Chapter 4
REDUCTION-OXIDATION TITRATIONS
A reduction-oxidation reaction (Redox reaction) involves the transfer of an electron(s) from
species in the reaction to another. Redox reactions allow titrations of analytes that cannot be
determined by acid-base reactions. These include:
• Metal ions such as Cu2+
• Anions such as nitrite
• Neutral or weakly acidic organic species such as Vitamin C
You should revise the following terms and concepts which were introduced in your
Introductory chemistry classes, as they form the basis of the theory of Redox reactions
Terminology:
• Oxidation
• Reduction
• Oxidant
• Reductant
• Oxidising agent
• Reducing agent
• Oxidation number
• Half equation
• Overall equation
You should be able to determine oxidation numbers and be able to use these to determine if a
species has been oxidised or reduced during a reaction.
Determine the oxidation number of the metal in the following species:
• Mn
• Mn2+
• MnO4• MnO2
You also need to be able to work with half equations by selecting the correct ones, combining
them to form balanced overall equations.
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Chapter 4
Important half-equations for redox titrations (especially the practical work associated with
this module)
Table 4.1 Redox half equations
As2O3 + 5H2O 2H3AsO4 + 4H+ + 4e
C2O42- 2CO2 + 2e
C6H8O6 (ascorbic acid) C6H6O6 + 2H+ + 2e
Ce4+ + e Ce3+
Cr2O72- + 14H+ + 6e 2Cr3+ + 7H2O
Fe2+ Fe3+ + e
H2O2 O2 + 2H+ + 2e
I2 + 2e 2IIO3- + 6H+ + 6e I- + 3H2O
MnO4- + 8H+ + 5e Mn2+ + 4H2O
NO2- + H2O NO3- + 2H+ + 2e
2S2O32- S4O62- + 2e
In Redox titrations a strong oxidising agent such as permanganate is able to oxidise more
species than a weaker oxidising agent such as iodine.
Sample Pre-treatment
Many species that cannot be analysed directly by redox can have their oxidation state changed
leading to a species which can be analysed by a redox titration. Examples include:
• Metal ions that are essentially inert in their cation form, Cr3+ and Mn2+, become
reactive if oxidised to a higher oxidation state in the form of an oxyanion, Cr2O72and MnO4-. Excess oxidant must be removed prior to titration to prevent
interference.
• Oxidants for the above changes include persulfate (S2O82-) and periodate (IO4-).
Persulfate requires silver nitrate as a catalyst and the excess is destroyed by boiling.
Periodate is removed by precipitation.
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Chapter 4
Reduction of the analyte can be done in a number of ways:
• Jones reductor is a column containing an amalgam of zinc and mercury through
which the analyte, such as iron (III) is passed. There are a number of problems
involved in this procedure including the reduction of other species in the sample
and the production of hydrogen gas with acidic solutions.
• Walden reductor in which the zinc amalgam is replaced by silver granules and
dilute HCl. The method reduces fewer analytes but is costly.
• Homogeneous reduction methods use the reducing agent to dissolve the sample.
The excess reducing agent must be destroyed prior to analysis. Tin II chloride is a
common reducing agent, excess is destroyed by addition of mercury II chloride.
OXIDISING TITRANTS
Potassium permanganate
Purple KMnO4- is a strong oxidising agent and is a common redox titrant. It is not a primary
standard even though it is solid, available in AR grade and has a relatively high formula
weight.
When used in acidic conditions it is reduced to colourless Mn2+ according to the equation
MnO4- + 8H+ Mn2+ + 4H2O
When used in neutral or alkaline conditions it forms manganese IV oxide, an insoluble solid
which is a catalyst for the auto-decomposition of permanganate. This poses a problem for the
shelf-life of the standardised permanganate.
MnO4- + 4H+ + 3e MnO2 (s) + 2H2O
And
MnO4- + 2H2O + 3e MnO2(s) + 4OHPreparation of a solution of permanganate involves heating on a steam bath for about an hour
and then filtering the solution before storing in a dark coloured glass storage bottle as light
promotes the decomposition of the solution.
