IB Chemistry home > Syllabus 2016 > Stoichiometry > Other titrations

Not all titrations involve acids and bases. Any reaction in solution which can be indicated in some form or another can be performed as a titration. The manner of indication could be chemical or physical, depending on the nature of the reaction.

Redox titrations

Redox titrations involve reduction - oxidation reactions, i.e. there is a transfer of electrons from one species to another. Some transition metals change from one oxidation state to another with an accompanying colour change. This means that they are self-indicating and no third substance needs to be added.

In the majority of cases however, some means of indicating the end-point of the titration is needed.


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Potassium manganate(VII) (self-indicating)

Potassium manganate(VII) is a strong oxidising agent. It cannot be obtained sufficiently pure for use as a primary standard; it is usually standardised against a primary standard such as sodium ethandioate (oxalate).

Potassium manganate(VII) is dark purple in colour due to the presence of the manganate(VII) ion, MnO4-. When this ion is reduced by a reducing agent such as iron(II) it turns to manganese(II), Mn2+, ions that are almost colourless (actually very pale pink). This means that the potassium manganate(VII) titration system is self-indicating with a colour change of dark purple to very pale pink.

MnO4- + 8H+ + 5Fe2+ Mn2+ + 5Fe3+ + 4H2O

The iron(II) and iron(II) ions involved in the reaction are also very pale in colour and so do not influence the dramatic colour change of the potassium manganate VII.

The potassium manganate(VII) solution is usually made up in dilute sulfuric acid for two reasons.

  1. It provides the hydrogen ions needed in the redox reaction
  2. It stabilises the solution and helps to prevent decomposition of the potassium manganate(VII) over time.

Example: 25 cm3 samples of an unknown solution of iron(II) sulfate were titrated against a 0.02M solution of potassium manganate(VII) in dilute sulfuric acid until the purple colour just disappeared. The following results were obtained:

titration
initial burette reading ±0.05 cm3
final burette reading ±0.05 cm3
titre (±0.1 cm3)
1
0.0
15.2
15.2
2
0.0
14.8
14.8
3
0.0
14.8
14.8

Calculate the molarity of the iron(II) sulfate solution

Average of concordant results = 14.8 cm3

Moles of manganate(VII) ions = molarity x volume (litres)

Moles manganate(VII) = 0.0148 x 0.02 = 2.96 x 10-4

from equation:

MnO4- + 8H+ + 5Fe2+ Mn2+ + 5Fe3+ + 4H2O

1 mole manganate(VII) = 5 moles iron(II)

moles iron(II) = 5 x 2.96 x 10-4 = 1.48 x 10-3

Volume of iron(II) sulfate solution = 25 cm3

Molarity of iron(II) sulfate solution = moles/volume (litres) = 0.059 M

Potassium manganate(VII) may also be used to determine:


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Sodium thiosulfate

This can't be used as a primary standard, since the sodium thiosulfate pentahydrate crystals have a variable water content. It must be first standardised against a primary standard, such as potassium iodate(V) or potassium manganate(VII).

Sodium thiosulfate is used in the determination of iodine and (indirectly) chlorine and bromine. It can also be used to find concentrations of copper(II) salts.

Sodium thiosulfate is a colourless reducing agent that gets oxidised to the tetrathionate ion:

2S2O32- S4O62- + 2e

It reacts with iodine in the following way:

2S2O32- + I2 S4O62- + 2I-

The indicator used to detect the presence of iodine is starch, which turns a deep blue/black colour in the presence of iodine. The freshly prepared starch solution is not usually added until near the end-point to prevent formation of a permanent blue-black colour.


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Potassium iodide

Iodide ions are reducing agents that can reduce copper(II) ions to copper I ions. This can be used in the determination of copper solutions.

2I- + 2Cu2+ I2 + 2Cu+

An excess of KI solution is added to the unknown copper(II) solution. The iodine liberated is then determined using a standardised sodium thiosulfate solution, which is added slowly until the colour of the iodine changes to pale yellow. Starch solution is then added to intensify the colour due to iodine and the titration continued until the blue-black colour is completely discharged.

Potassium iodide may be used with oxidising agents and the iodine liberated subsequently determined by titration using the sodium thiosulfate/starch method.


