Electrochemistry: Class 9 Chemistry

Key points

Electrochemistry studies the relation between chemical reactions and electricity. Some reactions produce electric current, as in cells and batteries, and electric current can be used to bring about chemical change, as in electrolysis.

Oxidation is the loss of electrons (or an increase in oxidation number) and reduction is the gain of electrons (or a decrease in oxidation number). Remember OIL RIG: oxidation is loss, reduction is gain. In a reaction both happen together, so it is called a redox reaction. The substance that is oxidised is the reducing agent, and the substance that is reduced is the oxidising agent.

Oxygen and hydrogen give older definitions. Oxidation is the gain of oxygen or loss of hydrogen. Reduction is the loss of oxygen or gain of hydrogen. For example, in CuO + H₂ → Cu + H₂O, copper oxide is reduced and hydrogen is oxidised.

Rules for oxidation numbers:

• The oxidation number of an atom in an uncombined element is 0 (Na, O₂, Cl₂).
• For a simple ion it equals the charge (Na⁺ is +1, Cl⁻ is −1).
• In compounds, hydrogen is +1 (except in metal hydrides) and oxygen is −2 (except in peroxides).
• Group 1 metals are +1 and group 2 metals are +2.
• The sum of oxidation numbers equals zero in a neutral compound and equals the charge in an ion.

Example of a redox reaction: Zn + Cu²⁺ → Zn²⁺ + Cu. Zinc goes from 0 to +2, so it is oxidised. Copper goes from +2 to 0, so it is reduced. The half equations are Zn → Zn²⁺ + 2e⁻ and Cu²⁺ + 2e⁻ → Cu.

An electrolyte is a compound that conducts electricity in the molten state or in solution because of free ions. Examples are NaCl, HCl and CuSO₄. Non-electrolytes such as sugar and alcohol do not. Strong electrolytes ionise almost completely, and weak electrolytes (ethanoic acid) ionise only slightly.

Electrolysis is the chemical decomposition of an electrolyte by passing direct current through it. The set-up is an electrolytic cell with two electrodes dipped in the electrolyte and connected to a battery. The anode is the positive electrode and the cathode is the negative electrode. Anions move to the anode and lose electrons (oxidation). Cations move to the cathode and gain electrons (reduction).

Electrolysis of molten sodium chloride gives sodium at the cathode and chlorine at the anode. Cathode: Na⁺ + e⁻ → Na. Anode: 2Cl⁻ → Cl₂ + 2e⁻. This is the industrial method (Downs cell) for sodium. In concentrated aqueous sodium chloride (brine), the products are hydrogen at the cathode, chlorine at the anode, and sodium hydroxide left in the solution.

Electroplating is coating one metal with a thin layer of another using electrolysis. The object to be plated is made the cathode, the plating metal is the anode, and the electrolyte contains a salt of the plating metal. To silver-plate a spoon: the spoon is the cathode, a silver bar is the anode, and the solution is silver nitrate. Plating is done to prevent corrosion, improve appearance and increase hardness.

A galvanic (voltaic) cell converts chemical energy into electrical energy by a spontaneous redox reaction. In a Daniell cell, a zinc rod in zinc sulphate solution and a copper rod in copper sulphate solution are linked by a salt bridge and an external wire. Zinc is the anode (negative), where oxidation takes place: Zn → Zn²⁺ + 2e⁻. Copper is the cathode (positive), where reduction takes place: Cu²⁺ + 2e⁻ → Cu. Electrons flow from zinc to copper through the wire.

The salt bridge, a tube of an inert salt solution such as KNO₃, completes the circuit by allowing ions to move, and it keeps each half cell electrically neutral. The cell potential of the Daniell cell is about 1.10 V under standard conditions. In an electrolytic cell the anode is positive, but in a galvanic cell the anode is negative. Both have oxidation at the anode.

Corrosion is the slow destruction of a metal by reaction with air, water or chemicals. Rusting of iron needs both oxygen and water, and is speeded up by salts. Rust is hydrated iron(III) oxide, Fe₂O₃·xH₂O. It can be prevented by painting, oiling, galvanising (coating with zinc), tin plating and alloying (stainless steel). Sacrificial protection uses a more reactive metal, such as magnesium or zinc, which corrodes in place of the iron.

Batteries are galvanic cells. A dry cell (Leclanché cell) uses a zinc case as the anode and a carbon rod as the cathode with a paste of ammonium chloride. The lead-acid battery used in cars has lead and lead(IV) oxide plates in sulphuric acid, and it can be recharged. Lithium-ion cells are used in phones and laptops.

Study tip: Learn OIL RIG and the rule that oxidation occurs at the anode in every cell. Practise oxidation number calculations by writing the sum equation, and always draw and label the cell diagrams (electrodes, electrolyte, direction of electrons and ions) because these carry marks.

Definitions and flashcards

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Short questions with model answers

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Numericals with worked steps

Common mistakes

• Mixing up anode and cathode charges between galvanic and electrolytic cells. Oxidation is always at the anode.
• Forgetting that the oxidation number of oxygen in H₂O₂ is −1.
• Writing the oxidising agent as the substance that is oxidised. It is the one that is reduced.
• Using the wrong electrode for plating: the object is the cathode.
• Saying that rusting needs only water or only air.

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