Redox Titrations
Chapter 16
Quantitative Chemical Analysis, Daniel C. Harris and Charles A. Lucy, © 2020 W. H. Freeman and Company
1
Chapter Outline
•
•
•
•
•
•
•
Section 16-1 The Shape of a Redox Titration Curve
Section 16-2 Finding the End Point
Section 16-3 Adjustment of the Analyte Oxidation State
Section 16-4 Oxidation with Potassium Permanganate
Section 16-5 Oxidation with Ce4+
Section 16-6 Oxidation with Potassium Dichromate
Section 16-7 Methods Involving Iodine
Quantitative Chemical Analysis, Daniel C. Harris and Charles A. Lucy, © 2020 W. H. Freeman and Company
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Redox Titrations
• Based on an oxidation-reduction reaction between analyte and titrant
• Used to determine analytes in
•
•
•
•
Chemistry (e.g., catalyst studies)
Biology (e.g., metabolism studies)
Environmental science (e.g., toxic heavy metal waste analysis)
Materials science (e.g., oxidation state of chromium in laser crystals)
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Large-Scale Redox Titrations Are Used for
Environmental Remediation
Example 1: Using iron to remediate toxic chromate
Example 2: Using ferrate to remediate toxic hydrogen sulfide
Quantitative Chemical Analysis, Daniel C. Harris and Charles A. Lucy, © 2020 W. H. Freeman and Company
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Table 16-1a Oxidizing agents (oxidants)
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Table 16-1b Reducing agents (reductants)
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Section 16-1
The Shape of a Redox Titration Curve
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Predicting the Shape of a Redox Titration Curve
(1 of 2)
Consider the titration of Fe2+ with
Ce4+ in 1 M HClO4.
How is the titration curve predicted?
Figure 16-2
Figure 16-1
Quantitative Chemical Analysis, Daniel C. Harris and Charles A. Lucy, © 2020 W. H. Freeman and Company
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Predicting the Shape of a Redox Titration Curve
(2 of 2)
A. Identify the titration reaction and the indicator half-reactions.
Figure 16-2
B. Recognize that the initial potential (before any titrant is added)
can be measured, but not accurately predicted.
C. To predict the rest of the titration curve, divide the redox
titration curve calculations into the following regions:
1. Before the equivalence point
2. At the equivalence point
3. After the equivalence point
Quantitative Chemical Analysis, Daniel C. Harris and Charles A. Lucy, © 2020 W. H. Freeman and Company
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Redox Titration Curve: Before Titrant Is Added (1 of 2)
How is the titration curve predicted for the titration of Fe2+ with Ce4+?
Figure 16-1
1. Identify the titration reaction.
Ce4+ + Fe2+ → Ce3+ + Fe3+
2. Identify the reduction half-reactions that
can occur at the Pt electrode, and look up
their standard reduction potentials.
Fe3+ + e−
Fe2+
E° = 0.767 V
Ce4+ + e−
Ce3+ E° = 1.70 V
• The potentials are in 1 M HClO4.
Quantitative Chemical Analysis, Daniel C. Harris and Charles A. Lucy, © 2020 W. H. Freeman and Company
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Box 16-1 Many Redox Reactions Are Atom-Transfer
Reactions
More realistic depiction of this titration reaction as an atom-transfer reaction in
which H serves as an electron carrier
Ce4+ + Fe2+ → Ce3+ + Fe3+
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Redox Titration Curve: Before Titrant Is Added (2 of 2)
Figure 16-1
Figure 16-2
The initial potential of
the analyte solution
(before any titrant is
added) is highly
sensitive to impurities
and cannot ordinarily
be accurately
calculated.
Quantitative Chemical Analysis, Daniel C. Harris and Charles A. Lucy, © 2020 W. H. Freeman and Company
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Redox Titration Curve: Before the Equivalence Point
Ce4+
+
Fe2+
→
Ce3+
+
Fe3+
Figure 16-2
• The titrant Ce4+ is completely consumed and creates an equal
number of moles of Ce3+ and Fe3+.
• The solution contains excess unreacted analyte Fe2+.
• Can easily find the concentrations of Fe2+ and Fe3+.
• Convenient to calculate the cell voltage using the analyte
reduction half-reaction.
Fe3+ + e−
Fe2+
E° = 0.767 V (in 1 M HClO4)
E  E (Fe3 Fe2 )  E (S.C.E.)

