CHEM 321
1
a.
Sample Problems-KEY
Test 2
Select the best answer or fill in the blank.
Which statement is FALSE about an EDTA titration of a metal ion.
The metal-EDTA complex should be less stable than the metal-indicator complex.
The metal-EDTA complex should be more stable than the metal-indicator complex.
Many metals cannot be titrated at high pH because they precipitate as hydroxides
A metal ion will generally give a better titration endpoint if Kf is high
b. Which of the following two component solutions is a good buffer
0.50 M NH3 and 1.00 M NH4Cl
0.05 M acetic acid and 1.00 M sodium acetate
0.5 M HCl and 1.0 M NaCl
0.100 M HPO42- and 0.1 M H3PO4
c. The pH at which the average charge of a polyprotic acid is zero is called the
isoelectric point
equivalence point
endpoint
isoionic point
d. Draw the principal form of glycine (NH2CH2COOH, pKa1 = 2.3 and pKa2 = 9.2) if it is dissolved
in a highly basic solution.
NH2CH2COO
e. The strongest base that can be used in the titration of an acid in CH3NH2 solution CH2NHf.
Which statement is FALSE
In a precipitation titration the endpoint is sharpest when the pK sp of the precipitate is small
An acid-base indicator is a weak acid or base
In all titrations a rapid reaction is desirable
An Volhard titration is a type of precipitation titration
g. Based on the formation constants for EDTA complexes for the following metal ion, which
could be titrated at the lowest pH if the other conditions for all of the titrations are the same.
Mg2+ ( log Kf = 8.79 ),
Sr2+ (log Kf = 8.73).
Ca2+ (log Kf = 10.69 )
h. In this type of titration a known excess of reagent is added to the analyte, and the excess is
then titrated with a second standard reagent.
displacement
blank
direct
back
i. If an acid base indicator has a pKa = 8, then its color change occurs approximately over the
pH range of _7_ to __9_.
j.
What is the pKb for NaF (pKa = 3.17 for HF) ____10.83________
k. Assume that you have solutions and a pH meter. 1.0 M NaH2PO4, 1.0 M Na2HPO4, 1.0 M
HCl, 1.0 M NaOH. How could you make a buffer with pH =7. For H3PO4 (pKa1 = 2.1, pKa2 = 7.2
pKa3 = 12.2)
Add 1.0 M NaOH to NaH2PO4 until pH =7
Add 1.0 M HCl to NaH2PO4 until pH =7
Add 1.0 M NaOH to Na2HPO4 until pH =7
Add 1.0 M HCl to NaOH until pH =7
1
Sample Problems-KEY
CHEM 321
l. Which of the following will produce an acidic solution if dissolved in water?
NaF
NaBr
CH3NH2
Test 2
NH4Cl
m. Compound H4A has the following pKa’s 1,2,5,9. What is its principle form at pH = 6
H4A
H3A
H2A2
HA3
n. What is the approximate pH that results if 10.0 mL of a 0.200 M NaOH solution is added to
20 mL of a 0.200 M NH4Cl solution? ( pKa for NH4+ = 9.25)
4.75
7.00
11.00
9.25
2. How many mL of 0.5 M NaOH must be added a 100 mL of a 1.0 M acetic acid solution in
order to make a buffer with a pH of 5.00. Ka = 1.810 for acetic acid.
First find the ratio required
nbase
nbase
n
 1.78
log base  5.00  4.75  0.25
nacid
nacid
nacid
The moles of weak acid available is found
nacid = (0.100L)(1.0 mol/L) = 0.100 mol
Next recognize that the moles of weak base will be equal to the moles of strong
base added and that the moles of weak acid will be reduced when the strong base is
added.
nNaOH = 0.0640
nNaOH  1.780.100  nNaOH 
nbase
nNaOH

 1.78
2.78nNaOH = 0.178
nacid 0.100  nNaOH
128 mL will be required
n
0.0640 mol
VNaOH  NaOH 
 0.128 L
1
MNaOH 0.50 mol L
pH  pK a  log
3.
Determine the approximate pH of the following solutions or at the point indicated.
(pKa1 = 2.148, pKa2 = 7.199, pKa3 = 12.15 for H3PO4)
a) 0.025 M NaH2PO4
FIRST INTERMEDIATE FROM
pH 
pK1  pK 2 2.148  7.199

