Chem 321
D. Miller
1.
(12)
Answer Key
Name
Using activities, calculate the pH of an aqueous solution formed
by the addition of 20.0 mL of 0.100 M HCl to 200. mL of 0.016 M
K2CO3. Assume that the ionic strength of the solution is fixed at
0.10 M and that the activity coefficient for CO 32- is 0.37.
CO32- + H3O+
init mol
0.0032
mol after rxn
0.0012
6 HCO3- + H2O
0.0020
0
0.0020
a HCO3-/CO32- buffer is formed with
-3
Since the 10 rules apply,
for μ = 0.10 M
γCO32- = 0.37
γHCO3- = 0.775
[HCO3-]init = 0.0020 mol/0.220 L = 0.00909 M
[CO32-]init = 0.0012 mol/0.220 L = 0.00545 M
9.79
2.
For each of the following multiple-choice questions, circle the
letter preceding the correct (best) answer. Only one answer will
be accepted for each.
(3 ea)
The ionic strength of 0.010 M Na 3PO4 aqueous solution is:
A.
0.040 M.
B.
0.12 M.
C.
0.020 M.
D.
0.060 M.
Activity coefficients generally increase as:
A.
the ionic strength increases.
B.
the charge on the ion increases.
C.
the hydrated ion size increases.
D.
all of the above.
If, in an aqueous solution,
[Fe 2+]=[Fe3+], one should expect that
2+
the activity of Fe is:
A.
smaller than the activity of Fe 3+.
B.
larger than the activity of Fe 3+.
C.
the same as the activity of Fe 3+.
Chem 321
D. Miller
2.
Name
Answer Key
(continued)
At the half-way point in the titration of NH3 with HCl, the pH of
the solution is about equal to:
A.
pKb of NH3.
B.
pKa of NH4+.
C.
both A and B.
D.
none of the above.
If K3PO4 is titrated with HCl, the pH of the solution at the
second equivalence point is about equal to:
3.
A.
9.6
B.
7.2
C.
4.7
D.
2.2
Short answer
If a
with
DATA
your
sample of formic acid (HCO2H, Ka = 1.80 x 10-4) is titrated
0.10 M NaOH, which of the acid-base indicators listed on the
SHEET would be the most appropriate one to use. Justify
choice of indicator.
(5)
HCO2H pKa = 3.74
At the equivalence point you expect
pH ~ (12.5 + 3.74)/2 ~ 8.1. A suitable indictor
undergoes a color change in a pH range that brackets the
equivalence point pH. Thus, thymol blue, with a pH
transition range of 8.0 - 9.6, is the best choice.
Why is a strong acid used instead of a weak acid to titrate a
weak base? Be specific.
(5)
The reaction between a strong acid and a weak base strongly favors the products (K is
very large). The reaction between a weak acid and a weak base does not favor the
products as much (it may even favor the reactants). Thus, such a reaction is not as
quantitative as is needed for a titration.
In solutions of weak acids the contribution to the H3O+
concentration from the autoionization of water is usually
insignificant. Indicate one situation where this contribution is
likely to be important.
(3)
The water contribution to the [H3O+] is about 10-7 M. This becomes important when (1)
the concentration of acid is very small (e.g., 10-9 M) and (2) the acid hydrolyzes to a very
limited extent (e.g., Ka = 10-11).
Chem 321
D. Miller
3.
Name
Answer Key
(continued)
If you wished to buffer a solution at pH=7.0, which of the
following weak acids would you choose to make your buffer?
Justify your choice.
(4)
H2SO3
HOBr
HC7H5O2
pKa1 = 1.91
pKa2 = 7.18
pKa = 8.64
pKa = 4.20
For an effective buffer, pH ~ pKa ± 1. Thus H2SO3 is the best choice.
Briefly describe how you would prepare a 0.10 M solution of this
buffer at the desired pH.
(3)
Measure out the necessary amount of H2SO3 to produce a 0,10 M solution and dissolve
in deionized water (use ~ 3/4 of the desired final volume), Monitor the pH of this solution
and add NaOH until the pH = 7.00. Dilute to the final volume with deionized water.
Indicate whether an aqueous solution of each of the following is
acidic, basic or neutral. Explain your choice in each case.
(3 ea)
NaHSO3 HSO3- is amphiprotic
ACIDIC
KClO4 -
NEUTRAL, because ClO4- is the conj. base of a strong acid and is
thus a neutral ion as is K+
When 15 mL of 0.10 M NaOH is mixed with 20 mL of 0.050 M H 2CO3,
what type of solution (weak acid, weak base, buffer, etc.)
results? Explain your answer.
(3)
reactions: H2CO3 + OHHCO3- + OH-
6 HCO3- + H2O
6 CO32- + H2O
Since there are 0.0015 mol of OH- and only 0.0010 mol of H2CO3, all of the H2CO3 will be
converted to HCO3- and then some of the HCO3- will be converted to CO32- forming a
HCO3- / CO32- buffer.
(4)
Briefly explain why the solubility of AgSCN in water increases
when KNO3 is added to the solution.
Adding KNO3 to the solution increases the ionic strength. This results in better shielding
of the Ag+ and SCN- ions, making it more difficult for them to recombine and allowing for
more AgSCN to dissolve.
Chem 321
D. Miller
4.
(10)
Name
Answer Key
Using activities, calculate the pH of a 0.010 M Ba(OH) 2 aqueous
solution.
0.010 M Ba(OH)2 ! [OH-] = 0.020 M
μ = ½[(0.010 M)(+2)2 + (0.020 M)(-1)2] = 0.030 M
12.23