Chapter 20
Applications of
Oxidation/Reduction Titrations
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Auxiliary Oxidizing and Reducing Reagents
• The analyte in an oxidation/reduction titration must be in a single
oxidation state at the outset; however, the steps preceding titration, such
as dissolving the sample and separating interferences, convert the analyte
to a mixture of oxidized states.
• For example,
 Iron solution usually contains a mixture of Fe2+ and Fe3+.
 Convert all of the ion to Fe2+ by treating the sample solution with
an Auxiliary Reducing Agent is important .
• If we plan to titrate with a standard reductant, pretreatment with an
auxiliary oxidizing reagent is needed.
• To be useful as a preoxidant or a prereductant, a reagent must react
quantitatively with the analyte and must be easily removable to avoid
interfering in the titration.
2
Auxiliary Reducing Reagents
• A number of metals are good reducing agents and have been used for
the prereduction of analytes.
 zinc, aluminum, cadmium, lead, nickel, copper, silver…
• Sticks or coils of the metal can be immersed directly in the analyte
solution.
• After reduction is judged complete, the solid is removed manually and
rinsed with water.
• The analyte solution must be filtered to remove granular or powdered
forms of the metal.
• After reduction of the analyte, the reducing agent is removed by
filtration of by use of a reductor.
3
A Jones Reductor
• A typical Jones reductor has a diameter of about 2 cm and
holds a 40- to 50-cm column of amalgamated zinc.
• Amalgamation is accomplished by allowing zinc granules
to stand briefly in a solution of mercury(II) chloride, where
the following reaction occurs:
2Zn(s) + Hg+2  Zn+2 + Zn(Hg)(s)
• Zinc amalgam is an effective reducing agent and has
the inhibiting the reduction of H+ by Zn.
• Solution are quite acidic can be passed through a
Jones reductor without significant hydrogen
formation.
4
Two Reductors
•
In a Walden reductor, granular metallic silver held in a narrow glass column is the
reductant.
•
Silver is not a good reducing agent unless chloride or some other ion that forms a
silver salt of low solubility is present.
•
Prereductions with a Walden reductor are generally carried out from hydrochloric
acid solutions of the analyte.
•
The coating of AgCl produced on the metal is removed periodically by dipping a
zinc rod into the solution that covers the packing.
•
The Walden reductor is somewhat more selective than Jones reductor.
5
Auxiliary Oxidizing Reagents
Sodium Bismuthate (NaBiO3) is a powerful oxidizing agent.
• Can convert Mn2+ to MnO4• Oxidations are performed by suspending the bismuthate in the analyte
solution and boiling for a brief period.
• The unused reagent is then removed by filtration.
NaBiO3(s) + 4H+ + 2e-  BiO+ + Na+ + 2H2O
Ammonium Peroxydisulfate (NH4)2S2O8 is a powerful oxidizing agent.
• In acidic solution, it converts Cr3+ to Cr2O72-, Ce3+ to Ce4+, Mn2+ to MnO4• The half reaction is:
S2O8-2 + 2e-  2SO4-2
• The oxidants are catalyzed by traces of Ag+.
• The excess reagent is easily decomposed by a brief period of boiling:
2 S2O82- + 2H2O -> 4SO42- + O2(g) + 4H+
6
Auxiliary Oxidizing Reagents
Sodium Peroxide and Hydrogen Peroxide are also oxidizing agents.
• Peroxide is a convenient oxidizing agent either as the solid Na salt or as a
dilute solution of the acid.
• The half reaction of H2O2 in acidic solution is:
H2O2 + 2H+ + 2e-  2H2O
E0 = 1.78 V
• After oxidation is complete, the solution is freed of excess reagent by
boiling:
H2O2 -> 2H2O + O2(g)
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Applying standard reducing agents
• Standard solution of most reductants tend to react with O2


they are seldom used for the direct titration of oxidizing analytes
Indirect methods are used instead
• The two most common reductants are:


Fe2+ solutions
Thiosulfate ions (S2O32-)
Iron(II) Solutions
• Iron(II) gets rapidly oxidized by air in neutral solutions but oxidation is
inhibited in the presence of acids, with the most stable preparations being
about 0.5 M in H2SO4.
