Board exam focus — Redox Reactions (CBSE & HBSE)

Redox reactions form the backbone of electrochemistry, metallurgy and analytical chemistry. This chapter develops the concept of oxidation and reduction from the classical (oxygen/hydrogen) view to the modern electron-transfer and oxidation-number views, classifies redox reactions, teaches two balancing methods, and applies the ideas to redox titrations. Both CBSE and HBSE Class 11 syllabi weight oxidation-number assignment, balancing, and disproportionation heavily.

Concepts of Oxidation, Reduction and Oxidation Number

Classical vs Electronic Concept

The classical concept defined oxidation as addition of oxygen (or any electronegative element) or removal of hydrogen (or any electropositive element), and reduction as the reverse. This view is limited because many redox reactions involve neither oxygen nor hydrogen.

The electronic concept is general:

• Oxidation = loss of electrons (LEO)
• Reduction = gain of electrons (GER)
• Mnemonic: OIL RIG — Oxidation Is Loss, Reduction Is Gain.

Oxidation and reduction always occur together; the species oxidised is the reducing agent and the species reduced is the oxidising agent.

Tip: A common CBSE trap is asking which species is the reducing agent. The reducing agent is itself oxidised (loses electrons) — students often pick the opposite.

Oxidation Number (Oxidation State)

The oxidation number is the charge an atom appears to have when all bonds are treated as fully ionic. It can be zero, positive, negative, or even fractional (an average).

Worked logic

In KMnO\u2084: K = +1, O = \u22122 (\u00d74 = \u22128). Let Mn = x. Then +1 + x \u2212 8 = 0, so x = +7.

In S\u2082O\u2083\u00b2\u207b (thiosulphate) the average oxidation number of S is +2, but the two sulphur atoms are actually inequivalent (one is +5, one is \u22121). This is why fractional/average values appear.

Types of Redox Reactions and Disproportionation

Four Major Types

1. Combination reactions — two species combine; redox if elements are involved.

• C(s) + O\u2082(g) \u2192 CO\u2082(g)

1. Decomposition reactions — a compound breaks into two or more; must give at least one element to be redox.

• 2H\u2082O(l) \u2192 2H\u2082(g) + O\u2082(g)

1. Displacement reactions — one element displaces another.

• Metal displacement: CuSO\u2084 + Zn \u2192 ZnSO\u2084 + Cu
• Non-metal displacement: 2Br\u207b + Cl\u2082 \u2192 Br\u2082 + 2Cl\u207b

1. Disproportionation reactions — the same element in a single oxidation state is simultaneously oxidised and reduced.

Disproportionation in Detail

A disproportionation reaction requires an element in an intermediate oxidation state that has both a higher and a lower accessible state.

Tip: Fluorine cannot disproportionate because \u22121 is its lowest state — it has no lower state to fall to. This is a favourite HBSE one-mark question.

Comproportionation

The reverse process, where two different oxidation states of an element form a single intermediate state, is comproportionation, e.g. 5Cl\u207b + ClO\u2083\u207b + 6H\u207a \u2192 3Cl\u2082 + 3H\u2082O.

Balancing Redox Equations and Redox Titrations

Method 1: Oxidation Number Method

1. Write the skeletal equation and assign oxidation numbers.
2. Identify the atoms whose oxidation number changes; compute the increase and decrease per atom.
3. Multiply each by suitable integers so that total increase = total decrease.
4. Balance the remaining atoms (O by H\u2082O, H by H\u207a in acidic medium), then check charge.

Method 2: Half-Reaction (Ion-Electron) Method

1. Split the reaction into oxidation and reduction half-reactions (ionic form).
2. Balance atoms other than O and H.
3. Balance O by adding H\u2082O; balance H by adding H\u207a (acidic) or by adding H\u2082O/OH\u207b (basic).
4. Balance charge by adding electrons.
5. Equalise electrons, add the halves, cancel common species.

Tip: In basic medium, first balance as if acidic, then add equal OH\u207b to both sides to neutralise the H\u207a, combining H\u207a + OH\u207b into H\u2082O. Forgetting this final step is the most common balancing error.

Redox Titrations

In a redox titration the equivalence point is found from the stoichiometry of electron transfer. KMnO\u2084 is a self-indicator (deep purple \u2192 colourless Mn\u00b2\u207a in acidic medium). In iodometric/iodimetric titrations, starch is the indicator (blue-black with I\u2082).

Frequently asked questions

Are these Redox Reactions notes free?

Yes — the Redox Reactions notes for Chemistry (Class 11) on Siksha Sarovar are completely free to read, with no account required.

Do these notes follow CBSE and HBSE?

Yes. The Redox Reactions notes are NCERT-aligned and include guidance for both CBSE and Haryana Board (HBSE), with important questions and MCQs for revision.

What does the Redox Reactions chapter cover?

Concept explanations, key formulas and definitions, fully solved examples and board-pattern practice questions for Redox Reactions.