Chemical Reactions and Equations

Light the kitchen stove and LPG bursts into a clean blue flame — within a second, new substances are born that were not there before. Chemists everywhere record such changes in one compact, universal line of shorthand called a chemical equation. Writing that line correctly — and balancing it so that not a single atom goes missing — is the first and most scoring skill of Class 10 chemistry.

Definition: A chemical equation is the symbolic representation of a chemical reaction using chemical formulae: the reactants are written on the left and the products on the right, separated by an arrow (→) pointing towards the products.

Definition: A balanced chemical equation contains an equal number of atoms of each element on both sides of the arrow, as demanded by the Law of Conservation of Mass — mass can neither be created nor destroyed in a chemical reaction.

From word equation to skeletal equation

Hold a magnesium ribbon in a flame with a pair of tongs — after cleaning it with sandpaper, which removes the unreactive coating of magnesium oxide/carbonate so that the fresh metal beneath can react. It burns with a dazzling white flame and leaves behind a white powder, magnesium oxide. The change can first be recorded as a word equation:

Magnesium + Oxygen → Magnesium oxide

The substances that react — magnesium and oxygen — are the reactants; the new substance formed — magnesium oxide — is the product. The plus sign joins substances on the same side, and the arrow, read as "gives" or "forms", separates reactants from products and shows the direction of the change.

Word equations are long, and they change from language to language. Chemistry solves this by replacing names with chemical formulae:

Mg + O2 → MgO

An equation written like this, before any atom-checking, is called a skeletal equation. Note that oxygen is written as O2, never as a lone O — hydrogen, oxygen, nitrogen and chlorine all exist as diatomic molecules (H2, O2, N2, Cl2). A formula equation is compact, understood worldwide, and shows exactly which elements take part in the reaction.

The law that forces us to balance

Count the atoms in Mg + O2 → MgO. Left side: 1 Mg and 2 O. Right side: 1 Mg but only 1 O. One oxygen atom has quietly vanished — and that is forbidden. The Law of Conservation of Mass (established by Antoine Lavoisier) states that mass is neither created nor destroyed in a chemical reaction: atoms are only rearranged into new combinations. Two consequences follow directly:

• the number of atoms of each element must be equal on the two sides, and
• the total mass of the reactants must equal the total mass of the products.

An equation that passes the atom count is a balanced chemical equation; the raw skeletal form is unbalanced. Balancing the magnesium equation gives 2Mg + O2 → 2MgO. Verify with masses (Mg = 24 u, O = 16 u): reactants 2 × 24 + 32 = 80 u; products 2 × 40 = 80 u. The chemical account book tallies to the last atom.

Balancing by hit and trial — the five-step drill

NCERT balances equations by the hit-and-trial method: try the smallest whole-number coefficients until every element's count matches. Watch the full method on the burning of hydrogen.

Step 1 — Write the skeletal equation. H2 + O2 → H2O

Step 2 — Count each element on both sides. Left: H = 2, O = 2. Right: H = 2, O = 1. Hydrogen already matches; oxygen does not. (NCERT draws a box around each formula while counting — the box is a reminder that nothing inside it may be altered.)

Step 3 — Start with the compound having the most atoms and fix its unbalanced element. Here that compound is H2O. Make oxygen equal by placing the coefficient 2 before it: H2 + O2 → 2H2O. Remember, a coefficient multiplies the entire formula — 2H2O means 4 H atoms and 2 O atoms.

Step 4 — Rebalance whatever got disturbed; keep lone elements for the end. Hydrogen is now 2 on the left but 4 on the right, so place 2 before H2: 2H2 + O2 → 2H2O. Free elements such as O2, H2 or a lone metal are balanced last, because changing their coefficient disturbs no other element.

Step 5 — Verify, reduce, and polish. H: 4 = 4 ✓; O: 2 = 2 ✓. The coefficients 2 : 1 : 2 are already the smallest whole numbers. With state symbols: 2H2(g) + O2(g) → 2H2O(l).

Golden rule: balance only with coefficients written in front of formulae. Never touch a subscript. Writing H2O2 in place of H2O would "balance" the oxygen instantly — but H2O2 is hydrogen peroxide, a completely different substance, so the equation would then describe a different (and wrong) reaction.

Making the equation more informative — states and conditions

A board answer earns full marks when the balanced equation also communicates the physical states of the substances and the conditions of the reaction.

State symbols, written in brackets after each formula: (s) solid, (l) liquid, (g) gas, and (aq) aqueous solution — a substance dissolved in water.

• Zn(s) + H2SO4(aq) → ZnSO4(aq) + H2(g)
• 2H2(g) + O2(g) → 2H2O(l)
• 3Fe(s) + 4H2O(g) → Fe3O4(s) + 4H2(g) — the water here is steam, so it is H2O(g), not H2O(l).

Reaction conditions — temperature, pressure, catalyst, light — are written above or below the arrow:

• CO(g) + 2H2(g) --(340 atm)--> CH3OH(l)
• 6CO2 + 12H2O --(sunlight, chlorophyll)--> C6H12O6 + 6O2 + 6H2O (photosynthesis)
• Heating is commonly shown by the Greek letter delta (Δ) written over the arrow.

Many examiners also accept an upward arrow (↑) after a gas that escapes and a downward arrow (↓) after a precipitate that settles, in place of (g) and (s).

Five traps that cost marks

1. Changing a subscript to balance. Illegal — it changes the substance itself (turning H2O into H2O2 converts water into peroxide).
2. Forgetting diatomic elements. Reactant oxygen is O2, hydrogen is H2, nitrogen is N2, chlorine is Cl2 — never a lone atom.
3. Counting polyatomic ions atom by atom. When a group such as SO4 or NO3 remains intact on both sides, count the whole group as one unit — faster and far safer.
4. Leaving fractions or non-minimal coefficients. 4H2 + 2O2 → 4H2O is atom-balanced but must be reduced to 2H2 + O2 → 2H2O; a fraction such as 7/2 must be cleared by multiplying the whole equation by 2.
5. Reading too much into the equation. A balanced equation is only an atom account — it does not tell you the speed of the reaction or the conditions required unless these are written on the arrow.

Work through the examples below in order. They begin with straight conversions and finish with the twists CBSE papers love — polyatomic ions, odd oxygen counts, subscript traps and mass audits.

