Chemical Reactions and Equations

1. What is a chemical reaction?

Think of everyday changes: milk turning sour in summer, an iron pan rusting in humid air, grapes fermenting, food being cooked, food being digested, and us breathing (respiration). In each case the identity of the starting substance changes — a brand-new substance with new properties is formed. Whenever such a change happens, we say a chemical reaction has taken place.

The substances we start with are called reactants; the new substances formed are called products. So a chemical reaction is simply: reactants → products.

How do we know a reaction has occurred? NCERT performs three activities (burning magnesium ribbon, mixing lead nitrate + potassium iodide, and zinc + dilute acid). From these, any of the following observations signals a chemical reaction:

• Change in state (e.g. solid magnesium → powdery ash).
• Change in colour (e.g. lead iodide gives a yellow precipitate).
• Evolution of a gas (e.g. zinc + acid bubbles off hydrogen).
• Change in temperature (the test tube becomes warm or cold).

2. From words to a chemical equation

Describing a reaction in a full sentence is long and clumsy. We shorten it in two steps. First a word-equation, then a chemical (symbol) equation.

For Activity 1.1 the word-equation is:

Replacing names with chemical formulae gives the symbolic equation:

Now count the atoms on each side. If the number of atoms of every element is the same on both sides, the equation is balanced. If not, it is a skeletal (unbalanced) equation. In Mg + O₂ → MgO the oxygen atoms are 2 on the left but only 1 on the right — so it is unbalanced.

3. Why equations must be balanced

The law of conservation of mass (from Class 9) states that mass can neither be created nor destroyed in a chemical reaction. So the total mass of products must equal the total mass of reactants. In atomic terms: the number of atoms of each element must be the same before and after the reaction.

An unbalanced equation would mean atoms appeared from nowhere or vanished — which is impossible. Hence every chemical equation must be balanced.

4. Balancing step by step (hit-and-trial method)

NCERT balances Fe + H₂O → Fe₃O₄ + H₂ by the hit-and-trial method, using the smallest whole-number coefficients. The steps:

• Step I: Draw boxes around each formula. Never change anything inside a formula while balancing — only put numbers (coefficients) in front.
• Step II: List the number of atoms of each element on both sides.
• Step III: Start with the compound having the most atoms — here Fe₃O₄ (oxygen). O is 4 on RHS, 1 on LHS, so write 4 in front of H₂O.
• Step IV: Balance hydrogen — 4H₂O gives 8 H on LHS, so put 4 in front of H₂ (4 × 2 = 8).
• Step V: Balance iron — Fe is 3 on RHS, so put 3 in front of Fe.
• Step VI: Re-count to verify all elements match.
• Step VII: Add physical states — (s) solid, (l) liquid, (g) gas, (aq) aqueous (dissolved in water).

Here (g) is written with H₂O because water is used as steam. Reaction conditions (heat, pressure, catalyst) are written above/below the arrow:

5. Combination reactions

When two or more substances combine to form a single product, it is a combination reaction.

More examples:

• Burning of coal: C(s) + O₂(g) → CO₂(g)
• Formation of water: 2H₂(g) + O₂(g) → 2H₂O(l)

Exothermic vs endothermic: Reactions that release heat along with products are exothermic (e.g. slaked lime, burning of natural gas CH₄ + 2O₂ → CO₂ + 2H₂O, respiration, decomposition of vegetable matter into compost). Reactions that absorb energy are endothermic.

Whitewashing: slaked lime reacts slowly with CO₂ in air to form a shiny layer of calcium carbonate (also the formula for marble):

6. Decomposition reactions

In a decomposition reaction a single reactant breaks down into two or more simpler products. It is the opposite of combination. The energy needed to break the reactant down can be supplied as heat, light or electricity, and these reactions are usually endothermic.

By heat (thermal decomposition):

By electricity (electrolysis of water, Activity 1.7): water breaks into hydrogen and oxygen; volume of hydrogen collected is double that of oxygen (because water is H₂O).

By light (photo-decomposition): white silver chloride turns grey in sunlight (used in black-and-white photography):

7. Displacement reactions

When a more reactive element displaces (removes) a less reactive element from its compound, it is a displacement reaction.

Iron is more reactive than copper, so it displaces copper. Other examples (zinc and lead are also more reactive than copper):

8. Double displacement reactions

When two compounds exchange their ions (atoms or groups of atoms) to form two new compounds, it is a double displacement reaction. If one of the products is an insoluble solid (a precipitate), it is also called a precipitation reaction.

The white BaSO₄ precipitate forms from SO₄²⁻ and Ba²⁺ ions; sodium chloride stays in solution. The ions of the two reactants simply swap partners.

9. Oxidation, reduction and redox reactions

Oxidation = gain of oxygen OR loss of hydrogen.
Reduction = loss of oxygen OR gain of hydrogen.

In any reaction where one substance is oxidised, another is simultaneously reduced — together these are oxidation-reduction (redox) reactions.