• Oxidizing agents are conveniently determined by treatment of the analyte
solution with a measured excess of standard iron(II) followed by
immediate titration of the excess Fe2+ with a standard solution of K2Cr2O7
or Ce4+.
8
Applying standard reducing agents
Sodium Thiosulfate
• Thiosulfate ion (S2O3-2) is a moderately strong reducing agent that is used to
determine oxidizing agents by an indirect procedure in which iodine is an
intermediate.
• With iodine, thiosulfate ion is oxidized quantitatively to tetrathionate ion
according to the half-reaction:
2S2O3-2  S4O6-2 + 2eOther oxidants can oxidize the tetrathionate ion to sulfate ion.
Determine oxidizing agents:
1. Adding an unmeasured excess of KI to a slightly acidic solution of analyte
2. Reduction of the analyte and produce the equivalent amount of I2
3. I2 will be titrated with Na2S2O3
 Example: determine sodium hypochlorite in bleaches
OCl- + 2I- + 2H+ -> Cl- + I2 + H2O
(unmeasured excess KI)
I2 + 2S2O32- -> S4O62- + 2I9
Applying standard reducing agents
Sodium Thiosulfate
Detecting End Points in Iodine/Thiosulfate Titrations
• Iodine titrations are often performed with starch as an indicator.
• The deep blue color develops in the presence of I2 with b-amylose.
• Red adduct forms when a-amylose with I2, but it is not reversible and is
undesirable.
• The commercial soluble starch contains b-amylose only.
• Aqueous starch decompose fast in the air due to the bacterial action and
it can be inhibited by adding Hg2+ or CHCl3 as bacteriostat.
• Starch decomposes in the solution with high I2 concentration. In titrations
of excess I2 with Na2S2O3, addition of the indicator must be deferred until
most of the I2 has been reduced.
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Applying standard reducing agents
Sodium Thiosulfate
Stability of Sodium Thiosulfate Solutions
• Although Na2S2O3 solutions are resistant to air oxidation, it decompose to
give sulfur and hydrogen sulfite ion:
S2O3-2 + H+  HSO3- + S(s)
• pH, the presence of microorganisms, the concentration of the solution,
the presence of copper(II) ions, and exposure to sunlight affect the
reaction rate.
• The decomposition reaction increases when it becomes acidic.
• The most important cause for the instability of Na2S2O3 is bacteria that
metabolize thiosulfate ion to sulfite and sulfate or elemental sulfur.
• Sterile solution, keep pH between 9 and 10, or the presence of bactericide
such as chloroform or Hg2+ can increase stability.
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Applying standard reducing agents
Sodium Thiosulfate
Standardizing Thiosulfate Solutions
• Potassium iodate (KIO3) is an excellent primary standard for thiosulfate
solutions.
• The KIO3 and the excess of KI in the acidic solution:
IO3- + 5I- + 6H+ -> 3 I2 + 3 H2O
• The liberated I2 is then titrated with Na2S2O3,
I2 + 2S2O32- -> S4O62- + 2I So 1 mol IO3- = 3 mol I2 = 6 mol S2O32• Other primary standards for sodium thiosulfate are potassium
dichromate, potassium bromate, potassium hydrogen iodate, potassium
hexacyanoferrate(III), and metallic copper.
• All these compounds liberate I2 when treated with excess KI.
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13
Applying standard reducing agents
Sodium Thiosulfate
Applications of Sodium Thiosulfate Solutions
• Several substances can be determined by the indirect method
involving titration with Na2S2O3.
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Applying Standard Oxidizing Agents
• Table shows 5 of the
most widely used
volumetric oxidizing
reagents.
• Standard potential
ranges from 0.5 to 1.5V.
• The choice of agent depends on the strength of the analyte as a reducing
agent, the rate of reaction between oxidant and analyte, the stability of
the standard oxidant solutions, the cost, and the availability of a
satisfactory indicator.