Example 1 — Word equation to chemical equation

Q: Write the balanced chemical equation, with state symbols, for: zinc + dilute sulphuric acid → zinc sulphate + hydrogen.
Given: Formulae — Zn, H2SO4, ZnSO4, H2.
Solve: Skeletal equation: Zn + H2SO4 → ZnSO4 + H2. Count both sides: Zn 1 = 1; H 2 = 2; S 1 = 1; O 4 = 4. Every element already matches — some skeletal equations are born balanced, but you must still verify by counting.
Answer: Zn(s) + H2SO4(aq) → ZnSO4(aq) + H2(g). Quick check: acid and salt are in solution (aq), hydrogen escapes as a gas (g). ✓

Example 2 — Burning magnesium

Q: Balance Mg + O2 → MgO and add the condition and state symbols.
Solve: O: left 2, right 1 → place 2 before MgO: Mg + O2 → 2MgO. Mg: left 1, right 2 → place 2 before Mg: 2Mg + O2 → 2MgO. Verify: Mg 2 = 2, O 2 = 2.
Answer: 2Mg(s) + O2(g) --heat--> 2MgO(s). Sanity check with masses: 48 + 32 = 80 u on both sides. ✓

Example 3 — Combustion of methane (count-table method)

Q: Balance the equation for the burning of methane (natural gas): CH4 + O2 → CO2 + H2O, and add state symbols.
Solve: Count first — C: 1 = 1 ✓; H: 4 vs 2 ✗; O: 2 vs 3 ✗. Balance H by placing 2 before H2O: CH4 + O2 → CO2 + 2H2O (H: 4 = 4). Recount O on the right: 2 + 2 = 4. Balance O last (it stands alone on the left) with 2 before O2: CH4 + 2O2 → CO2 + 2H2O. Verify: C 1 = 1, H 4 = 4, O 4 = 4.
Answer: CH4(g) + 2O2(g) → CO2(g) + 2H2O(g). In a flame the water leaves as vapour, hence H2O(g). ✓

Example 4 — NCERT favourite: iron and steam

Q: Balance Fe + H2O → Fe3O4 + H2 (iron reacting with steam) and add state symbols.
Solve: The biggest compound is Fe3O4 — start with its oxygen. O: left 1, right 4 → 4H2O: Fe + 4H2O → Fe3O4 + H2. H: left 8, right 2 → 4H2: Fe + 4H2O → Fe3O4 + 4H2. Fe: left 1, right 3 → 3Fe. Verify: Fe 3 = 3, H 8 = 8, O 4 = 4.
Answer: 3Fe(s) + 4H2O(g) → Fe3O4(s) + 4H2(g). Trap alert: the water is steam, so its state symbol is (g), not (l). ✓

Example 5 — Tricky: polyatomic ions as single units

Q: Balance BaCl2 + Al2(SO4)3 → AlCl3 + BaSO4, the reaction in which solutions of barium chloride and aluminium sulphate give a white precipitate of barium sulphate.
Solve: The sulphate group SO4 stays intact on both sides — count it as one unit. SO4: left 3, right 1 → 3BaSO4. Ba: right 3, left 1 → 3BaCl2. Al: left 2, right 1 → 2AlCl3. Cl: left 3 × 2 = 6, right 2 × 3 = 6 — it balances automatically, a good sign the earlier steps were correct.
Answer: 3BaCl2(aq) + Al2(SO4)3(aq) → 2AlCl3(aq) + 3BaSO4(s). Verify: Ba 3 = 3, Cl 6 = 6, Al 2 = 2, SO4 3 = 3. ✓

Example 6 — Tricky: the odd-oxygen fraction trick

Q: Balance the combustion of ethane: C2H6 + O2 → CO2 + H2O.
Solve: C: 2 → 2CO2. H: 6 → 3H2O. Oxygen needed on the right = 2 × 2 + 3 × 1 = 7 — an odd number, but O2 supplies oxygen only in pairs. Temporarily allow a fraction: C2H6 + 7/2 O2 → 2CO2 + 3H2O. Fractions are not allowed in a final equation, so multiply every coefficient by 2.
Answer: 2C2H6 + 7O2 → 4CO2 + 6H2O. Verify: C 4 = 4, H 12 = 12, O 14 = 14. ✓

Example 7 — Tricky: why changing a subscript is illegal

Q: To balance H2 + O2 → H2O, student A writes H2 + O2 → H2O2 and student B writes 2H2 + O2 → 2H2O. Both versions have equal atoms on the two sides. Who is correct, and why?
Solve: H2O2 is hydrogen peroxide — a different compound with different properties, not water. By changing the subscript, student A changed the identity of the product, so the equation no longer describes the burning of hydrogen at all. Subscripts are fixed by a compound's composition; balancing may adjust only the coefficients placed in front of formulae.
Answer: Student B is correct: 2H2(g) + O2(g) → 2H2O(l). Rule of thumb: coefficients — yes; subscripts — never.

Example 8 — Tricky: mass audit using conservation of mass

Q: 3.0 g of magnesium ribbon is burnt completely in exactly 2.0 g of oxygen. What mass of magnesium oxide forms? Name the law used, and show that the balanced equation supports the numbers. (Mg = 24 u, O = 16 u)
Formula: Total mass of reactants = total mass of products.
Solve: By conservation of mass, mass of MgO = 3.0 + 2.0 = 5.0 g. Cross-check with 2Mg + O2 → 2MgO: the mass ratio is 48 : 32 : 80, which simplifies to 3 : 2 : 5 — so 3 g of Mg needs exactly 2 g of O2 and yields 5 g of MgO. The data sits perfectly on the balanced ratio.
Answer: 5.0 g of MgO, by the Law of Conservation of Mass. Sanity check: 48 + 32 = 80 u on each side of the balanced equation. ✓

:::keypoints

• A chemical equation shows reactants → products using chemical formulae; the arrow means "gives".
• Balanced equation: atoms of each element are equal on both sides — Law of Conservation of Mass.
• Balance by hit and trial: adjust coefficients only; never change a subscript.
• Start with the compound having the most atoms; balance lone elements (O2, H2) last.
• A coefficient multiplies the whole formula: 2H2O = 4 H atoms + 2 O atoms.
• State symbols: (s) solid, (l) liquid, (g) gas, (aq) aqueous; steam is H2O(g).
• Conditions (heat Δ, pressure, catalyst, sunlight) are written above or below the arrow.
• Final coefficients must be the smallest whole numbers — clear fractions by multiplying through.
  :::

:::memory
Balancing is weighing — think SCALES: Skeletal equation first, Count atoms both sides, Adjust coefficients only, Lone elements last, Equal? verify every element, States and conditions to finish.
:::

:::recap

• A chemical equation is chemistry's balance sheet: atoms in = atoms out.
• The skeletal equation names the substances; balancing makes it obey conservation of mass.
• Only coefficients may be changed — subscripts define the substance itself.
• Finish every equation with state symbols and the reaction conditions on the arrow.
  :::

Watch a wall being whitewashed before a festival: the coating goes on dull, yet two or three days later the wall gleams with a shiny finish. Drop a little water on a lump of quicklime and it hisses, crackles and turns dangerously hot. Dip an iron nail into blue copper sulphate solution and the nail slowly turns brown while the blue itself fades away. Each of these is a chemical reaction — and every reaction you will meet in Class 10 fits into a small set of patterns that this lesson trains you to recognise on sight.

Definition: A combination reaction is a reaction in which two or more reactants combine to form a single product (pattern: A + B → AB). A decomposition reaction is its exact reverse — a single compound breaks down into two or more simpler products (pattern: AB → A + B) using heat, light or electricity.

Definition: In a displacement reaction, a more reactive element displaces (pushes out) a less reactive element from its compound (pattern: A + BC → AC + B). In a double displacement reaction, two compounds exchange their ions (pattern: AB + CD → AD + CB), often producing an insoluble solid called a precipitate.

The four patterns at a glance — count the reactants and products first, then look for free elements.

Combination reactions — many reactants, one product

The give-away of a combination reaction is the product count: however many substances go in, only one product comes out.

The NCERT flagship example is quicklime reacting with water:

CaO(s) + H2O(l) → Ca(OH)2(aq) + heat

Calcium oxide (quicklime) reacts vigorously with water to form calcium hydroxide (slaked lime), releasing such a large amount of heat that the container becomes hot to touch. Two reactants, one product — a combination reaction that is also exothermic.