 [Fe2 ]  
E  0.767  0.059 16 log  3    0.241
 [Fe ]  

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Redox Titration Curve: Half Equivalence Point
Ce4+ + Fe2+ → Ce3+ + Fe3+
Figure 16-2
• When the volume of titrant added is half of the amount
required to reach the equivalence point (V  1 2 Ve ),[Fe3 ]  [Fe2 ].
• In General, at 1 2 Ve :
E = E°analyte − Ereference
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Redox Titration Curve: At the Equivalence Point (1 of 3)
•
•
•
•
•
Ce4+ + Fe2+ → Ce3+ + Fe3+
Exactly enough Ce4+ has been added to react with all Fe2+.
Virtually all cerium is in the form Ce3+, and virtually all iron is
in the form Fe3+.
Tiny amounts of Ce4+ and Fe2+ are present at equilibrium.
Both half-reactions are in equilibrium at the Pt electrode.
Convenient to use both reactions to describe the cell voltage.
Figure 16-2
 [Fe2 ] 
E   0.767  0.059 16 log  3 
 [Fe ] 
 [Ce3 ] 
E   1.70  0.059 16 log  4 
 [Ce ] 
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Redox Titration Curve: At the Equivalence Point (2 of 3)
Ce4+ + Fe2+ → Ce3+ + Fe3+
Figure 16-2
 [Fe2 ] 
 [Ce3 ] 
2E   0.767  1.70  0.059 16 log  3   0.059 16 log  4 
 [Fe ] 
 [Ce ] 
 [Fe2 ][Ce3 ] 
2E   2.467  0.059 16 log  3
4 
[Fe
][Ce
]

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Redox Titration Curve: At the Equivalence Point (3 of 3)
In general, the voltage is independent of concentrations at
the equivalence point:
E
nEanalyte  nE titrant
ntotal
Figure 16-2
 Eref
where n is the number of electrons transferred in the
analyte or titrant half-reaction and ntotal is the sum of the
number of electrons transferred in the two half-reactions.
Quantitative Chemical Analysis, Daniel C. Harris and Charles A. Lucy, © 2020 W. H. Freeman and Company
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Redox Titration Curve: After the Equivalence Point
Ce4+ + Fe2+ → Ce3+ + Fe3+
Figure 16-2
• All iron has been converted to Fe3+ and the moles of Ce3+ equal the
moles of Fe3+.
• The solution contains a known excess of unreacted titrant Ce4+.
• Can easily find the concentrations of Ce4+ and Ce3+.
• Convenient to calculate the cell voltage using the titrant reduction
half-reaction.
Ce4+ + e−
Ce3+
E° = 1.70 V (in 1 M HClO4)
E  E (Ce4 Ce3 )  E (S. C. E.)

 [Ce3 ]  
E  1.70  0.059 16 log  4    0.241
 [Ce ]  

Quantitative Chemical Analysis, Daniel C. Harris and Charles A. Lucy, © 2020 W. H. Freeman and Company
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Redox Titration Curve: Twice the Equivalence Point
Ce4+ + Fe2+ → Ce3+ + Fe3+
Figure 16-2
• When the volume of titrant added is twice the amount required
to reach the equivalence point (V = 2Ve), [Ce3+] = [Ce4+].
• In General, at 2Ve:
E = E°titrant − Ereference
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Example: Potentiometric Redox Titration (1 of 4)
Suppose that we titrate 100.0 mL of 0.050 0 M Fe2+
with 0.100 M Ce4+, using the cell in Figure 16-1. The
equivalence point occurs when VCe  50.0 mL.
Figure 16-1
4
Calculate the cell voltage at 36.0, 50.0, and 63.0 mL.
Quantitative Chemical Analysis, Daniel C. Harris and Charles A. Lucy, © 2020 W. H. Freeman and Company
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Example: Potentiometric Redox Titration (2 of 4)
Solution: At 36.0 mL: This is 36.0/50.0 of the way to the equivalence point.
Therefore, 36.0/50.0 of the iron is in the form Fe3+ and 14.0/50.0 is in the
form Fe2+. Putting [Fe2 ] /[Fe3 ]  14.0 / 36.0 into Equation 16-6 gives E = 0.550 V.

