Think of how chemists first understood oxidation in the kitchen and the lab. When magnesium ribbon burns brightly in air, it grabs oxygen from air to become MgO. That addition of oxygen is called oxidation. In the same sense, when H₂S + Cl₂ → 2HCl + S, hydrogen is snatched away from H₂S, so H₂S is oxidised - removal of hydrogen is also oxidation.

On the other side, reduction is just the reverse. Heating mercuric oxide, 2HgO → 2Hg + O₂, oxygen is removed, so HgO is reduced. And when sulphur combines with hydrogen, S + H₂ → H₂S, you are adding hydrogen, so sulphur gets reduced. So classical idea is simple - oxidation is adding electronegative element or removing electropositive one, reduction is just the opposite.

Now let us make this idea sharper. Oxidation number (ON) is not a real charge, it is an imaginary charge we assign by pretending that every covalent bond is completely ionic. The more electronegative atom gets the negative charge. This bookkeeping helps us track who lost electrons and who gained.

The rules are like traffic rules, easy once you remember. Free element is 0 like Na, O₂, P₄. Ion's ON = its charge like Na⁺ is +1, O^2- is -2. Oxygen is -2 generally, but in peroxides like H₂O₂ it is -1, in superoxide KO₂ it is -1/2, and in OF₂ it is +2. Hydrogen is +1 normally but -1 in metal hydrides. Fluorine is always -1. Alkali metals +1, alkaline earth +2. And most important - sum of ON in a neutral molecule is 0, in an ion it equals ion charge.

Just see how useful this is. In KMnO₄, K is +1, O four times -2 = -8, so Mn must be +7 to make sum zero. That huge +7 tells you it can strongly oxidise others!

Let us practice calculating ON like we do hisaab at the shop - quick and accurate. Take K₂Cr₂O₇. K is +1 each, so +2 total, O is -2 each so -14, total so far -12, therefore two Cr must give +12, so each Cr is +6. In H₂SO₄, H₂ is +2, O₄ is -8, so S must be +6.

Ammonium ion NH₄⁺ - H₄ is +4, net charge is +1, so N is -3. Now Fe₃O₄ is tricky. Average ON comes as +8/3, but you should not get confused. Actually it is a mixed oxide FeO·Fe₂O₃, so one Fe is +2 and two Fe are +3. In organic molecules like acetic acid CH₃COOH, the methyl carbon CH₃ is -3 while the carboxyl carbon COOH is +3. So same molecule can have different ON for same element!

Now comes the modern electronic concept, which is used everywhere in Class 11 and 12. Oxidation is Loss of electrons, Reduction is Gain of electrons. Just remember OIL RIG - Oxidation Is Loss, Reduction Is Gain. It is that simple.

Who collects electrons and who donates? The oxidising agent is the taker - like a strict sahukar. It gains electrons, so it itself gets reduced. The reducing agent is the giver - like a generous dukanwala. It loses electrons, so it itself gets oxidised. For example, Zn → Zn²⁺ + 2e⁻ - Zn loses electrons, so it is a reducing agent and gets oxidised. Cu²⁺ + 2e⁻ → Cu - Cu²⁺ gains electrons, so it is oxidising agent and gets reduced.

Redox reactions have families, just like our festivals have categories. The simplest is Combination: A + B → C. Most combinations where two elements form one product are redox. See C + O₂ → CO₂ - carbon goes from 0 to +4, oxidised, oxygen 0 to -2, reduced. So combination of a metal and non-metal is always redox.

Second is Decomposition: AB → A + B. This is opposite of combination - one breaks into two. Example 2HgO → 2Hg + O₂ - Hg goes +2 to 0 (reduced), O -2 to 0 (oxidised). All decompositions of that type are redox.

Third is Displacement: A + BC → AC + B. A stronger reagent kicks out the weaker from its salt, just like a stronger student taking front seat. Zn + CuSO₄ → ZnSO₄ + Cu - Zn is more reactive than Cu, it displaces Cu. Zn 0 to +2 oxidised, Cu +2 to 0 reduced. This power ranking you will later learn as electrode potential series.

This one is very special for exams and looks like magic. Disproportionation is when one and the same element becomes both king and beggar - it is oxidised and reduced at the same time. Reaction starts with 1 form and gives 2 forms.

