Chemical Reaction

Cooking rice, rusting of an iron gate, curd forming from milk, a match burning and a lemon turning a tea light in colour are all chemical reactions. In each of them old substances disappear and new ones appear. This chapter teaches you to describe a reaction in the language of chemistry: how to write it, how to balance it, how to classify it, how fast it goes, and how the acids, bases and salts that you meet every day take part in it.

15.1 Types of reactions

Physical and chemical change

Physical change	Chemical change
No new substance is formed.	One or more new substances are formed.
Usually easy to reverse.	Usually difficult to reverse.
Only physical properties (shape, state) change.	Composition and properties both change.
Melting of ice, dissolving sugar	Burning of wood, rusting, digestion of food

Chemical equation

A chemical equation is the short way of describing a chemical reaction using symbols and formulae. The substances on the left are reactants; the substances on the right are products. The arrow means "gives".

Magnesium burns in oxygen to give magnesium oxide:

2Mg + O₂ → 2MgO

Symbols in brackets show the state: (s) solid, (l) liquid, (g) gas, (aq) dissolved in water. An upward arrow ↑ after a product means a gas is released, and a downward arrow ↓ means a solid (precipitate) settles.

Evidence that a chemical reaction has taken place

• change of colour (iron nail in copper sulphate solution turns brown);
• evolution of a gas (bubbles when zinc is put in acid);
• formation of a precipitate (white solid when silver nitrate is added to salt solution);
• change of temperature (heat released when quicklime is added to water);
• change of state or smell.

Types of chemical reactions

1. Combination (synthesis) reaction

Two or more substances join to form one product.

• C + O₂ → CO₂
• 2H₂ + O₂ → 2H₂O
• CaO + H₂O → Ca(OH)₂ (slaking of lime, heat is given out)

2. Decomposition reaction

One compound breaks into two or more simpler substances. It needs heat, light or electricity.

Type	Example
Thermal decomposition	CaCO₃ --Δ--> CaO + CO₂ ↑
Thermal decomposition	2HgO --Δ--> 2Hg + O₂ ↑
Electrolysis	2H₂O --electricity--> 2H₂ ↑ + O₂ ↑
Photolysis (light)	2AgCl --sunlight--> 2Ag + Cl₂ ↑

The decomposition of limestone (CaCO₃) in a kiln is the first step of making cement and quicklime.

3. Displacement (single replacement) reaction

A more reactive element replaces a less reactive element from its compound.

• Fe + CuSO₄ → FeSO₄ + Cu (the blue solution turns pale green and brown copper settles on the nail)
• Zn + CuSO₄ → ZnSO₄ + Cu
• Zn + 2HCl → ZnCl₂ + H₂ ↑

4. Double displacement (double decomposition) reaction

Two compounds exchange their ions to form two new compounds. If one product is an insoluble solid, it is called a precipitation reaction.

• AgNO₃ + NaCl → AgCl ↓ + NaNO₃ (white precipitate)
• BaCl₂ + Na₂SO₄ → BaSO₄ ↓ + 2NaCl (white precipitate)
• NaOH + HCl → NaCl + H₂O (neutralisation is also double displacement)

5. Exothermic and endothermic reactions

Exothermic reaction	Endothermic reaction
Heat is released to the surroundings.	Heat is absorbed from the surroundings.
The container becomes warm.	The container becomes cold.
Burning of fuel, respiration, CaO + H₂O	Photosynthesis, decomposition of CaCO₃

6. Oxidation and reduction (redox)

In the simple school definition, oxidation is the gain of oxygen or the loss of hydrogen; reduction is the loss of oxygen or the gain of hydrogen. Both happen together in one reaction.

CuO + H₂ → Cu + H₂O

Here CuO loses oxygen, so it is reduced; H₂ gains oxygen, so it is oxidised. The substance that is reduced is the oxidising agent (CuO) and the substance that is oxidised is the reducing agent (H₂).

15.2 Balancing equations

Law of conservation of mass

Antoine Lavoisier showed that in a chemical reaction, mass is neither created nor destroyed. The total mass of the reactants equals the total mass of the products. This means that atoms of every element on the left must equal those on the right. An equation that obeys this is balanced.

Steps to balance an equation (hit and trial method)

1. Write the correct formulae of all reactants and products (skeleton equation).
2. Count the atoms of each element on both sides.
3. Balance the element that appears in the most complex formula first; leave hydrogen and oxygen for last.
4. Put a coefficient (a number) in front of a formula. Never change a subscript inside a formula.
5. Recount all atoms. Check that the equation is balanced.

15.3 Rate of reaction

Some reactions are fast (burning of petrol, firecrackers) and some are slow (rusting, ripening of fruit). The rate of reaction is the speed at which reactants are used up or products are formed.

