Electrochemistry

Syllabus points

• Calculate oxidation numbers of elements in compounds and ions.
• Use oxidation numbers to identify oxidation, reduction, oxidising agents and reducing agents.
• Explain and use the term disproportionation.
• Construct redox equations from half-equations, balancing for both atoms and charge.
• Describe electrolysis, and predict the products at each electrode.

Oxidation number

Oxidation number is a bookkeeping device: it records how electrons have been shared out, and it makes redox reactions countable rather than a matter of judgement.

The rules, applied in order:

• An uncombined element has oxidation number 0. That includes O₂, Cl₂ and P₄.
• A simple ion has the oxidation number of its charge, so Na⁺ is +1 and S²⁻ is -2.
• Group 1 is always +1, Group 2 always +2, fluorine always -1.
• Hydrogen is +1, except in metal hydrides such as NaH where it is -1.
• Oxygen is -2, except in peroxides such as H₂O₂ where it is -1, and in OF₂ where it is +2.
• The oxidation numbers in a neutral compound sum to 0; in an ion they sum to the charge.

Worked example

Find the oxidation number of manganese in MnO₄⁻.

Each oxygen is -2, and there are four of them, giving -8. The overall charge is -1, so

Mn + (-8) = -1, therefore Mn = +7

The same method gives +6 for chromium in Cr₂O₇²⁻: fourteen oxygens contribute -14, the ion carries -2, so the two chromiums share +12 and each is +6.

Oxidation, reduction and agents

The mnemonic OIL RIG covers the first column: oxidation is loss, reduction is gain.

The agents are where marks are lost, because the naming feels backwards:

• An oxidising agent oxidises something else, so it is itself reduced and gains electrons.
• A reducing agent reduces something else, so it is itself oxidised and loses electrons.

Common oxidising agents are MnO₄⁻ in acid, Cr₂O₇²⁻ in acid, and the halogens. Common reducing agents are metals, I⁻, Fe²⁺ and SO₂.

Disproportionation

Disproportionation is a reaction in which the same element is both oxidised and reduced.

The classic example is chlorine with cold dilute sodium hydroxide:

Cl₂ + 2NaOH → NaCl + NaClO + H₂O

Chlorine starts at 0. In NaCl it is -1, so it has been reduced; in NaClO it is +1, so it has been oxidised. One element, both directions, in a single reaction.

To identify disproportionation, work out the oxidation number of the element in question on both sides and look for it going up and down from the same starting value.

Building redox equations

Half-equations are combined by making the electrons cancel.

1. Write each half-equation, balancing atoms other than O and H first.
2. Balance oxygen by adding H₂O, and hydrogen by adding H⁺.
3. Balance the charge by adding electrons.
4. Multiply each half-equation so the electrons are equal, then add and cancel.

Worked example

Manganate(VII) oxidising iron(II) in acid.

MnO₄⁻ + 8H⁺ + 5e⁻ → Mn²⁺ + 4H₂O

Fe²⁺ → Fe³⁺ + e⁻

The second must be multiplied by 5 so that five electrons appear on each side:

MnO₄⁻ + 8H⁺ + 5Fe²⁺ → Mn²⁺ + 4H₂O + 5Fe³⁺

Check both: the atoms balance, and the charge is (-1) + (+8) + (+10) = +17 on the left and (+2) + (+15) = +17 on the right. A redox equation that balances for atoms but not for charge is wrong, and checking the charge catches it.

Electrolysis

Electrolysis is the decomposition of a molten or aqueous ionic compound by passing electricity through it. The ions must be free to move, which is why a solid ionic compound cannot be electrolysed.

• The cathode is negative and attracts cations, which gain electrons, so reduction happens there.
• The anode is positive and attracts anions, which lose electrons, so oxidation happens there.

Remembering that reduction is at the cathode and oxidation is at the anode covers most of what is asked.

Molten compounds

With a molten compound only the two ions of the compound are present, so the products are simply the element from each. Molten lead(II) bromide gives lead at the cathode and bromine at the anode:

Pb²⁺ + 2e⁻ → Pb

2Br⁻ → Br₂ + 2e⁻

Aqueous solutions

Water is present as well, so there is a competition, and the general rules are:

• At the cathode, hydrogen is released unless the metal is below hydrogen in reactivity, in which case the metal is deposited. Copper is deposited; sodium is not, so hydrogen comes off instead.
• At the anode, a halide gives the halogen. If no halide is present, oxygen is released from water or hydroxide ions.

Concentration matters too. Concentrated sodium chloride solution gives chlorine at the anode; very dilute solution gives oxygen.

Common mistakes

• Saying an oxidising agent is oxidised. It is reduced, because it takes electrons from something else.
• Giving oxygen -2 in hydrogen peroxide. It is -1 there.
• Balancing a redox equation for atoms only. The charge must balance as well.
• Saying the cathode is positive. In electrolysis it is negative, and reduction happens there.
• Calling any reaction where one thing is oxidised and another reduced disproportionation. It has to be the same element doing both.
• Forgetting that a solid ionic compound does not conduct, so it must be molten or in solution.