 [Fe2 ]  
E  0.767  0.059 16 log  3    0.241
 [Fe ]  

Quantitative Chemical Analysis, Daniel C. Harris and Charles A. Lucy, © 2020 W. H. Freeman and Company
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Example: Potentiometric Redox Titration (3 of 4)
Solution: At 50.0 mL: Equation 16-11 tells us that the cell voltage at the
equivalence point is 0.99 V, regardless of the concentrations of reagents for this
particular titration.
E = E+ − E(S.C.E.)
At 63.0 mL: The first 50.0 mL of cerium were converted into Ce3+ . There is an
excess of 13.0 mL of Ce4+, so [Ce3 ] /[Ce4 ]  50.0 / 13.0 in Equation 16-12, and
E = 1.424 V.

 [Ce3 ]  
E  1.70  0.059 16 log  4    0.241
 [Ce ]  

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Example: Potentiometric Redox Titration (4 of 4)
Test Yourself: Find E at VCe4  20.0 and 51.0 mL.
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Titration Curve Symmetry Near Equivalence
Point
Figure 16-2
• When the titration stoichiometry is
1:1 (Figure 16-2), the titration curve
is symmetric about the equivalence
point.
• When the reaction stoichiometry is
not 1:1 (Figure 16-3), the titration
curve is not symmetric about the
equivalence point.
• Example: titration of thallium(I)
with iodate
• 2:1 stoichiometry
Figure 16-3
2Tl  IO3  2Cl  6H  2Tl3  ICl2  3H2O
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Demonstration 16-1 Potentiometric Titration of

2+
Fe with MnO4
MnO4  Fe2  8H  Mn2  5Fe3  4H2O
• Dissolve 0.60 g Fe(NH4)2(SO4)2∙6H2O in 400 mL of 1 M H2SO4. Titrate with 0.02 M
KMnO4 (Ve ≈ 15 mL) using Pt and calomel electrodes with a pH meter as a
potentiometer.
Fe3+ + e−
Fe2+
MnO4  8H  5e
Before the equivalence point:
E° = 0.68 V in 1 M H2 SO4
Mn2  4H2O E  1.507 V
After the equivalence point:
E  E   E (calomel)
E  E   E (calomel)

 [Fe2 ]  
E  0.68  0.059 16 log  3    0.241
 [Fe ]  


 [Mn2 ]  
0.059 16
E  1.507 
log 
 0.241

 8 
5
 [MnO4 ][H ]  

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Demonstration 16-1: Potentiometric Titration of

2+
Fe with MnO4
At the equivalence point:
 [Fe2 ] 
E   0.68  0.059 16 log  3 
 [Fe ] 
Add the two equations to get

 [Mn2 ]  
0.059 16
5E   5 1.507 
log 

 8 
5
 [MnO4 ][H ]  

 [Mn2 ][Fe2 ] 
6E   8.215  0.059 16 log 

3
 8 
[MnO
][Fe
][H
] 
4

At the equivalence point, [Fe3 ]  5[Mn2 ] and[Fe2 ]  5[MnO4 ].
 1 
6E   8.215  0.059 16 log   8 
 [H ] 
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Section 16-2
Finding the End Point
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Redox Indicators (1 of 2)
A redox indicator is a compound that changes colors when going from its oxidized
to reduced state.
In(oxidized) + 𝑛e−
In(reduced)
0.059 16  [In (reduced)] 
E E 
log 

n
[In
(oxi
d
i
ze
d)
]