If hydrogen gas is passed over heated CuO, the black coating turns brown copper again:

More redox examples:

10. Effects of oxidation in everyday life — corrosion and rancidity

Corrosion: when a metal is attacked by substances around it such as moisture, acids, etc., it is said to corrode. The reddish-brown coating on iron is rusting; silver gets a black coating; copper gets a green coating. Corrosion damages car bodies, bridges, iron railings and ships, and replacing rusted iron costs enormous money each year. Prevention: painting, oiling, greasing, galvanising (coating with zinc).

Rancidity: when fats and oils are oxidised, they become rancid — their smell and taste change (food "goes off"). Prevention: adding antioxidants, keeping food in air-tight containers, refrigeration, and flushing packets (e.g. chips) with nitrogen gas so oxygen cannot reach the food.

11. NCERT in-text questions — answered

Page 5 — Q1. Why should a magnesium ribbon be cleaned before burning in air? Magnesium reacts slowly with air to form a layer of magnesium oxide on its surface. This unreactive layer would stop the ribbon from burning properly, so it is cleaned (rubbed with sandpaper) to remove the oxide layer.

Q2. Write the balanced equation for:

• (i) H₂ + Cl₂ → 2HCl
• (ii) 3BaCl₂ + Al₂(SO₄)₃ → 3BaSO₄ + 2AlCl₃
• (iii) 2Na + 2H₂O → 2NaOH + H₂

Q3. Balanced equation with state symbols:

• (i) BaCl₂(aq) + Na₂SO₄(aq) → BaSO₄(s) + 2NaCl(aq)
• (ii) NaOH(aq) + HCl(aq) → NaCl(aq) + H₂O(l)

Page 10 — Q1. A solution of substance 'X' is used for whitewashing. (i) 'X' is calcium oxide (quick lime), CaO. (ii) CaO(s) + H₂O(l) → Ca(OH)₂(aq).

Q2. Why is the gas collected in one test tube in Activity 1.7 double the other? Name it. Water is H₂O — it has twice as many hydrogen atoms as oxygen atoms. On electrolysis it gives twice the volume of hydrogen compared to oxygen. The doubled gas is hydrogen.

Page 13 — Q1. Why does the colour of copper sulphate change when an iron nail is dipped in it? Iron is more reactive than copper, so it displaces copper from copper sulphate, forming iron sulphate (pale green) and free copper. The blue colour of copper sulphate fades: Fe + CuSO₄ → FeSO₄ + Cu.

Q2. Give an example of a double displacement reaction other than Activity 1.10. 2KI(aq) + Pb(NO₃)₂(aq) → PbI₂(s) + 2KNO₃(aq) (a yellow precipitate of lead iodide forms).

Q3. Identify the substances oxidised and reduced:

• (i) 4Na(s) + O₂(g) → 2Na₂O(s): sodium gains oxygen, so Na is oxidised (and O₂ is reduced).
• (ii) CuO(s) + H₂(g) → Cu(s) + H₂O(l): CuO loses oxygen so CuO is reduced; H₂ gains oxygen so H₂ is oxidised.

12. NCERT Exercises — answered (Q1–Q11)

Q1. For 2PbO(s) + C(s) → 2Pb(s) + CO₂(g): lead oxide is reduced (correct), carbon is oxidised (correct). The statements that are incorrect are (a) "lead is getting reduced" and (b) "carbon dioxide is getting oxidised". So the answer is (i) (a) and (b).

Q2. Fe₂O₃ + 2Al → Al₂O₃ + 2Fe is a (d) displacement reaction (aluminium displaces iron — the thermite reaction).

Q3. When dilute HCl is added to iron filings, (a) hydrogen gas and iron chloride are produced: Fe + 2HCl → FeCl₂ + H₂. (Iron chloride forms, not chlorine gas.)

Q4. What is a balanced chemical equation? Why balance? A balanced equation has an equal number of atoms of each element on both sides. We balance because of the law of conservation of mass — atoms (and hence mass) cannot be created or destroyed in a reaction.

Q5. Translate and balance:

• (a) 3H₂ + N₂ → 2NH₃
• (b) 2H₂S + 3O₂ → 2H₂O + 2SO₂
• (c) 3BaCl₂ + Al₂(SO₄)₃ → 2AlCl₃ + 3BaSO₄
• (d) 2K + 2H₂O → 2KOH + H₂

Q6. Balance:

• (a) 2HNO₃ + Ca(OH)₂ → Ca(NO₃)₂ + 2H₂O
• (b) 2NaOH + H₂SO₄ → Na₂SO₄ + 2H₂O
• (c) NaCl + AgNO₃ → AgCl + NaNO₃
• (d) BaCl₂ + H₂SO₄ → BaSO₄ + 2HCl

Q7. Write balanced equations for:

• (a) Ca(OH)₂(aq) + CO₂(g) → CaCO₃(s) + H₂O(l)
• (b) Zn + 2AgNO₃ → Zn(NO₃)₂ + 2Ag
• (c) 2Al + 3CuCl₂ → 2AlCl₃ + 3Cu
• (d) BaCl₂ + K₂SO₄ → BaSO₄ + 2KCl

Q8. Balance and identify the type:

• (a) 2KBr(aq) + BaI₂(aq) → 2KI(aq) + BaBr₂(s) — double displacement.
• (b) ZnCO₃(s) → ZnO(s) + CO₂(g) — decomposition.
• (c) H₂(g) + Cl₂(g) → 2HCl(g) — combination.
• (d) Mg(s) + 2HCl(aq) → MgCl₂(aq) + H₂(g) — displacement.