 4.67
2
2
b) 0.020 M Na3PO4 (weak base)
Kb 
OH  

Kw
10 14
 12.15  1.41 10 2
K a 10
KbC0 
1.41 10 0.020  0.0168
2
2
Sample Problems-KEY
CHEM 321
Test 2
c) a solution made by adding 40 mL of 0.1 M NaH2PO4 to 80 mL of 0.1 M Na2HPO4
BUFFER

0.040 L 0.100
molNaH2PO4 
  0.0040 molNaH2PO4
L


mol

0.080 L 0.100 Na2HPO4
L

pH  pK a  log

  0.0080 molNa2HPO4

[base]
0.0080
 7.20  log
 7.50
[acid]
0.0040
4.
Draw a rough titration curve for the titration of 50 mL of 0.20 M H2A (pKa1 = 5 and pKa2 = 9)
with 0.10 M NaOH. Mark the equivalence points and the buffer regions. Which indicators are
appropriate for the equivalence points.
  
0 mL
50 mL
100 mL
150 mL
200 mL

weak acid, H  10 5 0.2  1.41 10 3
pH = 2.85
middle of first buffer region, pH = 5
59
7
first equiv point, intermediate form pH 
2
middle of second buffer region, pH = 9


 0.2 
4
second equiv point, weak base, OH  105 
  6.32  10
5


pOH=3.3 , pH = 11.7
indicator for first equiv point is p-nitrophenol
indicator for second equivalence point is Nitramine
5. A solution is 0.200 M in EDTA and 0.200 M in Ca and is buffered at pH = 9. The
formation constant for CaY is 4.9x1010. The fraction of EDTA in the form Y at pH = 9 is
0.054. What is the concentration of free Ca in this solution?
init
equil
Ca2+
0
x
+
EDTA
0
x
CaY 

Ca EDTA
2
K
'
f
2
0.200
x
 8.7  106
9
2.65  10

CaY
0.200
0.200 - x


K'f  Y4 K f  0.054 4.9  1010  2.65  109
K 'f 
0.200  x   0.200  2.65  109
x x 
x2
[Ca2+] = x = 8.7 x 10-6
3
Sample Problems-KEY
CHEM 321
Test 2
6. A 50.00 mL sample containing Fe2+ and Fe3+ required 13.7 mL of 0.0120 M EDTA when
titrated at pH =2 and 29.6 mL when titrated at pH = 6. What are the concentrations of Fe2+ and
Fe3+ in this sample?
CEDTA  0.0120 
VEDTA  13.7  mL
nEDTA  VEDTA  CEDTA
nFe3  nEDTA
nEDTA  1.644  10
nFe3  1.644  10
4
nFe3andFe2  nEDTA
mole
L
nFe2  nFe3andFe2  nFe3
nFe2  1.908  10
CFe2  3.816  10
Vsamp
mol
nEDTA  3.552  10  4 mol
nEDTA  VEDTA  CEDTA
nFe2
4
mol
CEDTA  0.0120 
VEDTA  29.6  mL
CFe2 
Vsamp  50  mL
mole
L
4
mol
 3 mol
CFe3 
L
nFe3
CFe3  3.288  10
Vsamp
 3 mol
L
7. A solution is 0.200 M in EDTA and 0.100 M in Mg and is buffered at pH = 10. What is the
concentration of free Mg in this solution?
init
equil
Mg2+
0
x
+
EDTA
0.100
0.100 + x