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Potassium dichromate(VI)

Potassium dichromate(VI) can be obtained very pure. This means that it is suitable for use as a primary standard and may be used to standardise solutions of most reducing agents.

When it behaves as an oxidising agent, the orange dichromate ions change to green chromium(III) ions. Although the colour change is extreme, it is not suitable to use for end-point determination and an indicator is usually employed in conjuction with the dichromate.

The half-equation for the reaction is:

Cr2O72- + 14H+ + 6e Cr3+ + 7H2O

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Hydrogen peroxide

Hydrogen peroxide H2O2 cannot be obtained easily with a high degree of purity as the solution decomposes slowly at room temperature according to the equation:

H2O2 H2O + O2

Once it has been standardised, however, it can be used in redox titrations with other reducing or oxidising agents.

It gets oxidised (loses electrons) when it reacts with strong oxidising agents such as potassium manganate(VII) and it gets reduced (gains electrons) when it reacts with strong reducing agents.

With a strong oxidising agent the oxygen atoms go from oxidation number -1 to 0:

H2O2 2H+ + O2 + 2e

Wth a strong reducing agent the oxygen atoms change from -1 to -2:

H2O2 + 2e 2OH-

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Precipitation titration

Precipitation titrations involve the addition of one reagent to another producing a precipitation of an insoluble substance. A classical example is the determination of silver ions in a solution.

Silver nitrate solution of known molarity is added dropwise to a specific volume of a standardised solution containing chloride ions. A small amount of chromate ions (usually potassium chromate(VI) solution) having previously been added. Once all of the chloride ions have been precipitated a characteristic brick-red colour appears caused by the formation of a precipitate of silver chromate (Ag2CrO4). The silver chromate cannot form until all of the chloride ions have been used up.

AgNO3 + KCl AgCl + KNO3

Example: A 25cm3 sample of potassium chloride solution was titrated against 0.02M silver nitrate solution. The following results were obtained. Calculate the concentration of the potassium chloride solution.

titration
initial burette reading ±0.05cm3
final burette reading ±0.05cm3
titre (±0.1cm3)
1
0.0
18.8
18.8
2
0.0
18.5
18.5
3
0.0
18.5
18.5

Average titre = 18.5cm3

Moles of AgNO3 = 0.0185 x 0.02 = 3.7 x 10-4

This is equivalent to 3.7 x 10-4 moles of potassium chloride

Moles = molarity/volume(litres)

Molarity of KCl = 3.7 x 10-4 /0.025 = 0.0148 M


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EDTA titrations

These involve the use of EDTA - Ethylene Diamine Tetra Acetic acid (by its old name)

This compound is a hexadentate chelating agent which reacts with metal ions of many different kinds producing colour changes. The actual ions detected depend in the pH of the solutions. EDTA titrations are not required for IB students.


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Physical detection methods

The end point of a chemical reaction does not have to be determined by a colour change. Other factors that could be used identify the end point may be:

Of these methods the only ones that concern IB students are energy change and pH change.

Energy detection

Titrations involving temperature measurements are called thermometric titrations. The energy change is evaluated by measuring the heat of the solution over the course of the titration.

The only advantage thermometric titrations have over conventional titration is that much a higher solution concentration can be used. This is not an advantage 'per se' but it saves the time taken to dilute the more concentrated solutions to that normally used in conventional titrations. Thermometric titrations are, however, not very accurate.

Example: Thermometric titration of ethanoic acid using 2M sodium hydroxide.

5cm3 aliquots of 2M sodium hydroxide are added from a burette to 25cm3 of unknown molar ethanoic acid in a calorimeter and the mixture stirred.

The highest temperature is recorded after each addition.

Additions are continued until no more temperature rise is seen

The results are plotted on a graph of base added against temperature.

Treatment of data

As a hot liquid tends to cool down naturally, the ascending and descending slopes of the graph are back extrapolated and the point where the two lines cross is taken as the highest temperature attained. The volume of base added at this ppint corresponds to the neutralisation point of the ethanoic acid.

pH change

pH change can be recorded using a pH meter either attached to a data logging computer or recorded manually. Direct pH measurement is useful for those titrations where there is more than one acidic proton involved and consequently more than one point of inflexion.

Examples include phosphoric acid, citric acid and ethandioic acid titrations.


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