Sulfuric acid (0.5-1.0M) is the acid of choice for a redox reaction with permanganate. Using
hydrochloric acid poses a problem as the permanganate may react with the chloride ion
forming chlorine gas.
Standardisation of permanganate may be against the primary standards, sodium oxalate or
arsenic III oxide (very toxic), although a number of secondary standards could also be used.
The reaction between oxalate and permanganate is initially slow, but is catalysed by the
presence of Mn2+ formed during the reaction.
Use the half-equations of Table 4.1 to write the full redox equation for the reaction
between acidic permanganate and oxalate.
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Chapter 4
The McBride method of standardisation involved titration at 80-90oC. The high temperatures
also cause the oxalate to be decomposed.
The Fowler and Bright method of standardisation recommends rapid addition of 90-99% of
the required volume at room temperature, followed by gentle warming to about 55-60OC
before completing the titration.
The colour differences between the two species MnO4-, purple, and Mn2+, colourless, mean
the reaction can be self-indicating and no further indicator is required.
Analysis of iron using permanganate
Iron III cannot be oxidised by permanganate, however, if the Fe3+ is reduced to Fe2+ oxidation
is then possible
Use table 4.1 to write the overall equation for the reaction between permanganate
and iron II.
For many samples containing Fe such as steel, hydrochloric acid is usually the acid that has
been used for the dissolving process. This can cause problems as chlorine gas and
hypochlorous acid may be formed. Changes to the method allow iron to be determined.
Changes include:
• Diluting the sample
• Adding Mn2+to the reaction
• Titrating cold solutions
• Titrating slowly
• Adding phosphoric acid to form a colourless soluble complex with Fe3+
• In practice Zimmerman-Reinhardt reagent, MnSO4, phosphoric acid and sulfuric
acid is added after preliminary reduction of the iron.
Potassium dichromate
Dichromate is a weaker oxidising agent than permanganate and can therefore oxidise fewer
species. This has the benefit of decreasing unwanted reactions in the titration. Dichromate
will only react with chloride in acidic conditions and at chloride levels in excess of 1-2M.
Dichromate in acidic conditions reacts according to the equation below
Cr2O72- + 14H+ + 6e 2Cr3+ + 7H2O
Yellow/orange
green
Potassium dichromate is a primary standard and can be prepared directly from the solid by
accurate weighing. It can be standardised against a pure iron II salt or metallic iron converted
to iron II.
The strong colour of both the oxidised and reduced form of dichromate mean that an external
indicator is required for a titration.
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Chapter 4
True redox indicators change colour at particular solution voltage (or potential) ie the
indicator is oxidised (or reduced) at a certain potential. Examples include:
• N-phenylanthranilic acid
• Sodium diphenylamine sulfonate (in the presence of phosphoric acid)
Use Table 4.1 to write the overall equation for the reaction between dichromate and
ironII.
Cerium (IV)
Cerium IV, often as the sulfate, has an oxidising strength between permanganate and
dichromate. It is extremely stable in solution (in the presence of 1 M sulfuric acid) and is not
affected by storage or light. It is however, 100 times more expensive than permanganate.
Ce4+ + e Ce3+
Yellow
colourless
The solution needs to be standardised, usually against sodium oxalate or arsenic (III) oxide.
Indicators of choice include ferroin and N-phenylanthranilic acid. Cerium IV works well in
the oxidation of organic species such as alkanols and alkanals.
Use Table 4.2 to write the overall equation for the reaction between cerium IV and iron II.
Iodine
Iodine is a weak oxidising agent, only reacting with species that are easily oxidised. Solution
conditions are normally neutral or acidic since iodine reacts with hydroxide. Where iodine is
used to directly titrate the analyte the reaction is termed iodimetric.