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Redox Indicators (2 of 2)
 [In (reduced)]  10
 .
The color of In(reduced) is observed when 

 [In (oxidized)]  1
The color of In(oxidized) is observed when  [In (reduced)]   1 .
 [In (oxidized)]  10
The color change will occur over the range E   E 

0.059 16 
 volts.
n

A redox indicator will give a satisfactory visual end point when the difference in
formal potentials of the analyte and titrant is  0.4 V.
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Table 16-2 Redox indicators
Color
Indicator
Oxidized
Reduced
E°
Phenosafranine
Red
Colorless
0.28
Indigo tetrasulfonate
Blue
Colorless
0.36
Methylene blue
Blue
Colorless
0.53
Diphenylamine
Violet
Colorless
0.75
4′-Ethoxy-2,4-diaminoazobenzene
Yellow
Red
0.76
Diphenylamine sulfonic acid
Red-violet
Colorless
0.85
Diphenylbenzidine sulfonic acid
Violet
Colorless
0.87
Tris(2,2′-bipyridine)iron
Pale blue
Red
1.120
Tris(1,10-phenanthroline)iron (ferroin)
Pale blue
Red
1.147
Tris(5-nitro-1,10-phenanthroline)iron
Pale blue
Red-violet
1.25
Tris(2,2′-bipyridine)ruthenium
Pale blue
Yellow
1.29
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Gran Plot
• A Gran plot uses data from well before Ve to locate Ve.
Figure 16-4
• Potentiometric data taken close to Ve are not accurate
because electrodes are slow to equilibrate when one
member of the redox couple is nearly gone.
V  10nE /0.059 16  Ve  10 n(E E )/0.059 16  V  10 n(E E )/0.059 16
• Plot V  10nE /0.059 16 (y) versus V (x).
• Straight line with x-intercept = Ve.
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Starch-Iodine Complex
• Many analytical procedures are based on redox titrations involving iodine.
• Starch is the indicator of choice because the starch-iodine complex is an intense blue color.
• The structure of iodine in amylose remains an open
question.
Figure 16-5
• Theoretical interpretation of the visible absorption
spectrum suggests that iodine in amylose exists as I6
units with an I–I bond length of 0.30 nm.
• Vibrational frequencies of iodine in the Raman
spectrum are the same as those of a crystal whose
structure is know to contain nearly linear, indefinitely
long chains of iodine atoms with I–I bond lengths
near 0.31 nm.
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Section 16-3
Adjustment of the Analyte
Oxidation State
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Preadjustment of Analyte Oxidation State
Sometimes the analyte oxidation state needs to be adjusted prior to titration.
Oxidation state adjustment is especially useful for analytes that contain an element in
multiple oxidation states.
Example: Iron samples often contain some iron in both the +2 and +3 states
For a total iron analysis in such samples, the oxidation state is adjusted so that the iron is
either all Fe2+ (prereduction) or all Fe3+ (preoxidation), prior to analysis with a redox
titration.
•
Preadjustment (preoxidation or prereduction) must be quantitative.
•
Excess preadjustment reagent must be eliminated so that it does not interfere in the
subsequent titration.
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Preoxidation
Oxidants Used for Preoxidation
2
1. Peroxydisulfate, S2O8 (persulfate) in the presence of Ag+
•
Excess oxidant is destroyed by boiling the solution.
2. Silver(I,III) oxide, AgIAgIIIO2, (usually written as AgO)
•
Excess reagent is destroyed by boiling the solution.
3. Solid sodium bismuthate (NiBiO3)
•
Excess oxidant is removed by filtration.
4. Hydrogen peroxide, H2O2, is a good oxidant in basic solution
•
Excess oxidant is removed by boiling.
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Prereduction
Reductants Used for Prereduction
1. Stannous chloride (SnCl2) reduces Fe3+ to Fe2+ in hot HCl. Excess reductant is
destroyed by adding HgCl2.
Sn2+ + 2HgCl2 → Sn4+ + 2Hg2Cl2 +2Cl−
2. Chromous chloride is sometimes used to prereduce analyte to a lower
oxidation state.
•
Excess Cr2+ is oxidized by atmospheric oxygen.
3. Sulfur dioxide and hydrogen sulfide are mild reducing agents expelled by
boiling an acidic solution.
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Prereduction Columns
Figure 16-6
An important prereduction technique uses a packed column with a solid
reducing agent.
Jones reductor: a column packed with zinc coated with a zinc amalgam
• Zinc is a powerful reducing agent.
• Not very selective.
• Mercury is a toxic waste hazard, so its use should be minimized.
Walden reductor: a column filled with solid silver and 1 M HCl
• It is more selective than the Jones reductor.
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Finding Environmentally Friendly Replacements