This reaction has a beautiful follow-up that boards love. A suspension of slaked lime is used for whitewashing walls. The calcium hydroxide then reacts slowly with the carbon dioxide present in air to form a thin layer of calcium carbonate on the walls:

Ca(OH)2(aq) + CO2(g) → CaCO3(s) + H2O(l)

Calcium carbonate is formed two to three days after whitewashing and gives the walls their shiny finish. Interestingly, marble has the same chemical formula, CaCO3.

Other combination reactions to remember:

• Burning of coal: C(s) + O2(g) → CO2(g)
• Formation of water: 2H2(g) + O2(g) → 2H2O(l)
• Burning of magnesium ribbon: 2Mg(s) + O2(g) → 2MgO(s) — it burns with a dazzling white flame, leaving white magnesium oxide powder

A combination can be element + element, element + compound, or compound + compound. The only rule that never bends is the single product.

Decomposition reactions — one compound splits apart

A decomposition reaction is the mirror image of combination: one reactant, two or more products. Breaking chemical bonds needs an energy supply, so NCERT names each kind of decomposition after the energy used — heat, electricity or light.

1. Thermal decomposition (energy from heat). Heat light green ferrous sulphate crystals in a dry boiling tube:

2FeSO4(s) --heat--> Fe2O3(s) + SO2(g) + SO3(g)

The crystals first lose their water of crystallisation and their green colour changes; on stronger heating they decompose to a brown residue of ferric oxide (Fe2O3), while sulphur dioxide and sulphur trioxide escape with the characteristic smell of burning sulphur.

Heating limestone is a thermal decomposition used on an industrial scale (for example in cement manufacture):

CaCO3(s) --heat--> CaO(s) + CO2(g)

Heating lead nitrate gives a dramatic observation:

2Pb(NO3)2(s) --heat--> 2PbO(s) + 4NO2(g) + O2(g)

The brown fumes you see are of nitrogen dioxide (NO2) — a favourite one-mark observation question.

2. Electrolytic decomposition (energy from electricity). Pass an electric current through water (with a few drops of dilute sulphuric acid added to make the water a good conductor):

2H2O(l) --electric current--> 2H2(g) + O2(g)

Hydrogen gas collects at the cathode and oxygen at the anode, and the volume of hydrogen is double the volume of oxygen — exactly the 2 : 1 ratio the balanced equation predicts.

3. Photolytic decomposition (energy from sunlight). Keep white silver chloride in sunlight and it slowly turns grey as silver metal forms:

2AgCl(s) --sunlight--> 2Ag(s) + Cl2(g)

Light yellow silver bromide behaves the same way: 2AgBr(s) --sunlight--> 2Ag(s) + Br2(g). These light-sensitive reactions are used in black and white photography.

Because decomposition reactions must absorb energy to happen, they are typically endothermic.

Displacement reactions — the more reactive element wins

In a displacement reaction a free element attacks a compound and kicks out a less reactive element, taking its place: A + BC → AC + B.

The classic NCERT activity: dip an iron nail in copper sulphate solution and leave it for about 20 minutes.

Fe(s) + CuSO4(aq) → FeSO4(aq) + Cu(s)

Two observations fetch the marks: (i) the nail acquires a brownish coating of copper, and (ii) the blue colour of the solution fades towards light green, because blue copper sulphate is being replaced by light green ferrous sulphate. Iron has displaced copper from its salt, which proves iron is more reactive than copper.

Other standard examples:

• Zn(s) + CuSO4(aq) → ZnSO4(aq) + Cu(s)
• Pb(s) + CuCl2(aq) → PbCl2(aq) + Cu(s)
• Zinc with dilute acid: Zn(s) + 2HCl(aq) → ZnCl2(aq) + H2(g) — here zinc displaces hydrogen, and brisk bubbles of hydrogen gas appear

Zinc and lead are more reactive elements than copper, so both displace copper from its compounds. The reverse never happens: a copper strip left in zinc sulphate or ferrous sulphate solution shows no reaction, because a less reactive element cannot displace a more reactive one. (You will meet the full reactivity series in the chapter on Metals and Non-metals.)

Double displacement reactions — an exchange of ions

Here two compounds react by exchanging their ions: AB + CD → AD + CB. Mix solutions of sodium sulphate and barium chloride:

Na2SO4(aq) + BaCl2(aq) → BaSO4(s) + 2NaCl(aq)

A white precipitate of barium sulphate appears instantly, formed when the barium ions meet the sulphate ions; sodium chloride stays dissolved in the solution. Any reaction in which an insoluble solid (precipitate) forms is called a precipitation reaction, and such precipitation reactions are double displacement reactions.

Common mistake alert: students regularly confuse displacement with double displacement. In displacement, a free element replaces another element in a single compound. In double displacement, two compounds swap their ions and no free element needs to appear at all. Scan the equation: a lone element on the reactant side signals displacement; two compounds facing two compounds signals double displacement.

The energy view — exothermic and endothermic reactions

Cutting across all four patterns is a second, independent classification: does the reaction give out heat or take in energy?

Exothermic reactions release heat along with the products. NCERT's own examples:

• Burning of natural gas: CH4(g) + 2O2(g) → CO2(g) + 2H2O(g) + energy
• Respiration: the glucose from digested food combines with oxygen in the cells of our body and provides energy — C6H12O6(aq) + 6O2(aq) → 6CO2(aq) + 6H2O(l) + energy
• The decomposition of vegetable matter into compost
• Quicklime reacting with water (seen above)

Endothermic reactions absorb energy. Most decomposition reactions are endothermic — that is exactly why they must be fed heat, light or electricity to keep going. Photosynthesis, which absorbs sunlight, is nature's great endothermic reaction.

On an energy diagram, an exothermic reaction ends with products at lower energy than the reactants — the difference escapes as heat. An endothermic reaction ends with products at higher energy — the difference was absorbed from the surroundings.

Exothermic: products settle at lower energy and the difference is released as heat. Endothermic: products sit at higher energy because energy was absorbed.

Exam tip: "exothermic" is a property, not a fifth type. Every reaction is classified twice — once by pattern (combination, decomposition, displacement, double displacement) and once by energy (exothermic or endothermic). Burning coal, for instance, is both a combination reaction and exothermic.

:::compare

Type	General form	NCERT example	Quick clue
Combination	A + B → AB	CaO + H2O → Ca(OH)2	many reactants, ONE product
Decomposition	AB → A + B	CaCO3 → CaO + CO2	ONE reactant, many products; needs heat/light/electricity
Displacement	A + BC → AC + B	Fe + CuSO4 → FeSO4 + Cu	free element replaces an element
Double displacement	AB + CD → AD + CB	Na2SO4 + BaCl2 → BaSO4 + 2NaCl	two compounds swap ions; precipitate common
:::

Work through these examples in order. The early ones train the basic "count and classify" habit; the Tricky ones rehearse the observation-plus-reason questions that board papers love.

Example 1 — Quicklime meets water

Q: What is observed when water is added to a small amount of quicklime? Write the chemical equation and name the type of reaction.
Given: Reactants — calcium oxide (quicklime) and water.
Formula: Combination pattern A + B → AB.
Solve: CaO(s) + H2O(l) → Ca(OH)2(aq) + heat. The mixture hisses and the vessel becomes very hot; a single product, calcium hydroxide (slaked lime), is formed from two reactants.
Answer: It is a combination reaction, and since a large amount of heat is released it is also exothermic. Check: exactly one product on the right — combination confirmed.