Collision theory in simple words

A reaction takes place when particles collide with enough energy. The more often they collide, and the harder they collide, the faster the reaction.

Factors affecting the rate

Factor	Effect on rate	Reason / example
Nature of reactants	Depends on the substance	Sodium reacts violently with water; iron reacts slowly.
Concentration	Higher concentration, faster	More particles in the same volume, so more collisions. Marble chips react faster with strong HCl than with dilute HCl.
Surface area (size of particles)	Smaller pieces, faster	More surface is exposed. Powdered chalk reacts faster than a lump; coal dust can explode.
Temperature	Higher temperature, faster	Particles move faster and collide harder. Food cooks faster in a pressure cooker; milk is stored in a fridge to slow souring.
Catalyst	Changes the rate, itself unchanged	MnO₂ speeds up the decomposition of H₂O₂ to H₂O and O₂.
Light	Some reactions speed up	Photosynthesis; AgBr in photographic film.

15.4 Acids, bases and salts

Acids

An acid is a substance that gives hydrogen ions (H⁺) in water. It tastes sour, turns blue litmus red and reacts with metals to release hydrogen.

Acid	Formula	Source / use
Hydrochloric acid	HCl	gastric juice; cleaning
Sulphuric acid	H₂SO₄	car batteries; fertiliser
Nitric acid	HNO₃	fertilisers, explosives
Acetic acid	CH₃COOH	vinegar
Citric acid	C₆H₈O₇	lemon, orange
Lactic acid	C₃H₆O₃	curd
Carbonic acid	H₂CO₃	soda water

Properties of acids

• With metals: Zn + H₂SO₄ → ZnSO₄ + H₂ ↑
• With metal carbonates: CaCO₃ + 2HCl → CaCl₂ + H₂O + CO₂ ↑ (lime water turns milky)
• With bases: HCl + NaOH → NaCl + H₂O

Bases and alkalis

A base is a substance that reacts with an acid to give salt and water. Bases that dissolve in water are called alkalis and give hydroxide ions (OH⁻). Alkalis feel soapy and taste bitter, turn red litmus blue.

Base	Formula	Use
Sodium hydroxide (caustic soda)	NaOH	soap making
Potassium hydroxide	KOH	soft soap
Calcium hydroxide (lime water)	Ca(OH)₂	whitewash, testing CO₂
Ammonium hydroxide	NH₄OH	cleaning, fertiliser
Magnesium hydroxide	Mg(OH)₂	milk of magnesia (antacid)

Indicators

An indicator is a substance that shows by a change in colour whether a solution is acidic or basic.

Indicator	In acid	In base
Blue litmus	red	blue (no change)
Red litmus	red (no change)	blue
Phenolphthalein	colourless	pink
Methyl orange	red	yellow
Turmeric	yellow	red-brown

The pH scale

The pH scale measures how acidic or basic a solution is. It runs from 0 to 14.

• pH less than 7: acidic (the lower the number, the stronger the acid)
• pH equal to 7: neutral (pure water)
• pH more than 7: basic (the higher the number, the stronger the base)

Universal indicator and pH paper show the pH by colour.

Neutralisation

The reaction between an acid and a base to form a salt and water is neutralisation.

Acid + Base → Salt + Water

• HCl + NaOH → NaCl + H₂O
• H₂SO₄ + 2NaOH → Na₂SO₄ + 2H₂O
• 2HCl + Ca(OH)₂ → CaCl₂ + 2H₂O

Applications in daily life: an antacid tablet (Mg(OH)₂) removes extra acid in the stomach; lime (CaO) is added to acidic soil; baking soda soothes a bee sting (acid) while vinegar soothes a wasp sting (alkali); toothpaste (a base) neutralises acids produced by bacteria in the mouth.

Salts

A salt is a compound formed when the hydrogen of an acid is replaced by a metal (or ammonium ion). Common salt is NaCl.

Type of salt	Formed from	Example
Normal salt	complete neutralisation	NaCl, K₂SO₄, CaCO₃
Acid salt	partial replacement of H of an acid	NaHCO₃, NaHSO₄
Basic salt	partial neutralisation of a base	Mg(OH)Cl

Common salts and uses

Salt	Formula	Use
Sodium chloride	NaCl	food, preservative, making NaOH and Cl₂
Sodium bicarbonate (baking soda)	NaHCO₃	baking, antacid, fire extinguisher
Sodium carbonate (washing soda)	Na₂CO₃·10H₂O	washing, softening hard water, glass
Calcium sulphate (plaster of Paris)	CaSO₄·½H₂O	plaster for broken bones, statues
Bleaching powder	CaOCl₂	disinfecting water, bleaching
Blue vitriol	CuSO₄·5H₂O	fungicide

Preparation of salts

A salt can be prepared by (i) acid + metal, (ii) acid + base, (iii) acid + carbonate, or (iv) double displacement. For example, zinc sulphate crystals: add excess zinc to dilute H₂SO₄ until hydrogen stops, filter off the extra zinc, heat the filtrate to saturation and cool it so that crystals of ZnSO₄·7H₂O form.