Iodine has a solubility in water of only 0.304 g/100 mL of water. In the presence of an excess
iodide ions soluble triiodide ion I3- is formed.
I2 + I- I3 A 0.05M iodine solution has an iodide concentration of 0.15M. This can be standardised
against arsenic (III) oxide or sodium thiosulfate
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Chapter 4
Iodine is volatile making its concentration very variable with time. In some of your practical
work you will be instructed to work on samples individually rather than three at a time.
Why?
Iodine is red-brown and forms colourless iodide ion in its redox reaction. With experience this
endpoint can be identified without the presence of an indicator. The endpoint is enhanced
with the addition of a starch based indicator. Starch forms an intense blue colour in the
presence of excess iodine. Starch is not a true redox indicator but a specific indicator for
iodine.
Iodine can be used to titrate easily oxidised species without reacting with other species. It is
common for analysis of :
• Sulfur dioxide
• Sulfide
• Ascorbic acid
Back titration may be used if the analyte reacts with oxygen in solution. This occurs in the
case of Vitamin C which decomposes when in solution.
With sulfide and SO2 reaction is best in acidic solution, but this risks the loss of the analyte
as a partly soluble gas (H2S or SO2). With these analytes the titration may be performed in the
reverse ie the sample in the burette and the standard in the flask.
Reducing Titrants
Sodium Thiosulfate (Na2S2O3)
Analytes that exist in a fully oxidised state:
• Copper II
• Periodate (IO4-)
• Hypochlorite (ClO-)
cannot be titrated with oxidising titrants such as permanganate or dichromate. They can be
titrated with reducing agents such as sodium thiosulfate. Many other reducing agents are
affected by oxidation in the air and the standardised figure is only valid for a short period of
time.
Thiosulfate reacts quantitatively only with iodine, however thiosulfate is standardised using
potassium iodate as the primary standard. The iodate does not react with the thiosulfate, but in
the presence of an acidic solution and iodide ions it forms iodine which can be titrated.
Formation of iodine
IO3- + 5I- + 6H+ 3I2 + 3H2O
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Chapter 4
Reaction of iodine with thiosulfate
Using the half equations in Table 4.1 write the balanced equation for the reaction between
iodine and thiosulfate.
What is the reaction ratio between the iodate ion and thiosulfate?
In acidic conditions or in the presence of bacteria thiosulfate can decompose to hydrogen
sulfite and elemental sulfur. Anti-bacterial agents such as trichloromethane or sodium
benzoate should be added to maintain solution stability. The solution is also prone to layering
when stored. The container should therefore be shaken thoroughly before use.
When iodine is being titrated (ie present in significant quantities) an insoluble and irreversible
complex forms with starch resulting in loss of iodine and therefore an artificially low endpoint
value. Starch should therefore not be added to the reaction vessel until almost all the iodine
has been reduced. Alternative indicators such as Vitex do not have this problem.
Iodometry is the process used for the quantitative use of thiosulfate to analyse a range of
oxidised species. The analyte reacts quantitatively with an excess of iodide ion to form iodine,
which can then be titrated with a standardised thiosulfate solution, as shown in Fig 4.1.
Analyte + Iodide
Reduced analyte + Iodine
Iodine titrated with
standardised thiosulfate
Fig 4.1 The process of iodometry
The reaction does not involve a back titration as the iodide (added as KI) is not added
quantitatively, simply in excess. The titration determines the moles of iodine which indicates
the moles of analyte.
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Chapter 4
Write the equations for the reaction of copper II ions with iodide forming a precipitate of
copper II iodide and iodine followed by the reaction of the iodine formed with thiosulfate.
The excess iodide can be a problem since it can be oxidised by the oxygen in air and in
solution to form iodine introducing an error. The error is known as the oxygen error and can
be minimised by avoiding high concentrations of acid and titrating as soon as possible.