Example 2 — Limestone in the kiln

Q: Limestone is heated strongly. Classify the reaction, write the balanced equation, and give one industrial use of the solid product.
Given: Reactant — calcium carbonate (limestone); energy supplied — heat.
Formula: Decomposition pattern AB → A + B.
Solve: CaCO3(s) --heat--> CaO(s) + CO2(g). One compound has split into two simpler products with the help of heat.
Answer: Thermal decomposition (an endothermic change). The product calcium oxide (quicklime) is used in the manufacture of cement. Check: one reactant, two products — decomposition confirmed.

Example 3 — The iron nail in blue solution

Q: An iron nail is kept immersed in copper sulphate solution for about 20 minutes. State two observations, write the equation, and state which of the two metals is more reactive.
Given: Iron nail; blue copper sulphate solution.
Solve: Fe(s) + CuSO4(aq) → FeSO4(aq) + Cu(s). Observation 1: the nail acquires a brownish coating of copper. Observation 2: the blue colour of the solution fades towards light green (ferrous sulphate forms).
Answer: A displacement reaction; iron is more reactive than copper, because it displaces copper from its salt solution. Check: a free element (Fe) stands on the reactant side — displacement confirmed.

Example 4 — Spot the precipitate

Q: Solutions of sodium sulphate and barium chloride are mixed. What is observed? Write the balanced equation and classify the reaction.
Given: Na2SO4 solution and BaCl2 solution.
Formula: Double displacement pattern AB + CD → AD + CB.
Solve: Na2SO4(aq) + BaCl2(aq) → BaSO4(s) + 2NaCl(aq). The barium ions and sulphate ions meet to form insoluble barium sulphate, while sodium chloride remains dissolved.
Answer: A white precipitate of BaSO4 appears — a double displacement reaction, specifically a precipitation reaction. Check: two compounds exchanged partners and an insoluble solid formed.

Example 5 — Tricky: classify all four in one go

Q: Classify each reaction, giving a reason:
(A) 2H2O(l) --electricity--> 2H2(g) + O2(g)
(B) Zn(s) + 2HCl(aq) → ZnCl2(aq) + H2(g)
(C) 2Mg(s) + O2(g) → 2MgO(s)
(D) AgNO3(aq) + NaCl(aq) → AgCl(s) + NaNO3(aq)
Solve: (A) One compound splits into two products using electricity — electrolytic decomposition. (B) The free element zinc replaces hydrogen from the acid — displacement. (C) Two reactants give a single product — combination (also exothermic). (D) Two compounds exchange ions and insoluble white silver chloride precipitates — double displacement (precipitation).

Reaction	Type	Key clue
A	Decomposition	one reactant → many products
B	Displacement	free element replaces an element
C	Combination	many reactants → one product
D	Double displacement	ions exchanged, precipitate forms

Answer: A — decomposition; B — displacement; C — combination; D — double displacement. Tip: count reactants and products first, then look for free elements.

Example 6 — Tricky: the green crystals puzzle

Q: Light green crystals are heated in a dry boiling tube. The colour changes, a brown solid is left behind, and a gas with the smell of burning sulphur is released. Identify the crystals, write the balanced equation, and classify the reaction.
Given: Colour clues — green reactant, brown residue; smell of burning sulphur.
Solve: The clues fit ferrous sulphate crystals (FeSO4·7H2O). On heating they first lose their water of crystallisation; on stronger heating: 2FeSO4(s) --heat--> Fe2O3(s) + SO2(g) + SO3(g). The brown residue is ferric oxide, and the smell comes from the sulphur oxides.
Answer: Thermal decomposition of ferrous sulphate. Watch the trap: two different gases (SO2 and SO3) are released, not one — many students lose a mark by writing only SO2.

Example 7 — Tricky: gas volumes in electrolysis

Q: During the electrolysis of water, 40 mL of gas collects over one electrode and 20 mL over the other. Identify each gas with its electrode, explain the 2 : 1 ratio, and state why a few drops of dilute sulphuric acid are added to the water.
Given: Gas volumes — 40 mL and 20 mL.
Formula: 2H2O(l) --electric current--> 2H2(g) + O2(g).
Solve: The balanced equation gives 2 volumes of hydrogen for every 1 volume of oxygen. So the 40 mL gas is hydrogen, collected at the cathode, and the 20 mL gas is oxygen, collected at the anode. Dilute sulphuric acid is added to make water a good conductor of electricity.
Answer: 40 mL of H2 at the cathode and 20 mL of O2 at the anode; the ratio simply follows the 2 : 1 mole ratio in the equation. Check: 40 = 2 × 20 — consistent.

Example 8 — Tricky: exothermic is not a type of its own

Q: A student writes: "CH4 + 2O2 → CO2 + 2H2O releases a lot of heat, so it is a combination reaction." Point out and correct the error. Also state, with reason, whether respiration is exothermic or endothermic.
Solve: Releasing heat only makes a reaction exothermic; the pattern decides its type. Here the products are TWO substances (CO2 and H2O), so it cannot be a combination reaction — it is the burning (combustion) of natural gas, an exothermic oxidation. For the second part: during respiration, glucose combines with oxygen in our cells — C6H12O6 + 6O2 → 6CO2 + 6H2O + energy — and energy is released.
Answer: The reaction is exothermic but not a combination reaction, because more than one product forms. Respiration is exothermic, since the breakdown of glucose releases energy. Check: combination demands a single product — this has two.

:::keypoints

• Combination: two or more reactants → ONE product; e.g. CaO + H2O → Ca(OH)2 + heat.
• Decomposition: ONE reactant → two or more products; energised by heat (thermal), electricity (electrolytic) or light (photolytic).
• Key decompositions: 2FeSO4 → Fe2O3 + SO2 + SO3; CaCO3 → CaO + CO2; 2Pb(NO3)2 → 2PbO + 4NO2 + O2; 2AgCl → 2Ag + Cl2.
• Electrolysis of water gives hydrogen (cathode) and oxygen (anode) in a 2 : 1 volume ratio.
• Displacement: a more reactive FREE element replaces a less reactive one; Fe + CuSO4 → FeSO4 + Cu (blue fades, nail turns brown).
• Double displacement: two compounds exchange ions, often giving a precipitate; Na2SO4 + BaCl2 → BaSO4↓ + 2NaCl (white).
• Exothermic = heat released (burning, respiration, compost); endothermic = energy absorbed (most decompositions).
• Classify every reaction twice: once by pattern, once by energy.
  :::

:::memory
Marriage – Divorce – Snatch – Swap. Combination = marriage (two become one). Decomposition = divorce (one splits apart). Displacement = snatch (the stronger element walks off with the partner). Double displacement = swap (two couples exchange partners). Run "M-D-S-S" in your head and the four types line themselves up.
:::

:::recap

• Count reactants and products first: many → one is combination; one → many is decomposition.
• A free element on the reactant side signals displacement; two compounds swapping ions signals double displacement.
• Decomposition always needs energy — heat, light or electricity — and is typically endothermic.
• Colour and precipitate observations (blue fades, brown residue, white BaSO4, grey AgCl) are the board's favourite clues.
  :::

Leave an iron gate out through one Indian monsoon and it grows a flaky brown crust. Leave a packet of chips open overnight and by morning it tastes stale and faintly bitter; a silver chain forgotten in a drawer slowly turns black. Three different objects, one shared chemistry — a slow reaction with the oxygen of air, the everyday face of oxidation and its inseparable partner, reduction.