Q9. Exothermic and endothermic reactions. Exothermic reactions release heat (e.g. burning of natural gas CH₄ + 2O₂ → CO₂ + 2H₂O; respiration; CaO + H₂O → Ca(OH)₂). Endothermic reactions absorb heat (e.g. decomposition of CaCO₃ → CaO + CO₂ on heating; barium hydroxide + ammonium chloride).

Q10. Why is respiration exothermic? During respiration, glucose combines with oxygen in our cells and breaks down to give CO₂, water and energy. Because energy/heat is released, respiration is an exothermic reaction: C₆H₁₂O₆(aq) + 6O₂(aq) → 6CO₂(aq) + 6H₂O(l) + energy.

Q11. Why are decomposition reactions the opposite of combination? In a combination reaction two or more substances join to form a single product; in decomposition a single substance splits into two or more products — exactly reversed. Combination usually releases energy (exothermic); decomposition usually absorbs energy (endothermic). Example pair: 2H₂ + O₂ → 2H₂O (combination) and 2H₂O → 2H₂ + O₂ (decomposition by electrolysis).

13. NCERT Exercises — answered (Q12–Q20)

Q12. One decomposition equation each for heat, light, electricity.

• Heat: CaCO₃(s) →^Heat CaO(s) + CO₂(g)
• Light: 2AgCl(s) →^Sunlight 2Ag(s) + Cl₂(g)
• Electricity: 2H₂O(l) →^Electricity 2H₂(g) + O₂(g)

Q13. Difference between displacement and double displacement (with equations). In displacement, a more reactive element replaces a less reactive one from its compound — only one element moves: Fe + CuSO₄ → FeSO₄ + Cu. In double displacement, two compounds exchange their ions to form two new compounds: Na₂SO₄ + BaCl₂ → BaSO₄ + 2NaCl.

Q14. Refining of silver — silver recovered from silver nitrate by copper. Copper is more reactive than silver, so it displaces silver: 2AgNO₃(aq) + Cu(s) → Cu(NO₃)₂(aq) + 2Ag(s). This is a displacement reaction.

Q15. What is a precipitation reaction? Examples. A reaction in which two solutions react to form an insoluble solid (precipitate) is a precipitation reaction. Examples: Na₂SO₄(aq) + BaCl₂(aq) → BaSO₄(s) + 2NaCl(aq); Pb(NO₃)₂(aq) + 2KI(aq) → PbI₂(s) + 2KNO₃(aq).

Q16. Explain oxidation and reduction in terms of oxygen (two examples each).

• Oxidation (gain of oxygen): 2Cu + O₂ → 2CuO; 2Mg + O₂ → 2MgO.
• Reduction (loss of oxygen): CuO + H₂ → Cu + H₂O (CuO reduced); ZnO + C → Zn + CO (ZnO reduced).

Q17. A shiny brown element 'X' turns black on heating in air. 'X' is copper (Cu); the black compound is copper(II) oxide (CuO): 2Cu + O₂ →^Heat 2CuO.

Q18. Why do we apply paint on iron articles? Paint forms a barrier that keeps out air and moisture, preventing the iron from corroding (rusting). This protects the iron and extends its life.

Q19. Why are oil and fat food items flushed with nitrogen? Nitrogen is unreactive and replaces the oxygen inside the packet. Without oxygen, the fats and oils cannot be oxidised, so the food does not become rancid and stays fresh longer.

Q20. Explain with one example each: (a) Corrosion — slow eating away of a metal when attacked by moisture, air or acids, e.g. rusting of iron forming reddish-brown iron oxide. (b) Rancidity — oxidation of fats and oils in food causing a bad smell and taste, e.g. butter or oil left open going off.

14. Common mistakes to avoid

• Changing formulae while balancing — never change subscripts (e.g. don't write H₂O₄); only adjust the coefficients in front.
• Forgetting to balance the equation, or leaving out state symbols when the question asks for them.
• Mixing up oxidation and reduction — remember OIL RIG: Oxidation Is Loss (of hydrogen / gain of oxygen), Reduction Is Gain (of hydrogen / loss of oxygen).
• Confusing displacement (one element moves) with double displacement (ions are exchanged).
• Calling every single-product reaction "combination" while forgetting some are also redox.
• Writing the arrow conditions wrong — heat/light/electricity go above the arrow, not as reactants.

15. Quick revision checklist

• Reaction signs: change in state, colour, temperature, or gas evolved.
• Balance every equation — law of conservation of mass; only change coefficients.
• Five types: combination, decomposition, displacement, double displacement, redox.
• Combination = many → one; Decomposition = one → many (heat/light/electricity).
• Exothermic releases heat; endothermic absorbs energy.
• Oxidation = +O / −H; Reduction = −O / +H; they occur together (redox).
• Corrosion (rusting) and rancidity are everyday oxidations — keep out air/moisture.