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.

• Potassium manganate(VII)
• Sodium thiosulfate
• Potassium iodide
• Potassium dichromate(VI)
• Hydrogen peroxide

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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:

• ethandioate (oxalate) ion (COO^-)₂
• hydrogen peroxide H₂O₂

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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:

2S₂O₃^2- S₄O₆^2- + 2e

It reacts with iodine in the following way:

2S₂O₃^2- + I₂ S₄O₆^2- + 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:

Cr₂O₇^2- + 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

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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:

• Energy change
• pH change
• Spectroscopic methods
• Conductivity
• Potential difference
• Refractive index
• Optical rotation

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.

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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