Definition: Oxidation is the gain of oxygen or the loss of hydrogen by a substance during a chemical reaction. Reduction is the exact opposite — the loss of oxygen or the gain of hydrogen.

Definition: A redox reaction (reduction + oxidation) is a reaction in which one substance is oxidised while another is reduced, both in the same reaction. Oxidation never happens alone — every gain has a giver.

The oxygen story and the hydrogen story

At Class 10 level you decide "oxidised or reduced?" by tracking just two elements.

Track oxygen first. If a substance gains oxygen during a reaction, it is oxidised; if it loses oxygen, it is reduced.

• 2Cu + O₂ --heat--> 2CuO — copper gains oxygen, so copper is oxidised to copper(II) oxide.
• C + O₂ → CO₂ — carbon gains oxygen on burning, so carbon is oxidised.
• CuO + H₂ --heat--> Cu + H₂O — copper oxide loses oxygen, so CuO is reduced to copper.

Track hydrogen when there is no oxygen to follow. If a substance loses hydrogen it is oxidised; if it gains hydrogen it is reduced.

• H₂S + Cl₂ → 2HCl + S — hydrogen sulphide loses hydrogen, so H₂S is oxidised to sulphur; chlorine gains hydrogen, so Cl₂ is reduced to HCl.

Notice that this last equation has no oxygen anywhere, yet it is a perfectly good redox reaction. Board papers use exactly this twist to separate students who memorised one definition from students who know both. (In Class 11 you will meet the most general definition — transfer of electrons — but every question in your Class 10 syllabus can be answered with the oxygen and hydrogen definitions above.)

One reaction, two halves — the redox idea

NCERT introduces redox through a lovely activity. Take about 1 g of copper powder in a china dish and heat it. The shiny brown powder becomes coated with a black substance — copper(II) oxide — because copper is oxidised by the oxygen of the air:

2Cu + O₂ --heat--> 2CuO

Now pass hydrogen gas over this hot black material. The black coating turns brown again, because the reverse change takes place and copper is obtained back:

CuO + H₂ --heat--> Cu + H₂O

Look closely at the second equation. Copper oxide loses oxygen: it is reduced. Hydrogen gains oxygen: it is oxidised. Both changes happen inside one single reaction — one substance's loss is literally the other's gain. That is why oxidation and reduction can never occur separately, and why such reactions are called oxidation-reduction or redox reactions.

Common misconception: students write "CuO is reduced" and stop. A substance cannot be oxidised without something else being reduced at the same time — they are two halves of one process. In any redox question, always name both partners; an answer that has oxidation with no matching reduction is chemically impossible.

Oxidising and reducing agents

Once you can spot who is oxidised and who is reduced, the "agent" vocabulary is one small step away.

• The oxidising agent is the substance that causes oxidation of the other reactant — it gives oxygen (or removes hydrogen). In doing this job, it is itself reduced.
• The reducing agent is the substance that causes reduction of the other reactant — it removes oxygen (or supplies hydrogen). In doing this job, it is itself oxidised.

This cross-over trips up many students, so anchor it to the copper reaction:

In CuO + H₂ → Cu + H₂O	CuO	H₂
What happens to it	loses oxygen → reduced	gains oxygen → oxidised
Job done to the other	gives oxygen → oxidising agent	takes away oxygen → reducing agent

Rule of thumb: an agent is named for the job it does to the other substance, and it suffers the opposite fate itself. The oxidising agent ends up reduced; the reducing agent ends up oxidised. This is not just vocabulary — cheap reducing agents such as carbon (coke) are the workhorses of metallurgy, pulling oxygen off metal oxides like ZnO to release the metal.

Corrosion — oxidation that eats metals

When a metal is attacked by substances around it — moisture, acids, gases in the air — the metal is said to corrode, and the process is called corrosion. It is slow, unwanted oxidation of the metal surface.

• Iron rusts: a new iron article is shiny, but kept in a damp place it acquires a flaky, reddish-brown coating of rust (hydrated iron(III) oxide). Rusting needs both air (oxygen) and moisture (water) together — remove either one and iron stays safe.
• Silver blackens: the black layer on old silver articles is silver sulphide, formed by the action of sulphur compounds present in air.
• Copper turns green: copper exposed to moist air slowly gains a green coating (basic copper carbonate) — the dull green film on old copper vessels and statues.

Corrosion damages car bodies, bridges, iron railings, ships and everything else made of iron; an enormous amount of money is spent every year replacing corroded iron. That is why we apply paint on iron articles — the paint layer cuts off contact between the iron surface and the air and moisture, so the oxidation cannot start. Oiling, greasing and galvanising protect iron in the same way (Chapter 3, Metals and Non-metals, takes this idea further).

Rancidity — oxidation that spoils food

Food containing fats and oils faces the same enemy. When fats and oils are oxidised, the food becomes rancid — its smell and taste change unpleasantly. Stale-smelling chips, butter or fried snacks left exposed to air are everyday examples. Manufacturers fight rancidity by keeping oxygen away from the fat:

• Antioxidants — substances which prevent oxidation — are added to foods containing fats and oils.
• Airtight containers slow down oxidation by limiting the supply of air.
• Nitrogen flushing — bags of chips are flushed with unreactive nitrogen gas so that oxygen never touches the chips.
• Refrigeration slows down the oxidation reactions, keeping food fresh for longer.

Why it matters: the whole game of this topic is controlling oxidation. We deliberately speed it up when it is useful (burning fuels for energy, using carbon to win metals from their oxides) and slow it down when it is harmful (rusting bridges, rancid food). Redox is not a definition to memorise — it is a lever that engineers, metallurgists and even chips manufacturers pull every day.

---

Work through these examples in order. The first four build the core skill — tracking oxygen and hydrogen — and the later ones are the twists CBSE papers love: redox with no oxygen, both definitions inside one equation, classify-and-apply questions, and activity-based reasoning. Keep both definition pairs in front of you as you solve.

Example 1 — Who is oxidised, who is reduced?

Q: In the reaction CuO(s) + H₂(g) --heat--> Cu(s) + H₂O(l), identify the substance oxidised and the substance reduced.
Given: the equation is balanced (Cu: 1 = 1, O: 1 = 1, H: 2 = 2).
Formula: gain of oxygen = oxidation; loss of oxygen = reduction.
Solve: CuO becomes Cu — it has lost its oxygen, so CuO is reduced. H₂ becomes H₂O — it has gained oxygen, so H₂ is oxidised.
Answer: H₂ is oxidised (to H₂O); CuO is reduced (to Cu). Sanity check: exactly one partner oxidised and one reduced — a complete redox pair. ✓

Example 2 — Agents at work (ZnO + C)

Q: In the reaction ZnO(s) + C(s) --heat--> Zn(s) + CO(g), identify the substance oxidised, the substance reduced, the oxidising agent and the reducing agent.
Given: the equation is balanced (Zn: 1 = 1, O: 1 = 1, C: 1 = 1).
Solve: ZnO becomes Zn — it has lost oxygen, so ZnO is reduced. C becomes CO — it has gained oxygen, so carbon is oxidised. ZnO supplied the oxygen, so it is the oxidising agent; carbon removed the oxygen, so it is the reducing agent.