MgY 
Mg EDTA
x
0.100
 Mg 2
4.29  107 0.100
2
K 'f 
2

MgY
0.100
0.100 - x

K'f  Y4 K f  0.054 108.9  4.29  107
K 'f 




0.100  x  0.100  4.29  107
x 0.100  x  x 0.100
[Mg2+] = 2.33 x 10-8
4
Sample Problems-KEY
CHEM 321
Test 2
8
A 35.00 mL sample containing Ni2+ is treated with 20.00 mL of 0.0200 M EDTA. All of
2+
the Ni is complexed leaving an excess of EDTA. Titration of the excess EDTA required 9.77
mL of 0.0212 M Mg2+ . What is the concentration of Ni2+ in the original solution?
L  liter
Vs  35  mL
VEDTA  20  mL
CEDTA  0.02 
nEDTA  VEDTA  CEDTA
nMg  VMg  CMg
VMg  9.77  mL
mol
L
CMg  0.0212 
mol
L
nEDTA  4  10  4 mol
nMg  2.071  10  4 mol
nEDTAxs  nMg
nNi  nEDTA  nEDTAxs
CNi 
nNi
Vs
nNi  1.929  10
4
mol
CNi  5.51074  10  3 mol L 1
9. A 0.9541 g sample of flour was analyzed by the Kjedhal method. The ammonia liberated
by the addition of NaOH was distilled into 25.00 mL of 0.06481 M HCl. The excess HCl was
back titrated with 4.12 mL 0f 0.05105 M NaOH. Calculate the percent nitrogen in the flour.
5
Sample Problems-KEY
CHEM 321
mass sample  0.9541  gm
Test 2
L  liter
moles of HCl at start_____________________________________
VHCl  25  mL
mole
L
moles HCl  VHCl  CHCl
CHCl  0.06481 
moles HCl  1.62025  10  3 mole
moles of NaOH titrated___________________________________
mole
VNaOH  4.12  mL
CNaOH  0.05105 
L
moles NaOH  CNaOH  VNaOH
4
moles NaOH  2.10326  10 mole
moles of unreacted HCl___________________________________
moles xsHCl  moles NaOH
moles HClreacted  moles HCl  moles xsHCl
moles xsHCl  2.10326  10  4 mole
moles HClreacted  1.40992  10
moles of NH3 liberated ________________________________
moles NH3  moles HClreacted
moles NH3  1.40992  10  3 mole
moles of N in sample _______________________________
moles N  moles NH3
mass of nitrogen ______________________________________
gm
2
mass N  14.007 
 moles N
mass N  1.97488  10 gm
mole
mass N
%N 
 100
%N  2.06989
mass sample
6
3
mole
Sample Problems-KEY
CHEM 321
Test 2
9. Answer the following questions based on the figure shown. Assume the acid can be
represented by H3A.
Distribution Diagram for a Triprotic System
1.0
0.9
0.8
0.7

0.6
0.5
0.4
0.3
0.2
0.1
0.0
0
1
2
3
4
5
6
7
8
9
10
11
12
13
14
pH
a)
Give approximate values for pKa1
10-3
b)
What are the principal species at pH = 9. Approximately, what is the fraction of
each species that is present? 9% A3-, 82% HA2-, 9% H2Ac)
What is the pH for a buffer made from an equal molar mixture of HA and H2A?
pH = 8
d)
Give the approximate value for Kb2 and write the equation that corresponds to
this constant.
HA2- +H2O  H2A- + OH-
Kb2 = Kw/Ka2 = 10-6
10. Consider the titration curves shown below.
a) Which curve best illustrates pH vs. Vbase for NaH2PO4 titrated with NaOH? B
b) What are the values of the pKa’s associated with the substance titrated in curve B.
3,8
c) Which curve is most likely to represent pMg vs. VEDTA for a titration of Mg with EDTA? C
d) Which curve best illustrates the titration of a weak base (pKb = 4) with strong acid.
7
F
Sample Problems-KEY
CHEM 321
A
Test 2
14
12
14
10
A
B
8
12
10
pX 6
8
pX
4
6
2
4
0
2
0
10
20
30
40
50
0
V
0
50
100
V
150
200
14
16
12
C
D
10
14
8
12
pX
pX 6
10
4
2
8
0
6
0
10
20
30
40
50
V
4
0
20
40
60
80
V
14
12
E
F
10
8
pX 6
4
2
0
0.0
10.0
20.0
30.0
40.0
50.0
V
ACID BASE INDICATORS
INDICATOR
Transition Range(pH)
cresol purple
Congo red
Methyl Red
p-nitrophenol
Cresol Red
Thymolphthalein
Alizaren Yellow
Nitramine
1.2 - 2.8
3.0 - 5.0
4.8 - 6.0
5.6 - 7.6
7.2 - 8.8
8.3 - 10.5
10.1 - 12.0
10.8 - 13.0
8
pK1  pK 2
2
[base]
pH  pK a  log
[acid]
pH 
[H ] 
K1K 2F  K1K w
K1  F
100