Chapter summary

• A chemical reaction forms new substances; it is shown by a balanced equation (reactants → products).
• Types: combination, decomposition, displacement, double displacement, exothermic/endothermic and redox.
• Mass is conserved: balance equations by writing coefficients, never by changing subscripts.
• Rate of reaction increases with concentration, surface area, temperature and a catalyst.
• Acids give H⁺, bases (alkalis) give OH⁻; pH below 7 is acidic, 7 is neutral, above 7 is basic.
• Acid + base → salt + water is neutralisation; Acid + metal → salt + H₂; Acid + carbonate → salt + water + CO₂.
• Important salts: NaCl, NaHCO₃, Na₂CO₃·10H₂O, plaster of Paris, bleaching powder.

Practice questions

Group A — Objective questions (1 mark each)

1. The reaction CaO + H₂O → Ca(OH)₂ is a [1]adecomposition reactionbcombination reactioncdisplacement reactionddouble displacement reaction
2. The white precipitate formed on mixing AgNO₃ and NaCl solutions is [1]aAgClbNaNO₃cAg₂OdNaCl
3. Phenolphthalein turns ______ in a base. [1]aredbcolourlesscpinkdyellow
4. A solution has pH 3. It is [1]astrongly basicbneutralcweakly basicdacidic
5. A catalyst [1]ais consumed in the reactionbchanges the rate of reaction without being used upcchanges the productsdalways slows down a reaction
6. Which is the balanced equation? [1]aH₂ + O₂ → H₂Ob2H₂ + O₂ → 2H₂OcH₂ + 2O₂ → 2H₂Od2H₂ + 2O₂ → 2H₂O

Answers to Group A

1. (b) — two substances join to form one.

2. (a) — AgCl is insoluble.

3. (c) — pink in alkali.

4. (d) — below 7.

5. (b) — definition of a catalyst.

6. (b) — 4 H and 2 O on each side.

Group B — Short answer (2 marks each)

1. Balance: Na + H₂O → NaOH + H₂. [2]
2. Why does powdered zinc react faster with acid than a lump of zinc? [2]
3. What is neutralisation? Write one equation. [2]
4. A student burns 12 g of carbon and obtains 44 g of CO₂. How much oxygen reacted? State the law used. [2]
5. Distinguish between exothermic and endothermic reactions with one example each. [2]

Answers to Group B

1. 2Na + 2H₂O → 2NaOH + H₂ (Na 2 = 2, O 2 = 2, H 4 = 4).

2. Powder has a larger surface area, so more particles are exposed to the acid and more collisions occur per second.

3. Reaction of an acid with a base to form a salt and water. HCl + NaOH → NaCl + H₂O.

4. Oxygen = 44 − 12 = 32 g. Law of conservation of mass.

5. Exothermic: releases heat (burning of coal). Endothermic: absorbs heat (decomposition of CaCO₃).

Group C — Long answer (4 marks each)

1. Explain with examples four factors that affect the rate of a chemical reaction. [4]
2. Describe the properties of acids with the help of equations. [4]
3. Write the steps to balance Fe + O₂ → Fe₂O₃ and give the balanced equation. [4]
4. Explain the pH scale and state how you would find whether a liquid is acidic, basic or neutral using litmus paper. [4]

Answers to Group C

1. Concentration (more concentrated acid reacts faster with marble); surface area (powdered chalk faster than a lump); temperature (food cooks faster at higher temperature); catalyst (MnO₂ speeds up decomposition of H₂O₂). Each is explained by more or harder collisions.

2. Sour taste; blue litmus turns red; reaction with metals gives hydrogen (Zn + 2HCl → ZnCl₂ + H₂); with carbonates gives CO₂ (CaCO₃ + 2HCl → CaCl₂ + H₂O + CO₂); with bases gives salt and water (HCl + NaOH → NaCl + H₂O).

3. Fe₂O₃ has 2 Fe and 3 O. Put 2 before Fe: 2Fe. Oxygen is 3 on the right and 2 on the left; take the lowest common multiple 6: 3O₂ on the left and 2Fe₂O₃ on the right; then Fe is 4, so 4Fe. 4Fe + 3O₂ → 2Fe₂O₃.

4. pH is a scale from 0 to 14 that shows acidity or basicity: below 7 acidic, 7 neutral, above 7 basic. With litmus: if blue litmus turns red, the liquid is acidic; if red litmus turns blue, it is basic; if neither changes, it is neutral.