Role	Substance
Oxidised	C (carbon)
Reduced	ZnO
Oxidising agent	ZnO
Reducing agent	C

Answer: C is oxidised, ZnO is reduced; ZnO is the oxidising agent and C is the reducing agent. This is exactly how zinc is obtained from its oxide in industry — carbon is the cheap, hard-working reducing agent.

Example 3 — When oxygen itself is a reactant

Q: Identify the substance oxidised and the substance reduced: 4Na(s) + O₂(g) → 2Na₂O(s). (NCERT in-text question)
Formula: the element that gains oxygen is oxidised; the oxygen that is used up is reduced.
Solve: Sodium combines with oxygen to form Na₂O — sodium has gained oxygen, so Na is oxidised. Oxygen is taken up by sodium, so O₂ is reduced.
Answer: Na is oxidised to Na₂O; O₂ is reduced. General rule worth memorising: whenever an element burns or combines with oxygen, the element is oxidised and oxygen itself is reduced (oxygen acts as the oxidising agent).

Example 4 — Everyday reasoning: the nitrogen-flushed packet

Q: Chips manufacturers flush bags of chips with nitrogen gas. Why? Name the process being prevented.
Solve: Chips contain fats and oils. Oxygen from ordinary air slowly oxidises these fats and oils, making the chips rancid — the smell and taste turn unpleasant. Nitrogen is an unreactive gas: filling the packet with nitrogen removes all contact with oxygen, so the oxidation cannot occur.
Answer: Nitrogen prevents oxidation of the fats and oils in the chips, i.e. it prevents rancidity. Related exam favourites: antioxidants, airtight containers and refrigeration all fight the same enemy — oxidation.

Example 5 — Tricky: redox with no oxygen anywhere

Q: In H₂S(g) + Cl₂(g) → 2HCl(g) + S(s), no species contains oxygen. Identify the substance oxidised and the substance reduced, with reasons.
Formula: loss of hydrogen = oxidation; gain of hydrogen = reduction.
Solve: H₂S becomes S — it has lost hydrogen, so H₂S is oxidised. Cl₂ becomes HCl — it has gained hydrogen, so Cl₂ is reduced. For the agents: Cl₂ removes hydrogen from H₂S, so Cl₂ is the oxidising agent; H₂S supplies the hydrogen, so it is the reducing agent.
Answer: H₂S is oxidised to S; Cl₂ is reduced to HCl. Trap: many students freeze when no oxygen appears — switch to the hydrogen definition and the question solves itself.

Example 6 — Tricky: both definitions in one equation

Q: For MnO₂ + 4HCl → MnCl₂ + 2H₂O + Cl₂, identify (i) the substance oxidised, (ii) the substance reduced, (iii) the oxidising agent, (iv) the reducing agent. (NCERT)
Solve: Follow both elements. MnO₂ becomes MnCl₂ — it has lost oxygen, so MnO₂ is reduced. HCl becomes Cl₂ — it has lost hydrogen, so HCl is oxidised. Therefore MnO₂, which causes the oxidation of HCl, is the oxidising agent, and HCl, which removes oxygen from MnO₂, is the reducing agent.
Answer: HCl is oxidised to Cl₂; MnO₂ is reduced to MnCl₂; MnO₂ is the oxidising agent and HCl is the reducing agent. Check: the oxygen leaving MnO₂ appears in the 2H₂O on the product side — every atom is accounted for. ✓

Example 7 — Tricky: classify twice, then apply

Q: Consider Fe₂O₃ + 2Al → Al₂O₃ + 2Fe. (i) Which substance is oxidised and which is reduced? (ii) What type of reaction is this? (iii) State one practical use of this reaction.
Solve: Al becomes Al₂O₃ — aluminium gains oxygen, so Al is oxidised (it is the reducing agent). Fe₂O₃ loses its oxygen to become Fe, so Fe₂O₃ is reduced (it is the oxidising agent). The more reactive aluminium displaces iron from its oxide, so this is a displacement reaction — and since oxidation and reduction occur together, it is a redox reaction too. It is highly exothermic: the iron is produced in the molten state.
Answer: Al is oxidised, Fe₂O₃ is reduced; it is a displacement (and redox) reaction; it is used in thermite welding of railway tracks. One equation, three marks — boards love stacking parts like this.

Example 8 — Tricky: the black-coating puzzle (activity-based)

Q: When copper powder is heated strongly in a china dish, its surface turns black. (i) Why does this happen? (ii) How can the black surface be turned brown again? Write the equations and name the processes involved.
Solve: (i) On heating, copper combines with oxygen from the air and is oxidised to black copper(II) oxide: 2Cu + O₂ --heat--> 2CuO. (ii) Pass hydrogen gas over the hot black material; CuO loses its oxygen and is reduced back to brown copper: CuO + H₂ --heat--> Cu + H₂O.
Answer: The black layer is CuO formed by oxidation of copper; passing H₂ over the hot material reduces CuO back to Cu. Sanity check: the same copper simply cycles Cu → CuO → Cu — oxidation on the way out, reduction on the way back. ✓

:::keypoints

• Oxidation = gain of oxygen or loss of hydrogen; reduction = loss of oxygen or gain of hydrogen.
• Oxidation and reduction always occur together in the same reaction — that is a redox reaction.
• The oxidising agent gives oxygen (or removes hydrogen) and is itself reduced.
• The reducing agent removes oxygen (or supplies hydrogen) and is itself oxidised.
• Anchor example: CuO + H₂ --heat--> Cu + H₂O — CuO is reduced, H₂ is oxidised.
• MnO₂ + 4HCl → MnCl₂ + 2H₂O + Cl₂ — HCl is oxidised (loses H), MnO₂ is reduced (loses O).
• Corrosion is slow oxidation of metals by moisture, acids and gases — iron rusts only when air and moisture are both present.
• Rancidity is oxidation of fats and oils — prevented by antioxidants, airtight packing, nitrogen flushing and refrigeration.
  :::

:::memory
"Oxidation adds Oxygen — both start with O." For hydrogen, flip it with OIL RIG: Oxidation Is Loss of hydrogen, Reduction Is Gain of hydrogen (in higher classes the same OIL RIG works for electrons). For agents: the agent does the job to the other and suffers the opposite itself — the oxidising agent ends up reduced, the reducing agent ends up oxidised.
:::

:::recap

• Follow the oxygen (or, if absent, the hydrogen): gain of O / loss of H means oxidised; loss of O / gain of H means reduced.
• Oxidation and reduction are two halves of one event — always name both partners in a redox answer.
• Agents cross over: the oxidising agent is itself reduced, and the reducing agent is itself oxidised.
• Corrosion of metals and rancidity of food are everyday oxidation — both are controlled by keeping air and moisture away.
  :::

Iron gates turn brown and flaky, silver chains go black, copper vessels grow a green crust, and an opened packet of chips smells odd within days. None of these is an accident — each one is a slow oxidation reaction happening quietly, in air and moisture, on a surface near you. NCERT gives these two everyday effects of oxidation their own names: corrosion for metals and rancidity for food.

Definition: Corrosion is the slow eating away (deterioration) of a metal by the action of air, moisture or chemicals such as acids on its surface. The rusting of iron is the most familiar example.