Classic example: 2H₂O₂ → 2H₂O + O₂. In H₂O₂, oxygen is -1. In H₂O it becomes -2 (gain of electron, reduced) and in O₂ it becomes 0 (loss of electron, oxidised). So same O is both oxidised and reduced. Similarly P₄ + 3OH⁻ + 3H₂O → PH₃ + 3H₂PO₂⁻ - P goes from 0 to -3 in PH₃ and +1 in H₂PO₂⁻. And Cl₂ + 2OH⁻ → Cl⁻ + ClO⁻ + H₂O - Cl 0 to -1 and +1.

Its reverse is comproportionation - where 2 forms become 1. Example 2H₂S (S -2) + SO₂ (S +4) → 3S (0) + 2H₂O. Two different ON of same element meeting to form middle value, just like two brothers compromising.

How do we name compounds when same metal has different oxidation states? IUPAC gave us Stock notation - we write oxidation number in Roman bracket right after element name. Fe(II), Fe(III), Mn(IV), Au(III) are all Stock names. So Fe₂O₃ is iron(III) oxide because each Fe is +3, and Cu₂O is copper(I) oxide because Cu is +1 there.

Older names you still see in shop labels also follow same logic. -ous means lower ON, -ic means higher ON. So ferrous is Fe(II), ferric is Fe(III), stannous is Sn(II), stannic is Sn(IV). So -ous is chota bhai with lower power, -ic is bada bhai with higher power.

In board exams you will surely get a redox equation to balance. Oxidation Number method is like balancing a seesaw - total increase in ON must equal total decrease in ON. The steps are very logical if you practice.

First write the skeleton. Find which atoms change ON and by how much. Suppose Cu goes from 0 to +2 (increase 2) and N from +5 to +4 (decrease 1). So to make increase = decrease, you need 2 Ns for each Cu. Multiply to equalize. Then balance the other atoms - add H₂O to balance O, H⁺ to balance H in acidic medium. If in basic medium, after acidic balancing, add same number of OH⁻ on both sides to neutralize H⁺ and make H⁺ + OH⁻ → H₂O. Final example from NCERT: Cu + 4HNO₃ → Cu(NO₃)₂ + 2NO₂ + 2H₂O.

This is the most systematic method, like dividing match into two teams. In half-reaction method you literally split the full redox into two halves - one oxidation half and one reduction half. Then you balance each half separately and finally join them.

For each half: first balance all atoms except O and H. Then balance O by adding H₂O, balance H by adding H⁺, then balance charge by adding e⁻. Now you have two half equations. Multiply them so electrons cancel (equalise). Add both halves, cancel common terms. Classic acidic example: MnO₄⁻ + 5Fe²⁺ + 8H⁺ → Mn²⁺ + 5Fe³⁺ + 4H₂O. See how clean it is!

In basic medium, trick is simple. First balance as if acidic. Then add same number of OH⁻ as H⁺ present on both sides. On side with H⁺, H⁺ + OH⁻ becomes H₂O. Cancel extra water. That is it - acidic first, then convert to basic.

Why do we study all this bookkeeping? Because it is used daily in titrations and battery factories. Redox is everywhere. KMnO₄ is special - it is self-indicator. Purple MnO₄⁻ becomes colourless Mn²⁺ in acid, so you know end point without extra indicator. K₂Cr₂O₇ needs diphenylamine as indicator. Iodometry uses starch which gives deep blue with I₂ and colourless when reduced - you must have seen in lab.

Now Electrode Potential E° is the power ranking list for oxidising and reducing strength. More positive E° means stronger oxidising agent - it loves to gain electrons. F₂ is +2.87 V is strongest oxidiser. More negative E° means stronger reducing agent - Li is -3.05 V is strongest reducer. Rule: lower E° gives electrons to higher E° spontaneously. So Zn E° = -0.76, Cu = +0.34, so Zn willingly gives to Cu²⁺: Zn + CuSO₄ → ZnSO₄ + Cu is spontaneous. This series decides which displacement will happen, which battery will work.

It is exactly like IPL ranking - highest rank team can pull players from lower rank easily.