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Applying Standard Oxidizing Agents
The Strong Oxidants: Potassium Permanganate and Cerium(IV)
• Solutions of permanganate ion and cerium(IV) ion are strong oxidizing
reagents whose applications closely parallel one another.
• The half-reactions are:
• The formal potential for Ce4+ in 1M perchloric acid and 1M nitric acid are
1.70 V and 1.61 V; however, Ce4+ is not stable in these two solution.
• In less acidic conditions, the product of MnO4- may be Mn3+, Mn4+, or
Mn6+ depending on conditions.
• Solutions of cerium(IV) in sulfuric acid are stable indefinitely, but
permanganate solutions decompose slowly and thus require occasional
restandardization.
• Cerium(IV) solutions in sulfuric acid do not oxidize chloride ion and can
be used to titrate hydrochloric acid solutions of analytes.
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Applying Standard Oxidizing Agents
The Strong Oxidants: Potassium Permanganate and Cerium(IV)
• Permanganate ion cannot be used with hydrochloric acid solutions unless
special precautions are taken to prevent the slow oxidation of chloride ion
that leads to overconsumption of the standard reagent.
• A further advantage of cerium(IV) is that a primary-standard-grade salt of
the reagent is available, thus making possible the direct preparation of
standard solutions.
• Despite the advantages, potassium permanganate is more widely used.
1. Color of MnO4- solution (deep violet), intense enough to be an
indicator in titration. (Mn2+: light pink, MnO42-: dark green, MnO43-:
deep blue)
2. Lower cost of KMnO4 solution.
3. Ce4+ may form precipitates of basic salts in solutions that are less
than 0.1M in strong acid.
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Applying Standard Oxidizing Agents
Detecting the End Points
• Potassium permanganate solution has an intense purple color, which is
sufficient to serve as an indicator for most titrations.
• When low conc. of KMnO4 solution is used, diphenylamine sulfonic acid or
the 1,10-phenanthroline complex of iron(II) can provide a sharper end point.
• The end point is not permanent because excess permanganate ions react
slowly with the relatively large concentration of manganese(II) ions present
at the end point.
2MnO4- + 3Mn2+ + 2H2O <-> 5MnO2(s) + 4H+
• The equilibrium constant is about 1047, indicating that the equilibrium conc.
of MnO4- is very small. But the reaction rate is slow to make the end point
fades gradually over 30 sec.
• Solutions of cerium(IV) are yellow-orange, but the color is not intense
enough to act as an indicator in titrations.
• Several ox/red indicators can be used for titration with Ce4+, such as iron(II)
complex of 1,10-phenanthroline.
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Applying Standard Oxidizing Agents
The Preparation and Stability of Standard Solutions
• Aqueous solutions of MnO4- are not stable due to water oxidation:
4MnO4- + 2H2O <-> 4MnO2(s) + 3O2(g) + 4OH-
• Permanganate solutions are moderately stable because the
decomposition reaction is slow.
• The decomposition is catalyzed by light, heat, acids, bases, Mn2+ and
particularly MnO2.
• MnO2 is a contaminant in the solid KMnO4.
• Removal of MnO2 by filtration and store in dark improves the KMnO4
stability.
• Filter paper cannot be used for filtering because KMnO4- reacts with it
to form MnO2.
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• The most widely used compounds for preparation of Ce4+ solution are listed.
20
Standardizing Permanganate and Ce(IV)
Solutions
• Sodium oxalate is a widely used primary standard.
• In acidic solutions, oxlate ions are converted to the undissociated acid.
2MnO4- + 5H2C2O4 + 6H+ -> 2Mn2+ + 10CO2(g) + 8H2O
• The reaction between permanganate ion and oxalic acid is complex and
proceeds slowly even at elevated temperature unless manganese(II) is
present as a catalyst.
• As the conc. of Mn2+ increases, the reaction proceeds more and more
rapidly as a result of autocatalysis.
• Sodium oxalate is also widely used to standardize Ce(IV) solutions.
2Ce4+ + H2C2O4 -> 2Ce3+ + 2CO2(g) + 2H+
• Cerium(IV) standardizations against sodium oxalate are usually performed
at 50°C in a hydrochloric acid solution containing iodine monochloride as
a catalyst.