Definition: Rancidity is the oxidation of the fats and oils present in food on exposure to air, which spoils the smell and taste of the food.

1. Corrosion — metals under slow attack

Most metals shine when freshly cut. Leave them in open air and the shine dulls, because the metal atoms on the surface slowly react with oxygen — often helped by water vapour and acidic gases — to form an oxide, sulphide or carbonate layer. The metal is oxidised (it loses electrons) while the attacking substance is reduced, so corrosion is really a redox reaction in slow motion on the surface of a metal.

Each metal corrodes with its own tell-tale colour, and board examiners love asking for the compound behind the colour:

• Iron → reddish-brown rust: hydrated iron(III) oxide, Fe₂O₃·xH₂O — the flaky layer on gates, tools and railings.
• Silver → black tarnish: silver sulphide (Ag₂S), formed when traces of hydrogen sulphide (H₂S) gas in the air act on silver ornaments.
• Copper → green coating: basic copper carbonate (CuCO₃·Cu(OH)₂), formed when copper reacts with moist air and carbon dioxide — the same chemistry that turns old statues and temple kalash green.

One instructive contrast: aluminium also reacts with oxygen, but its oxide layer (Al₂O₃) is tough and sticks firmly to the surface, sealing the metal below — which is why aluminium vessels stay usable for decades. Rust, on the other hand, is soft and flaky; it keeps falling off and exposing fresh iron, so the attack continues until the whole object crumbles.

2. Rusting of iron — the equation and the famous experiment

When iron is exposed to both oxygen and moisture, it forms hydrated iron(III) oxide — rust:

4Fe(s) + 3O₂(g) + 2xH₂O(l) → 2Fe₂O₃·xH₂O(s)

Read three exam-worthy facts straight off this equation. First, iron is oxidised (it gains oxygen) and oxygen is the oxidising agent. Second, water appears as a reactant — rusting is not a fight between iron and oxygen alone; moisture must join in. Third, the product is hydrated, which is why the x sits in the formula: the number of water molecules trapped inside rust varies.

NCERT proves the "both air and water" condition with a classic three test-tube activity. Take clean iron nails in three tubes. Tube A has ordinary water and is open to air — the nails get both air and water. Tube B has boiled distilled water with a layer of oil on top — boiling drives out the dissolved oxygen and the oil stops fresh air from entering, so the nails get water but no air. Tube C contains anhydrous calcium chloride (a drying agent) and is corked — the nails get air but no moisture. After a few days, only the nails in tube A rust; the nails in B and C stay shiny. Remove either oxygen or water, and rusting stops.

One more real-world layer: dissolved salts make water a much better electrolyte, so salt water speeds up rusting dramatically. That is why ships, and vehicles in coastal cities such as Mumbai and Chennai, rust far faster than those in dry inland towns.

3. Preventing corrosion — build a barrier or use a better metal

Every prevention method obeys one logic: keep oxygen and moisture away from the metal surface, or change the metal itself so that it resists attack.

• Painting, oiling and greasing — the cheapest barriers. Paint on gates and bridges, grease on machine parts, oil on a cycle chain: all of them block air and moisture.
• Galvanisation — coating iron or steel with a thin layer of zinc. Zinc is more reactive than iron, so even if the coating is scratched, the zinc corrodes first and continues to protect the iron beneath. NCERT states this directly: a galvanised article stays protected even if the zinc coating is broken. This is called sacrificial protection.
• Chromium or tin plating — shiny, non-corroding layers electroplated onto taps, cycle handlebars and food cans.
• Alloying — iron mixed with chromium and nickel gives stainless steel, which is hard and does not rust. Your kitchen utensils survive daily washing because of this alloy.

Why it matters: corrosion silently damages car bodies, bridges, iron railings, ships and railway tracks, and every year an enormous amount of money is spent replacing corroded iron. Prevention is far cheaper than replacement — which is why bridges are repainted on schedule and roofing sheets are sold galvanised.

4. Rancidity — when fats and oils go bad

Open a packet of chips left unsealed for a few days: the smell is sharp and unpleasant, the taste flat and bitter. That is rancidity. Fats and oils are long-chain molecules; when oxygen attacks them, they break down into smaller, foul-smelling compounds (including certain aldehydes and ketones). The food's smell and taste change — the two symptoms NCERT names — and it becomes unfit to eat.

Rancidity strikes fastest in fat-rich foods: fried snacks, namkeen, butter, ghee, cooking oils and biscuits. Three conditions speed it up — exposure to air, light and warmth. A tin of ghee on a sunny shelf spoils weeks before the same ghee kept sealed in a cool, dark cupboard.

5. Preventing rancidity — starve the fat of oxygen

Since rancidity is an oxidation, every remedy works by keeping oxygen (or the energy that speeds up oxidation) away from the fat:

• Antioxidants — substances that prevent oxidation — are added to packaged foods containing fats and oils (for example BHA and BHT, which you can spot on biscuit labels).
• Airtight containers limit the food's contact with air, slowing oxidation down. Your steel dabba for fried snacks works on this principle.
• Nitrogen flushing — chips manufacturers flush the packet with nitrogen, an unreactive gas, before sealing it. The puff in a chips packet is nitrogen, not air; with no oxygen inside, the oil on the chips cannot oxidise.
• Refrigeration — lowering the temperature slows the oxidation reaction, so ghee, butter and oils last longer in the fridge. Storing food away from light helps for the same reason.

Why it matters: rancidity connects classroom chemistry to food safety, packaging technology and shelf life — one concept explains the nitrogen puff in a chips packet, the antioxidant line on a biscuit label and why ghee lives in the fridge.

Common misconception: "Iron rusts because of oxygen alone." Correction: iron needs both oxygen and moisture to rust. The three test-tube experiment settles it — nails stay shiny in dry air (tube C) and in air-free water (tube B); only air plus water together (tube A) produce rust.

:::compare

Feature	Corrosion	Rancidity
What is oxidised	A metal (iron, silver, copper)	Fats and oils in food
Trigger	Air + moisture (sometimes acidic gases)	Air (oxygen), aided by light and warmth
Visible result	Coloured oxide/sulphide/carbonate layer	Bad smell, off taste
Common example	Rusting of iron gates	Stale chips, smelly ghee
Prevention	Paint, oil, galvanise, plate, alloy	Antioxidants, airtight pack, N₂ flush, fridge
:::

---

Work through these examples in order. They begin with straight definitions — the guaranteed board 1–2 markers — and climb to the higher-order twists CBSE likes to set: experiment analysis, reactivity reasoning and compare-and-explain chains.

Example 1 — Define and distinguish

Q: Define corrosion and rancidity, giving one example of each. (Board 2-marker)
Solve: Corrosion is an attack on a metal; rancidity is an attack on fats and oils in food. Both are oxidation effects, so name the substance attacked and the everyday result.
Answer: Corrosion is the slow eating away of a metal by the action of air, moisture or chemicals on its surface — e.g. rusting of an iron gate. Rancidity is the oxidation of fats and oils in food on exposure to air, which spoils its smell and taste — e.g. old chips or ghee smelling bad. Check: one attacks metals, the other attacks food — never swap the examples.

