an acid with a pKa of ~ 6 (Ka ~ 10–6) and its conjugate base in concentrations in excess of ~0.10 M are required. Hydroxlamine and hydroxylammonium chloride would be a good combination (Ka ~ 1.1×10–6).

b. This is 10 % of the buffer capacity so the pH will rise only slightly, perhaps 6.05

c. This is 10× the buffer capacity so the pH will be determined by the strong base, ~14

5. An assumption that we make about the group 2 hydroxides is that both hydroxides completely dissociate. This question will test that assumption using magnesium hydroxide as an example.

a. Write the hydrolysis reaction for the magnesium ion.

b. Find ΔG° for the reaction in part a.

c. Find the equilibrium constant at 25 °C for the reaction in part a.

d. What subscript should be used with the equilibrium constant found in part c?

e. Estimate the pH and per cent ionization of a 1 M solution of magnesium ion.

f. The per cent dissociation of the second hydroxide in magnesium hydroxide is equal to 100 – (per cent ionization of magnesium ion). Based on the result in part e, is it justifiable to assume that both hydroxides should be considered strong? Why or why not?

Answer

a. Mg2+(aq) + 2 H2O(l) → ← MgOH+(aq) + H3O+(aq)

b. ΔG° = [(–626.7) + (–237.2)] – [(–454.8) + 2(–237.2)] = 65.3 kJ/mol

c. Keq = e–ΔG°/RT = e–65300/(8.314×298) = 3.6×10–12

d. Keq = Ka since the reaction describes acid hydrolysis.

e. Mg2+(aq) + 2 H2O(l) → ← MgOH+(aq) + H3O+(aq)

Ka = [MgOH+]e[H3O+]e/ [Mg2+]e = 3.6×10–12

Initial100

Change–x+x+x

Equilibrium1 – xxx

Approximation is valid:

3.6×10–12 = x2/(1)

x = [H3O+] = 1.9×10–6

pH = –log(1.9×10–6) = 5.72

α = 1.9×10–6×100 %/(1) = 1.9×10–4 %

f. Per cent dissociation of MgOH+ = 100 – 1.9×10–4 % = 99.99981%. This is well beyond the accuracy of the measurement, so approximation as 100% dissociated is well justified.

6. This question will use available data to fill in a value in the Table of Thermodynamic Quantities.

a. Write the hydrolysis reaction for Co2+(aq).

b. Find ΔG° at 25 °C for the reaction in part a.

c. Find ΔGf° for the cobalt(II) monohydroxide cation.

d. A common mistake beginning students make is to write the hydrolysis reaction for Co2+(aq) to give cobalt(II) hydroxide. Find ΔG° for this hypothetical reaction and explain why the reaction does not occur.

Answer

a. Co2+(aq) + 2 H2O(l) → ← CoOH+(aq) + H3O+(aq)

b. ΔG° = –RTlnKa = –(8.314)(298)ln(1.3×10–9) = –50.7 kJ/mol

c. ΔG°: = [ΔGf°(CoOH+) + ΔGf°(H3O+)] – [ΔGf°(Co2+) + 2(ΔGf°(H2O))]

–50.7 = [ΔGf°(CoOH+) + (–237.2)] – [(–51.5) + 2(–237.2)]

ΔGf°(CoOH+) = –339.4 kJ/mol

d. Co2+(aq) + 4 H2O(l) → ← Co(OH)2(s) + 2 H3O+(aq)

ΔG° = [(–454.4) + 2(–237.2)] – [(–51.5) + 4(–237.2)] = +71.5 kJ/mol, which is significantly nonspontaneous.

7. Consider the cell:

Co(s) | Co2+(aq) pH = 0 || AgCl(s) | Cl–(aq) | pH = 0 | Ag(s)

a. Write the oxidation half–reaction.

b. Write the reduction half–reaction.

c. Write the net reaction.

d. Find the cell potential under standard conditions.

e. Is the reaction spontaneous or nonspontaneous? Why?

f. Write the mass action expression for the net reaction that would be used in the Nernst equation.

g. Will the spontaneity increase or decrease if the concentration of chloride ion is increased in the cell? Why?

Answer

a. Co(s) → Co2+(aq) + 2 e–

b. AgCl(s) + e– → Cl–(aq) + Ag(s)

c. 2 AgCl(s) + Co(s) → Co2+(aq) + 2 Cl–(aq) + 2 Ag(s)

d. E° = 0.2223 + (+0.277) = +0.499 V

e. Spontaneous, E° > 0

f. Q = [Co2+][Cl–]2

g. As the chloride concentration increases the reaction will shift to reactants, which is the nonspontaneous direction.