for Toxic Reductants
Determining NO3 in drinking water (EPA limit [NO3 ]  10 ppm)
Classical Nitrate Assay
• Uses metallic Cd to reduce NO3 to NO2 in a prereduction step.
•
However, Cd is toxic and creates a hazardous waste.
Environmentally Friendly Nitrate Assay
• Replaces Cd with the biological reducing agent NADH.
Nitrate reductase pH 7
NO3  NADH  H 
NO2  NAD  H2O
• A spectrophotometric assay is then used to detect the
nitrite.
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Section 16-4
Oxidation with Potassium
Permanganate
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Potassium Permanganate as an Oxidizing Titrant
•
•
KMnO4 is a strong oxidant with an intense violet color.
KMnO4 is reduced to:
•
Mn2+ (colorless) in strongly acidic solutions
MnO4  8H  5e
•
MnO2 (solid brown) in neutral or basic solutions
MnO4  4H  3e
•
•
Mn2  4H2O E  1.507 V
MnO2 (s)  2H2O E  1.692 V
MnO24 (green manganate ion) in strongly alkaline solutions
MnO24 E  0.56 V
MnO4  e
KMnO4 is not a primary standard. It is standardized by titration of sodium
oxalate.
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Preparation and Standardization of Potassium
Permanganate
•
•
KMnO4 is not a primary standard because traces of MnO2 are present.
Distilled water contains enough organic impurities to reduce some MnO4
to MnO2.
• Dissolve KMnO4 in distilled water, boil it, and filter through a sintered-glass
filter.
• Store in a dark glass bottle.
• Aqueous KMnO4 is unstable by virtue of the reaction
4MnO4  2H2O  4MnO2 (s)  3O2  4OH
which is slowed in the absence of MnO2, Mn2+, heat, light, acids, and bases.
•
KMnO4 is standardized by titration of sodium oxalate (Na2C2O4) or Fe.
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Table 16-3 Analytical applications of permanganate
titrations (1 of 3)
Species analyzed
H2C2O4
Oxidation reaction
Notes
Add 95% of titrant at 25°C, then complete
titration at 55°–60°C.
Titrate in boiling 2 M H2SO4 to remove Br2(g).
H2O2
Titrate in 1 M H2SO4.
HNO2
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Table 16-3 Analytical applications of permanganate
titrations (2 of 3)
Species analyzed
Oxidation reaction
Notes
Titrate in 2 M HCl.
Reduce W(VI) with Pb(Hg) at 50°C and titrate in
1 M HCl.
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43
43
Table 16-3 Analytical applications of permanganate
titrations (3 of 3)
Species analyzed
Oxidation reaction
Notes
(NH4)3PO4 ⋅ 12MoO3 is precipitated and
dissolved in H2SO4. The Mo(VI) is reduced
in a Jones reductor as above and titrated.
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Section 16-5
4+
Oxidation with Ce
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4+
Ce as
an Oxidizing Titrant
• Ce4+ binds anions ClO4 , SO24 , NO3 and Cl−.
• Variation of the Ce4+|Ce3+ formal potential is indicative of these interactions:
Ce4+ + e−
•
•
•
Ce3+
Ce4+ is yellow and Ce3+ is colorless, but the color change is not distinct.
Ferroin and other substituted phenanthroline redox indicators are well-suited
to titrations with Ce4+.
Ce4+ can be used in place of KMnO4 in most procedures.
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Preparation and Standardization of
4+
Ce
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Section 16-6
Oxidation with Potassium
Dichromate
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Potassium Dichromate as an Oxidizing Titrant
•
•
K2Cr2O7 is a primary standard for redox titrations and its solutions are stable.
In acid, orange dichromate ion is a powerful oxidant that is reduced to chromic ion.
Cr2O27  14H  6e
•
•
K2Cr2O7 is not as strong an oxidant as MnO4 or Ce4+.
Dichromate is converted into yellow chromate ion (CrO24 ), whose oxidizing power is nil.
Cr2O27  4H2O  3e
•
•
•
2Cr3  7H2O E  1.36 V
Cr(OH)3 (s,hydrated)  5OH E  0.12 V
Indicators such as diphenylamine sulfonic acid and diphenylbenzidine
sulfonic acid are used to find a dichromate end point.
Primarily used for the determination of Fe2+ or indirectly for species
that oxidize Fe2+ to Fe3+.
Reactions can also be monitored with Pt and calomel electrodes.
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Section 16-7
Methods Involving Iodine
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Iodimetry Versus Iodometry
• In iodimetry, a reducing agent is titrated with iodine (to produce I−).
• In iodometry, an oxidizing agent is added to excess I− to produce iodine, which is
then titrated with standard thiosulfate solution.
• Molecular iodine is only slightly soluble in water, but its solubility is enhanced by
complexation with iodide.
I2 (aq)  I

