Redox Reactions

Why Redox Matters for NDA

Redox is short for reduction–oxidation. These two processes always happen together: you can never have one without the other. If one substance loses electrons, another substance must gain those very same electrons. That single idea — a transfer of electrons from one species to another — makes the whole topic logical and scoring.

In NDA General Science (Chemistry), redox questions appear almost every year. They test simple recall: which substance is oxidised, what is the oxidation number of an element, which is the oxidising agent, or what kind of reaction is shown. None of these need long calculations — they only need clear rules, which you will master in this lesson.

Historically, the word oxidation meant “reaction with oxygen”, because that was the first kind chemists studied. Over time the meaning widened to cover loss of hydrogen and finally loss of electrons. Today the electron view is the most general and powerful. Because the rules are mechanical, once you practise a few examples you will rarely lose marks here. Treat this chapter as a banker for your Chemistry score, just like Cavalier toppers do.

What Oxidation and Reduction Mean

There are three classical ways to define oxidation and reduction. Knowing all three helps you answer any phrasing the examiner uses, because the same reaction can be described from any of these angles.

1. In terms of oxygen and hydrogen

• Oxidation = addition of oxygen OR removal of hydrogen.
• Reduction = removal of oxygen OR addition of hydrogen.

Example: 2Mg + O₂ → 2MgO. Magnesium gains oxygen, so magnesium is oxidised. In the reaction CuO + H₂ → Cu + H₂O, copper oxide loses oxygen, so it is reduced, while hydrogen gains oxygen and is oxidised. Notice again that both happen side by side.

2. In terms of electrons

• Oxidation = loss of electrons.
• Reduction = gain of electrons.

3. In terms of oxidation number

• Oxidation = increase in oxidation number.
• Reduction = decrease in oxidation number.

All three definitions describe the same event from different viewpoints. The electron and oxidation-number versions are the most reliable, because they work even when no oxygen or hydrogen is present — for example in 2Na + Cl₂ → 2NaCl, sodium loses an electron (oxidised) and chlorine gains one (reduced), with no oxygen anywhere in sight.

Oxidising and Reducing Agents

In any redox reaction, one species causes oxidation of another and is itself reduced. This is the part students confuse most, so go slow.

• An oxidising agent (oxidant) oxidises the other substance. To do this it accepts electrons, so it is itself reduced.
• A reducing agent (reductant) reduces the other substance. To do this it donates electrons, so it is itself oxidised.

Common oxidising agents: O₂, O₃, Cl₂, F₂, concentrated HNO₃, KMnO₄, K₂Cr₂O₇, H₂O₂.

Common reducing agents: H₂, C, CO, metals like Na, Mg, Zn, Fe, and H₂S.

Rules for Assigning Oxidation Numbers

The oxidation number (or oxidation state) is the charge an atom would have if all the bonds in the molecule were treated as fully ionic, that is, if the more electronegative atom were imagined to take both bonding electrons. It is a bookkeeping tool that lets us track electron movement. Learn these rules in order — they decide most NDA questions and take only minutes to apply.

• Free elements (uncombined) have oxidation number 0. Example: Na, O₂, P₄, H₂ all are 0.
• A monatomic ion has an oxidation number equal to its charge. Na⁺ is +1, Cl⁻ is −1.
• Oxygen is usually −2 (but −1 in peroxides like H₂O₂, and +2 in OF₂).
• Hydrogen is usually +1 (but −1 in metal hydrides like NaH).
• Fluorine is always −1. Group 1 metals are +1; Group 2 metals are +2.
• The sum of oxidation numbers in a neutral molecule is 0; in a polyatomic ion it equals the ion’s charge.

Finding Oxidation Numbers - Step by Step

Let us find the oxidation number of manganese (Mn) in potassium permanganate, KMnO₄ — an NDA classic.

Types of Redox Reactions

NDA papers sometimes ask you to name the type of redox reaction or to spot which one a given equation belongs to. There are four standard categories, and each has a simple identifying clue.

• Combination reaction: two species combine and at least one element changes oxidation state. Example: C + O₂ → CO₂.
• Decomposition reaction: a compound breaks into simpler products. Example: 2H₂O → 2H₂ + O₂.
• Displacement reaction: a more reactive element displaces a less reactive one. Example: Zn + CuSO₄ → ZnSO₄ + Cu.
• Disproportionation reaction: the same element is both oxidised and reduced. Example: Cl₂ + 2NaOH → NaCl + NaOCl + H₂O, where chlorine goes from 0 to both −1 and +1. Another classic is the decomposition of hydrogen peroxide: 2H₂O₂ → 2H₂O + O₂, in which oxygen (−1) is both reduced to −2 and oxidised to 0.

Balancing Redox Equations

The most reliable method for the exam is the oxidation-number method. It works in five clear steps.

1. Write the unbalanced equation and assign oxidation numbers to all atoms.
2. Identify which element is oxidised (number rises) and which is reduced (number falls).
3. Find the total increase and total decrease in oxidation number.
4. Multiply the species by suitable factors so total increase = total decrease (electrons lost = electrons gained).
5. Balance the remaining atoms (O and H) and check charges.

For reactions in acidic solution, add H₂O to balance oxygen and H⁺ to balance hydrogen. In basic solution, add OH⁻ ions instead. The half-reaction method is an alternative: split the reaction into an oxidation half and a reduction half, balance each separately for atoms and charge, then add them so the electrons cancel. Both methods give the same final equation, so use whichever feels faster to you.

Why bother balancing at all? Because chemistry obeys two conservation laws: matter is neither created nor destroyed (mass balance) and electric charge is conserved (charge balance). A correctly balanced redox equation respects both. If your final equation has unequal charges on the two sides, you have made an error somewhere.

Balancing - A Solved Example

Let us balance the reaction between iron(II) and permanganate ions, a textbook favourite.

Redox in Everyday Life

Examiners love linking redox to daily phenomena, because it tests whether you truly understand the idea rather than just memorising it. Keep these examples ready and be able to say what is oxidised and what is reduced in each.

• Corrosion (rusting): iron is oxidised by oxygen and moisture to form hydrated iron oxide (rust), Fe₂O₃·xH₂O.
• Combustion: burning of fuels is rapid oxidation releasing heat and light.
• Respiration: glucose is oxidised in our cells to release energy.
• Photosynthesis: carbon dioxide is reduced to glucose using light energy.
• Batteries and cells: chemical energy is converted to electrical energy through redox at the electrodes.
• Bleaching and disinfection: chlorine and hydrogen peroxide act by oxidation.

Food turning rancid is also a redox process: oils react with atmospheric oxygen and become stale, which is why packets of chips are filled with unreactive nitrogen gas to keep oxygen out. Antioxidants added to food work by getting oxidised themselves, sparing the food. These everyday connections turn dry theory into memorable facts.

Must-Know Facts and Revision

These quick facts pay off directly in the objective paper.

• Oxidation and reduction always occur together — hence “redox”.
• Oxidising agent gets reduced; reducing agent gets oxidised.
• Highest common oxidation states: Mn = +7 (KMnO₄), Cr = +6 (K₂Cr₂O₇), S = +6, N = +5.
• Oxygen is −2 except in peroxides (−1) and OF₂ (+2).
• Hydrogen is +1 except in metal hydrides (−1).
• Fluorine is always −1; it is the strongest oxidising element.

Previous-Year Style Practice

Try this NDA-pattern question before moving on. Cover the answer and attempt it yourself first.