Introduction to Oxidation Reduction Redox Reactions

Oxidation and Reduction Reactions: Electron Transfer and Redox Processes

Oxidation and reduction reactions — collectively called redox reactions — are among the most fundamental and widespread chemical processes in nature. The classical definition: oxidation is the gain of oxygen (or loss of hydrogen), and reduction is the loss of oxygen (or gain of hydrogen). The modern electronic definition: oxidation is the loss of electrons, and reduction is the gain of electrons. These two processes always occur together — when one substance loses electrons (is oxidised), another substance must gain those electrons (is reduced). This chapter in CBSE Class 10 Science extends the basic understanding of chemical reactions introduced in the earlier chapter and explores redox processes in detail, including their identification, the role of oxidising and reducing agents, and corrosion and rancidity as practical examples.

In a redox reaction, the oxidising agent (oxidant) is the substance that causes oxidation by accepting electrons itself — the oxidising agent is reduced in the process. The reducing agent (reductant) is the substance that causes reduction by donating its electrons — it is oxidised in the process. A simple mnemonic: OIL RIG — Oxidation Is Loss (of electrons), Reduction Is Gain (of electrons). For example, in the reaction Zn + CuSO₄ → ZnSO₄ + Cu, zinc loses two electrons (Zn → Zn²⁺ + 2e⁻) and is oxidised, while Cu²⁺ gains those electrons (Cu²⁺ + 2e⁻ → Cu) and is reduced. Zinc is the reducing agent and Cu²⁺ is the oxidising agent. This particular reaction is also a displacement reaction because zinc (more reactive than copper) displaces copper from its salt solution. The reactivity series determines whether such a displacement occurs — a more reactive metal displaces a less reactive metal from its salt solution, and this is always a redox reaction.

Corrosion is a natural redox process in which metals react with oxygen and moisture in the environment to form their oxides, hydroxides, or other compounds. The rusting of iron is the most familiar example: 4Fe + 3O₂ + 6H₂O → 4Fe(OH)₃, which further dehydrates to form reddish-brown Fe₂O₃·xH₂O (rust). Corrosion is an oxidation reaction that weakens metallic structures, bridges, ships, and buildings. Prevention methods include painting, greasing, galvanising (coating with zinc), forming alloys (stainless steel resists corrosion because chromium forms a protective oxide layer), and cathodic protection (using a sacrificial anode — attaching a more reactive metal like magnesium that corrodes instead). Silver tarnishes due to reaction with hydrogen sulphide in the air: 4Ag + 2H₂S + O₂ → 2Ag₂S + 2H₂O. Copper develops a green patina of basic copper carbonate: 2Cu + H₂O + CO₂ + O₂ → Cu₂CO₃(OH)₂. Rancidity is the oxidation of fats and oils in food, producing a foul odour and unpleasant taste. It occurs when the unsaturated carbon-carbon double bonds in lipids react with oxygen. Rancidity can be prevented by: storing food in airtight containers to exclude oxygen, refrigeration to slow down the reaction rate, adding antioxidants (substances that prevent oxidation, such as vitamin C and BHA/BHT), packaging in a nitrogen atmosphere to displace oxygen (used for potato chips and cooking oils), and adding anti-oxidant preservatives like sodium benzoate.

• Oxidation is loss of electrons (OIL), reduction is gain of electrons (RIG) — they always occur together in redox reactions.
• The oxidising agent accepts electrons (itself gets reduced); the reducing agent donates electrons (itself gets oxidised).
• Corrosion: rusting of iron requires oxygen and water; prevention includes painting, galvanising, alloying, and cathodic protection.
• Rancidity: oxidation of fats and oils in food producing foul smell; prevention includes airtight packaging, refrigeration, antioxidants, and nitrogen atmosphere.
• Reactivity series determines whether a metal will displace another from its salt — this is always a redox process.