3
I
• When we speak of using iodine as a titrant, we almost always mean that we
are using a solution of I2 plus excess I−.
Quantitative Chemical Analysis, Daniel C. Harris and Charles A. Lucy, © 2020 W. H. Freeman and Company
51
Starch as an Indicator
• In iodimetry, starch can be added at
the beginning of the titration.
• The first drop of excess I3 after the
Color Plate 12
equivalence point causes the solution to
turn dark blue.
• In iodometry, starch should not be
added until immediately before the
equivalence point.
• Otherwise, some iodine tends to
remain bound to the starch particles.
• Detected visually by fading of the I3 .
Quantitative Chemical Analysis, Daniel C. Harris and Charles A. Lucy, © 2020 W. H. Freeman and Company
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
3
Preparation and Standardization of I Solutions
(1 of 2)

• I3 is prepared by dissolving I2 in excess KI.
• Sublimed I2 is pure enough to be a primary standard, but is seldom used as a
standard because it evaporates while it is being weighed.
• The approximate amount is rapidly weighed, and the solution of I3 is standardized
with Na2S2O3.

• Acidic solutions of I3 are unstable because the excess I− is slowly oxidized by air.
6I  O2  4H  2I3  2H2O
• In neutral solutions, oxidation is insignificant in the absence of heat, light, and metal
ions.
• At pH  11, triiode disproportionates to hypoiodous acid, iodate, and iodide.
Quantitative Chemical Analysis, Daniel C. Harris and Charles A. Lucy, © 2020 W. H. Freeman and Company
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
3
Preparation and Standardization of I Solutions
(2 of 2)