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22
2 MnO4- + 5 C2O42- + 16 H+ -> 2 Mn2+ + 10 CO2(g) + 8 H2O
23
Using Permanganate and Ce(IV) Solutions
• The table lists some of the many applications of MnO4- and Ce4+ solutions
to the volumetric determination of inorganic species.
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Potassium Dichromate
• Dichromate ion is reduced to
Cr3+:
green color
Cr2O72- + 14H+ + 6e- <-> 2Cr3+ + 7H2O
Eo = 1.33V
• Dichromate titrations are generally carried out in solutions that are about
1 M in hydrochloric or sulfuric acid and the formal potential will be 1.0 ~
1.1V.
• Potassium dichromate solutions are indefinitely stable, can be boiled
without decomposition, and do not react with hydrochloric acid.
• The disadvantages are its lower electrode potential and the slowness of
its reaction with certain reducing agents.
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Potassium Dichromate
Preparing Dichromate Solutions
• Potassium dichromate can be dried at 150 ~ 200 oC before being
weighted.
• The orange color of a dichromate solution is not intense enough for use in
end-point detection.
• Diphenylamine sulfonic acid is an excellent indicator for titrations with this
reagent. The oxidized form is violet and the reduced form is colorless. So,
the color change observed is from green (Cr3+) to violet (oxidized form).
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Potassium Dichromate
Applying Potassium Dichromate Solutions
• The principal use of dichromate is for the volumetric titration of iron(II) in the
presence of moderate concentrations of hydrochloric acid.
Cr2O72- + 14H+ + 6Fe2+ <-> 2Cr3+ + 7H2O + 6Fe3+
•
Dichromate with Fe2+ has been widely used for INDIRECT determination of
oxidizing agents.
 A measured excess of Fe2+ solution is added to an acidic solution of analyte.
 The excess Fe2+ is back-titrated with standard K2Cr2O7
•
This method has been applied to the determination of nitrate, chlorate,
permanganate, and dichromate ions as well as organic peroxides and other
oxidizing agents.
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Iodine
• Iodine is a weak oxidizing agent used primarily for the determination of
strong reductants.
I3- + 2e- <-> 3IEo = 0.536 V
• Solutions are prepared by dissolving iodine in a concentrated solution of
potassium iodide.
• With smaller electrode potential, iodine solution has
 Relatively limited application
 Advantageous: a degree of selectivity makes possible the
determination of strong reducing agents in the presence of weak
ones
 Advantageous: a sensitive and reversible indicator for titrations.
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Iodine
• Iodine is not very soluble in water (0.001 M).
• Iodine is usually dissolved in moderately concentrated solutions of KI.
I2(s) + I- <-> I3K = 7.1 x 102
• Iodine solution lack stability because…
 Volatility of the solute: losses of iodine from an open vessel occur in a
relatively short time even in the presence of an excess of I Iodine slowly attacks most organic materials => do not use cork or
rubber stoppers to close containers of the iodine solution
 Air oxidation of iodide ion:
4I- + O2(g) + 4H+ <-> 2I2 + 2H2O
Air oxidation is promoted by acids, heat, and light.
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Potassium Bromate
• Primary-standard potassium bromate (KBrO3) is available from
commercial sources and can be used directly to prepare standard
solutions that are stable indefinitely.
 Not frequently used for direct titration
 Convenient and widely used stable source of bromine:
BrO3- + 5Br- + 6H+ -> 3Br2 + 3H2O
standard
solution
excess
 The primary use of standard KBrO3 is for the determination of
organic compounds that react with Br2.
 To determine the excess bromine, an excess of KI is introduced:
2I- + Br2 -> I2 + 2Br-
 The liberated iodine is then titrated with standard sodium
thiosulfate (Na2S2O3).
 Br2 is incorporated into an organic molecule either by substitution or
by addition.