O2 + 4I- + 4H+ 2I2 + 2H2O
Typical iodometric analyses include:
• Copper II
• Hydrogen peroxide
• Hypochlorite ion
For copper II analyses the formation of copper II iodide is prone to error as the iodine may
adsorb onto the precipitate. Thiocyanate is added to replace iodide in the precipitate and
prevent loss of iodine.
Iron (II)
The rapid reaction of Fe II with oxygen limits the use of Fe II as a reducing titrant as frequent
standardisation is required. However it is a widely used reducing titrant as it does not produce
toxic by-products and it reacts quantitatively.
Iron II is available in AR grade forms such as ammonium iron II sulfate or can be prepared by
dissolving pure iron acid and reducing the solution in a Jones reductor (Zn and Hg amalgam).
A major application of iron II is with dichromate for the determination of organic content of
water samples. The analysis known as the chemical oxygen demand (COD) involves
refluxing the sample in a excess of standard dichromate and back titrating with Fe II which
has been standardised against the dichromate.
Redox indicators
Table 4.3
Indicator
Titrants used with
Colour change
Starch
Iodine
Colourless to blue
N-phenylanthranilic acid
Cerium IV, dichromate
Purple to colourless
Ferroin
Cerium IV
Blue to red
Diphenylamine sulfonate
Dichromate
Purple to colourless
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Chapter 4
Concepts
1. Why can’t Fe3+ be analysed by titration with permanganate? How could it be treated to
allow titration with permanganate?
2. Explain why a true redox indicator could be used in the titration of iodine with thiosulfate,
but a specific indicator, such as starch could not be used in a titration of iron (II) with
permanganate
3. Give two reasons why an iodometric titration should be performed as soon as possible
after addition of the iodide
Applications
1. Permanganate solutions can be standardised by titrating against a primary standard oxalic
acid. Calculate the molarity of a KMnO4 solution if 0.2351g of oxalic acid, (COOH)2 , is
titrated to endpoint by 18.3 mL of KMnO4
2. An iodine disinfectant solution was analysed by titration with 0.0822 M thiosulfate
solution. The iodine solution was diluted from 10.0 to 100.0 mL and 25.0 mL aliquots
titrated to endpoint by 17.9 mL of the thiosulfate solution. Calculate the %w/v of the
original iodine solution
3. Calculate the SO2 content (mg/L) in 100 mL of a white wine sample which was titrated to
endpoint with 14.9 mL of 0.01505 M I2.
4. A pig iron sample weighing 2.2461g is dissolved, quantitatively converted to the iron (II)
state and the solution accurately made up to 250.0 mL. 25.0 mL aliquots of this are
titrated with an average of 16.4 mL of a 0.0626 M Ce(IV) solution. Calculate the
percentage of Fe in the sample
5. Calculate the molarity of a thiosulfate solution, given the following data. 1.1362g of
potassium iodate is dissolved and made up to 250 mL. 25 mL aliquots of the iodate are
made acidic and reacted with KI. Titration of the liberated iodine required an average of
22.4 mL of thiosulfate solution.
6. The concentration of oxygen in water can be determined by iodometry. Oxygen oxidises
the Mn(OH)2 to Mn(OH)3, and is itself reduced to water. The manganese (III) then reacts
in acidic conditions with iodide to form iodine. This is titrated with thiosulfate. Calculate
the concentration of oxygen in water if a 250 mL sample requires 7.8 mL of 0.0101 M
thiosulfate in the titration.
7. Calculate the cost of titrant in a titration of 25 mL of 0.1 M iron (II) with cerium (IV),
permanganate and dichromate. Costs for 500 g bottles are (NH4)2Ce(NO3)6 $300, K2Cr2O7
$50, KMnO4 $50.
8. A 25 mL aliquot of a nitrite solution was reacted with 50 mL of 0.0203 M KMnO4. The
excess permanganate was back-titrated with 23.6 mL of 0.103 M Fe2+. Calculate the
concentration of the nitrite solution (in g/L).
34