An excellent way to prepare standard I3 is to:
• Add a weighed quantity of the primary standard potassium iodate (KIO3) to a small
excess of KI
• Add excess strong acid to adjust the solution pH to ≈1
Under these conditions I3 is quantitatively formed via the following reaction:
103  8I  6H
3I3  H2O
The triiodide titrant solution prepared in this way should be used immediately,
because I3 gradually oxidizes in the presence of air.
Quantitative Chemical Analysis, Daniel C. Harris and Charles A. Lucy, © 2020 W. H. Freeman and Company
54
Use of Sodium Thiosulfate
•
Sodium thiosulfate is the almost universal titrant for triiodide. In neutral or acidic
solution, triioide oxidizes thiosulfate to tetrathionate:
•
In basic solution, I3 disproportionates to I− and HOI, which can oxidize S2O23 to SO24 .
•
•
The common form of thiosulfate, Na2S2O3∙5H2O is not a primary standard.
Thiosulfate titrant solutions are standardized immediately prior to use by titrating a
fresh solution of triiodide prepared from KIO3 (primary standard) and excess I−.
Quantitative Chemical Analysis, Daniel C. Harris and Charles A. Lucy, © 2020 W. H. Freeman and Company
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Iodimetric Determination of Vitamin C
Vitamin C is titrated with I3 until reaching the intense blue starch-iodine end point.
Oxidation half-reaction: Ascorbic acid + H2O
Reduction half-reaction:
2e  I3
Overall Reaction:
dehydroascorbic acid + 2H+ + 2e−
3I
Quantitative Chemical Analysis, Daniel C. Harris and Charles A. Lucy, © 2020 W. H. Freeman and Company
56
Table 16-4 Titrations with standard triiodide
(iodimetric titrations) (1 of 3)
Species analyzed
Oxidation reaction
Notes
Sn(IV) is reduced to Sn(II) with granular Pb or Ni
in 1 M HCl and titrated in the absence of
oxygen.
Titrate in NaHCO3 solution.
SO2
H2S
Quantitative Chemical Analysis, Daniel C. Harris and Charles A. Lucy, © 2020 W. H. Freeman and Company
57
Table 16-4 Titrations with standard triiodide
(iodimetric titrations) (2 of 3)
Species analyzed
Oxidation reaction
Notes
Cysteine,
glutathione,
thioglycolic acid,
mercaptoethanol
HCN
Titrate in carbonate-bicarbonate buffer, using pxylene as an extraction indicator.
H2C=O
Glucose (and
other reducing
sugars)
Quantitative Chemical Analysis, Daniel C. Harris and Charles A. Lucy, © 2020 W. H. Freeman and Company
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Table 16-4 Titrations with standard triiodide
(iodimetric titrations) (3 of 3)
Species analyzed
Oxidation reaction
Notes
Ascorbic acid
(vitamin C)
H3PO3
Titrate in NaHCO3 solution.
Quantitative Chemical Analysis, Daniel C. Harris and Charles A. Lucy, © 2020 W. H. Freeman and Company
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Iodometric Determination of Chlorine
1. Prereduction step: Quantitatively reduce the chlorine analyte with excess I−.
Cl2  3I  2Cl  I3
2. Titrate the I3 produced in the prereduction step with thiosulfate titrant (reducing).
3. Use the stoichiometry of the prereduction and titration reactions to determine
the amount of chlorine analyte from the titration equivalence point volume.
Quantitative Chemical Analysis, Daniel C. Harris and Charles A. Lucy, © 2020 W. H. Freeman and Company
60

3
Table 16-5 Titration of I produced by analyte
(iodometric titrations) (1 of 3)
Species
analyzed Reaction
Notes
Cl2
Reaction in dilute acid.
Reaction in 0.5 M H2SO4.
Br2
Reaction in dilute acid.
Reaction in 0.5 M H2SO4.
Reaction in 0.5 M HCl.
Reaction in 0.5 M HCl.
O2
H2O2
Reaction in 1 M H2SO4.
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
3
Table 16-5 Titration of I produced by analyte
(iodometric titrations) (2 of 3)
Species
analyzed Reaction
Notes
O3 is passed through neutral 2 wt% KI solution.
Add H2SO4 and titrate.
Reaction in 5 M HCl.
Reaction in neutral solution. Then acidify and
titrate.
NH4HF2 is used as a buffer.
Reaction in 1 M HCl.
Reaction in 0.1 M HCl.
Quantitative Chemical Analysis, Daniel C. Harris and Charles A. Lucy, © 2020 W. H. Freeman and Company
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
3
Table 16-5 Titration of I produced by analyte
(iodometric titrations) (3 of 3)
Species
analyzed Reaction
Notes
MnO2
Reaction in 0.5 M H3PO4 or HCl.
Reaction in 0.4 M HCl requires 5 min for
completion and is particularly sensitive to
air oxidation.
Reaction in 1 M H2SO4.
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