Example 2 — The chips-packet question (NCERT in-text)

Q: Why do chips manufacturers flush bags of chips with a gas such as nitrogen?
Solve: Step 1 — chips carry a layer of frying oil; oil + oxygen → oxidation → rancid smell and taste. Step 2 — nitrogen is an unreactive gas, so flushing pushes the oxygen out of the packet and fills it with nitrogen. Step 3 — with no oxygen inside, the oil cannot be oxidised.
Answer: Nitrogen flushing prevents the oxidation of fats and oils (rancidity), keeping the chips fresh and crisp for longer. Sanity check: the "puff" in the packet is a nitrogen cushion, not air.

Example 3 — Name the corrosion coatings

Q: State the colour and the chemical name (with formula) of the layer formed when (a) iron, (b) silver, (c) copper corrode in air.
Solve: Match each metal to its attacker and product: iron reacts with O₂ + moisture; silver with H₂S traces; copper with moist air and CO₂.
Answer: (a) Iron: reddish-brown hydrated iron(III) oxide, Fe₂O₃·xH₂O (rust). (b) Silver: black silver sulphide, Ag₂S. (c) Copper: green basic copper carbonate, CuCO₃·Cu(OH)₂.

Example 4 — Tricky: read the three test-tube experiment

Q: Clean iron nails are kept for a week in three test tubes as described below. (i) In which tube do the nails rust? (ii) Explain the result in each tube. (iii) What changes if tube A contains salt water instead?
Given: Tube A — ordinary water, open to air. Tube B — boiled distilled water with an oil layer on top, corked. Tube C — anhydrous calcium chloride (a drying agent), corked.
Solve: Rust needs BOTH air (O₂) and water. Tube A: both present → rust. Tube B: boiling removed the dissolved oxygen and the oil layer blocks fresh air, so water without air → no rust. Tube C: anhydrous CaCl₂ absorbs every trace of moisture, so air without water → no rust. Salt water is a better electrolyte, so it speeds up the attack.
Answer: (i) Only tube A. (ii) A has air + water; B lacks oxygen; C lacks moisture. (iii) With salt water the nails in A rust faster. Sanity check: knock out either requirement and rusting stops — exactly what tubes B and C demonstrate.

Example 5 — Tricky: scratched zinc vs scratched tin

Q: Two iron buckets are coated — one with zinc (galvanised), the other with tin. Both coatings get deeply scratched, exposing the iron. Which bucket rusts sooner, and why? (HOTS)
Given: Reactivity order: zinc is more reactive than iron; tin is less reactive than iron.
Solve: Zinc, being more reactive than iron, oxidises in preference to iron even after a scratch — sacrificial protection continues at the exposed spot. Tin, being less reactive than iron, protects only as a physical barrier; once the barrier is scratched, the exposed iron corrodes readily at the break.
Answer: The tin-plated bucket rusts sooner; the galvanised one stays protected even with a broken coating — the very reason galvanisation is preferred. Memory check: a more reactive coat is a bodyguard, a less reactive coat is only a raincoat.

Example 6 — Tricky: two cities, one gate

Q: Identical iron gates are installed in two cities. Which gate rusts faster? Give two reasons, and suggest one practical protection for the faster-rusting gate.
Given: City 1 — Jaisalmer: hot, dry desert air. City 2 — Mumbai: humid, salty coastal air.
Solve: Rusting needs moisture along with oxygen. Mumbai's air is humid (plenty of water) and carries sea-salt spray, which makes the surface film of water a better electrolyte and accelerates the redox attack. Jaisalmer's dry air starves the reaction of water, so rusting is very slow there.
Answer: The Mumbai gate rusts much faster — (1) high humidity supplies the moisture rusting needs, and (2) salt in coastal air speeds up the electrochemical attack. Protection: paint the gate or use galvanised iron. Sanity check: dry desert ≈ tube C of the experiment; salty coastal air ≈ tube A with salt water.

Example 7 — Redox inside the rust equation

Q: For the rusting reaction, identify (a) the substance oxidised, (b) the oxidising agent, and (c) state why corrosion is called an "effect of oxidation" in everyday life.
Formula: 4Fe(s) + 3O₂(g) + 2xH₂O(l) → 2Fe₂O₃·xH₂O(s)
Solve: (a) Iron gains oxygen (Fe → Fe₂O₃), so iron is oxidised. (b) Oxygen brings about this oxidation and is itself reduced, so O₂ is the oxidising agent. (c) The damage to the metal happens because the metal is oxidised at its surface by air and moisture.
Answer: (a) Iron; (b) oxygen (O₂); (c) corrosion is the slow, everyday oxidation of a metal — NCERT lists it, along with rancidity, under the effects of oxidation reactions in daily life. Check: in any rust question, iron is always the loser (oxidised), oxygen always the agent.

Example 8 — Tricky: the kitchen investigation

Q: Two halves of the same batch of ghee-fried snacks are stored for a week. Portion X develops a sharp, unpleasant smell; portion Y stays fresh. Explain both observations. Also name the method a chips factory uses and the chemical method a biscuit factory uses against the same problem.
Given: Portion X — open plate on a sunny kitchen shelf. Portion Y — airtight steel dabba kept in the fridge.
Solve: X: the fat is freely exposed to oxygen, and sunlight plus warmth speed up the oxidation, so the fat turns rancid quickly. Y: the airtight dabba cuts off fresh oxygen while the fridge's low temperature slows the reaction — two brakes applied together. Industry: chips packets are flushed with unreactive nitrogen; biscuits get added antioxidants (e.g. BHA/BHT).
Answer: X turned rancid (oxidation of fat, helped by air + light + warmth); Y was protected by airtight storage plus refrigeration. Factory methods: nitrogen flushing (chips) and antioxidants (biscuits). Check: every correct answer here is just one idea in disguise — keep oxygen away from the fat, or slow it down.

:::keypoints

• Corrosion is the slow eating away of metals by air, moisture or chemicals; rusting of iron is the key example.
• Rust is hydrated iron(III) oxide, Fe₂O₃·xH₂O: 4Fe + 3O₂ + 2xH₂O → 2Fe₂O₃·xH₂O — oxygen AND water are both needed.
• Colour clues: iron → reddish-brown rust; silver → black Ag₂S; copper → green CuCO₃·Cu(OH)₂.
• Prevent rusting: paint/oil/grease, galvanise (Zn), chromium or tin plating, alloy into stainless steel (Fe + Cr + Ni).
• Galvanised iron stays protected even when the zinc coat is scratched — zinc is more reactive and corrodes first.
• Rancidity is the oxidation of fats and oils in food; its smell and taste change.
• Prevent rancidity: antioxidants, airtight containers, nitrogen flushing, refrigeration (and dark storage).
• Both corrosion and rancidity are everyday effects of oxidation — link them to redox in every answer.
  :::

:::memory
Rust needs A + W — Air AND Water; miss either one and iron will not rust. To protect iron remember GOPAL: Galvanise, Oil/grease, Paint, Alloy, Layer of chrome or tin. To keep food fresh, switch on the FAAN: Fridge, Airtight dabba, Antioxidants, Nitrogen flush.
:::

:::recap

• Corrosion = slow oxidation of metals; rancidity = oxidation of fats and oils in food.
• Rusting needs both oxygen and moisture — proved by the three test-tube experiment.
• Zinc protects iron sacrificially (galvanisation); stainless steel (Fe + Cr + Ni) does not rust at all.
• Chips packets are flushed with nitrogen to prevent rancidity — NCERT's favourite question.
  :::