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Potassium Bromate as a Source of Bromine
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35
Potassium Bromate as a Source of Bromine
• Example of substitution reaction: the use of a bromine substitution
reaction to determine 8-hydroxyquinoline
• 8-hydroxyquinoline is an excellent precipitating reagent for cations:
Al3+ + 3HOC9H6N --(pH 4-9)--> Al(OC9H6N)3(s) + 3H+
Al(OC9H6N)3(s) --(hot 4M HCl)--> 3HOC9H6N + Al3+
3HOC9H6N + 6Br2 --> 3HOC9H4NBr2 + 6HBr
stoichiometric relationship: 1 mol Al3+ = 3mol HOC9H6N = 6mol Br2 = 2mol KBrO3
• Example of Addition Reaction:
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Determining Water with the Karl Fischer
Reagent
• It is used for the determination of water in various types of solids and
organic liquids.
• This important titrimetric method is based on an oxidation/reduction
reaction that is relatively specific for water.
• In an aprotic solvent (neither acidic nor basic), the reaction is:
I2 + SO2 + 2H2O  2HI + H2SO4
• The stoichiometry can vary from 2:1 to 1:1 depending on the presence of
acids and bases in the solution.
• Pyridine (C5H5N) was added in an anhydrous methanol as the solvent.
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Determining Water with the Karl Fischer
Reagent
• To stabilize the stoichiometry and shift the equilibrium further to the right,
pyridine (C5H5N) is added and anhydrous methanol is used as the solvent.
• A large excess of pyridine was used to complex I2 and SO2.
C5H5N ·I2 + C5H5N ·SO2 + C5H5N + H2O  2C5H5N ·HI + C5H5N ·SO3
C5H5N+ ·SO3- + CH3OH  C5H5N(H)SO4CH3
• Then the 2nd step is pyridinium sulfite consume water:
C5H5N+ ·SO3- + H2O  C5H5NH+SO4H• The 2nd step is (1) not as specific for water and (2) can be prevent by having a
large excess of methanol.
• The stoichiometry is: 1 mol I2 = 1 mol H2O
• For volumetric analysis, the classical Karl Fischer reagent consistes of I2, SO2,
pyridine and anhydrous MeOH.
38
Determining Water with the Karl Fischer
Reagent
• Pyridine-Free chemistry:
 Pyridine has objectionable odor.
 Other amines such as imidazole have replaced pyridine.
 The reaction is now believed to occur as follows:
1. Solvolysis:
2ROH + SO2  RSO3- + ROH2+
2. Buffering:
B + RSO3- + ROH2+  BH+SO3R- + ROH
3. Redox:
B ·I2 + BH+SO3R- + B + H2O  BH+SO4R- + 2BH+I-
• The stoichiometry: 1 mol I2 = 1 mol H2O
39
Determining Water with the Karl Fischer
Reagent
Interfering reactions
• Oxidizing agents such as Cu(II), Fe(III), nitrite, Br2, Cl2, or quinones produce I2,
which can react with H2O and cause determinations that are too low.
• The carbonyl groups on aldehydes and ketones can react with SO2 and H2O to form
bisulfite complexes.
• Oxidizable species such as ascorbic acid, ammonia, thiols, Tl+, Sn2+, In+, hydroxyl
amines, and thiosulfite can reduce iodine and cause water determinations that are
too high.
• Some interfering compounds react to produce H2O…
 carboxylic acid and alcohol produce ester and water
 Ketone and aldehyde react with alcoholic solvents to form ketals and acetals
R2C=O + 2CH3OH  R2C(OCH3)2 + H2O
• Phenolic derivatives and bicarbonates also cause reduction of I2.
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Determining Water with the Karl Fischer
Reagent
Detecting the End Point
• A Karl Fischer titration can be observed visually based on the brown color of the
excess reagent (C5H5N ·I2) or by electroanalytical measurements.
Reagent Properties
• Karl Fischer reagent decomposes on standing and should be prepared a day or two
before use.
Applications
• The Karl Fischer reagent can be used in the determination of water in many
organic acids, alcohols, esters, ethers, anhydrides, and halides.
• The hydrated salts of most organic acids, as well as the hydrates of a number of
inorganic salts that are soluble in methanol, can also be determined by direct
titration.
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