Types of Cell and Rusting

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Types of Cell and Rusting

This carbon-zinc dry cell is usually called a battery, however it is in fact a cell, with an e.m.f. of about 1.5V.

They are cheap and portable with a relatively short life, they must be thrown away when used up.

The electrode reactions that are occurring in this cell are as follows:

At the negative electrode:

Zn(s) → Zn2+(aq) + 2e-

At the positive electrode

2NH4+(aq) + 2e- → 2NH3(g) + H2(g)

The H2 produced is removed by the reaction with manganese (IV) oxide:

H2(g) + 2MnO2(s) → MnO3(s) + H2O(l)

Lead-acid batteries are most commonly found in cars. They can be recharged since they are made up of secondary cells. A typical battery is made up of 6 cells, each with an e.m.f of 2V.

Acting as electrodes are plates that separate the cells, immersed in dilute sulphuric acid.

The negative electrodes are plates made from lead, the positive electrodes are plates made from lead coated in lead (IV) oxide.

At the negative electrode:

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At the positive electrode:

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PbSO4 is made by both reactions. When the car moves, an electric current is generated. In other words, the PbSO4 made when current is drawn from the cell is decomposed when the car is running. Hence, the reversible reaction.

Rust is in fact the compound hydrated iron (III) oxide with the general formula, Fe2O3.xH2O.

Rusting occurs when both water and oxygen are present. It is an electrochemical process involving two redox half-equations.

The corrosion starts with the oxidation of iron to Fe2+:

Fe(s) → Fe2+(aq) + 2e-

E = -0.44V

The iron acts as a negative terminal. The released electrons travel through the iron until they make contact with oxygen. This forms the positive electrode.

O2(aq) + 2H2O(l) + 4e- → 4OH-(aq)

E = +1.23V.

The following step is that the Fe2+ and OH- ions combine to form Fe(OH)2(s). This is readily oxidised to iron (III) hydroxide, which then partially dehydrates to give rust, Fe2O3.xH2O.

  1. Coating the Iron: By adding a layer of paint, oil or grease to the iron, you prevent oxygen and water from coming into contact with the iron.
  2. Sacrificial Protection: A most effective way of preventing rust is to coat the iron with zinc, this is called galvanising. It works due to zinc's greater reactivity, i.e. zinc has a greater tendency to form ions, hence any Fe2+ present are reduced to Fe atoms.
  3. Alloying: Iron can be alloyed with nickel, chromium or carbon. The presence of other elements helps prevent Fe2+ forming and electrons been released, hence preventing the rusting process starting.

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