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Concept Labby Aryaman

Chemical Reactions & Equations

Understand how substances transform and how to balance chemical equations.

Chapter Notes

1. Chemical Change and its Signs

  • A chemical reaction is a change in which one or more substances (reactants) are converted into one or more new substances (products) with different properties.
  • Signs that a chemical reaction has occurred: change in state, change in colour, evolution of a gas, change in temperature, formation of a precipitate, and evolution of light or smell.
  • Example activities: magnesium ribbon burning with a dazzling white flame to give white MgO powder; lead nitrate + potassium iodide giving a yellow precipitate of lead iodide; zinc granules + dilute H₂SO₄ giving H₂ gas with warming of the flask.
  • Magnesium ribbon is cleaned with sandpaper before burning to remove the protective layer of magnesium oxide.
2Mg + O₂ →(burn) 2MgO

Watch it happen — Reaction Type Identifier

Controls

Live result

Type
Combination + Oxidation
Energy change
Exothermic
Observation
Dazzling white light, white MgO ash
Reaction Type Identifier — particle levelexothermic
2Mg + O₂2MgOCombination + Oxidation

Watch for colour change, gas evolution, precipitate, and temperature change as signs of a reaction.

Board-style practice questions

Q1.Why is the magnesium ribbon cleaned with sandpaper before burning it in air? Write the balanced equation for the reaction.

2 marks

2. Chemical Equations and Balancing

  • A word equation names the reactants and products; a chemical equation uses symbols and formulae. Reactants are written on the left, products on the right, separated by an arrow pointing towards the products.
  • A skeletal equation is unbalanced. Because of the Law of Conservation of Mass (mass is neither created nor destroyed in a chemical reaction), the number of atoms of each element must be equal on both sides.
  • Balancing by hit-and-trial: list atoms of each element, start with the compound having the maximum number of atoms, adjust coefficients (never change subscripts), and finally check H and O.
  • Make an equation more informative by writing the physical states — (s) solid, (l) liquid, (g) gas, (aq) aqueous — and by writing conditions such as heat (Δ), catalyst, temperature or pressure above/below the arrow.
3Fe(s) + 4H₂O(g) → Fe₃O₄(s) + 4H₂(g)

Watch it happen — Equation Balancer

Controls

1
1
1

Live result

Equation
1H₂ + 1O₂ → 1H₂O
H atoms (L/R)
2 / 2
O atoms (L/R)
2 / 1
Status
Not balanced
Equation Balancer — particle level
H₂ + O₂H₂O
Click to start animation

Balancing follows the Law of Conservation of Mass — atoms are rearranged, never created.

Board-style practice questions

Q1.Balance the equation Fe + H₂O → Fe₃O₄ + H₂ and state the law on which balancing is based. Add physical states.

3 marks

3. Combination Reactions

  • Two or more reactants combine to form a single product: A + B → AB.
  • Examples: quick lime with water giving slaked lime (used in whitewashing; slaked lime later reacts with CO₂ of air to form a shiny CaCO₃ film), burning of coal, and formation of water from H₂ and O₂.
  • Reactions that release heat along with products are exothermic. Respiration, decomposition of vegetable matter into compost, and burning of natural gas are exothermic.
CaO(s) + H₂O(l) → Ca(OH)₂(aq) + heat

4. Decomposition Reactions

  • A single reactant breaks down into two or more products: AB → A + B. These are mostly endothermic because energy must be supplied.
  • Thermal decomposition (by heat): ferrous sulphate crystals lose water and give Fe₂O₃, SO₂ and SO₃; lead nitrate gives brown NO₂ fumes; limestone gives quicklime (used in cement industry).
  • Electrolytic decomposition (by electricity): electrolysis of water gives H₂ at the cathode and O₂ at the anode, in a 2 : 1 volume ratio.
  • Photolytic decomposition (by light): silver chloride and silver bromide turn grey in sunlight — the basis of black-and-white photography.
CaCO₃(s) →(Δ) CaO(s) + CO₂(g)
Board-style practice questions

Q1.Write one equation each for thermal, electrolytic and photolytic decomposition.

2 marks

5. Displacement and Double Displacement Reactions

  • Displacement: a more reactive element displaces a less reactive one from its compound — A + BC → AC + B. Iron nail placed in copper sulphate solution turns brownish and the blue colour fades.
  • Double displacement: ions are exchanged between two compounds — AB + CD → AD + CB. Sodium sulphate + barium chloride gives a white precipitate of barium sulphate; such reactions are called precipitation reactions.
  • Neutralisation of an acid by a base to give salt and water is also a double displacement reaction.
Fe(s) + CuSO₄(aq) → FeSO₄(aq) + Cu(s)

Watch it happen — Displacement Predictor

Controls

Live result

Reaction occurs?
Yes — displacement happens
Reason
The added metal is more reactive
Metal in solution — displacement & electron lossReaction occurs?: Yes — displacement happens
e⁻ transfersalt solution

Reactivity series: K > Na > Ca > Mg > Al > Zn > Fe > Pb > Cu > Ag > Au.

Board-style practice questions

Q1.What happens when an iron nail is dipped in copper sulphate solution? Give the equation, type of reaction and the observation.

3 marks

6. Oxidation, Reduction and Redox

  • Oxidation is gain of oxygen or loss of hydrogen; reduction is loss of oxygen or gain of hydrogen. When both happen in the same reaction it is a redox reaction.
  • In CuO + H₂ → Cu + H₂O, copper oxide is reduced and hydrogen is oxidised. The substance that gets oxidised is the reducing agent; the one that gets reduced is the oxidising agent.
  • Copper turns black on heating in air (CuO forms) and turns brown again when hydrogen is passed over it.
CuO + H₂ →(Δ) Cu + H₂O

Watch it happen — Oxidation & Reduction Tracker

Controls

Live result

Oxidised
H₂ (gains oxygen)
Reduced
CuO (loses oxygen)
Note
H₂ is the reducing agent
Oxidation & Reduction Tracker — particle level
CuO + H₂Cu + H₂OH₂ (gains oxygen)

Oxidation is gain of oxygen / loss of electrons; reduction is the exact opposite. Both happen together.

7. Corrosion and Rancidity — Effects of Oxidation in Daily Life

  • Corrosion: a metal is attacked by moisture, air or chemicals from the surroundings. Iron rusts (reddish-brown), silver tarnishes black (Ag₂S), copper turns green (basic copper carbonate).
  • Rusting is prevented by painting, oiling, greasing, galvanisation (zinc coating), chrome plating, anodising and by making alloys such as stainless steel.
  • Rancidity: fats and oils get oxidised on standing, giving a bad smell and taste. Prevented by adding antioxidants, keeping food in air-tight containers, refrigeration, and packing chips in nitrogen gas.
4Fe + 3O₂ + xH₂O → 2Fe₂O₃·xH₂O (rust)

Watch it happen — Corrosion & Prevention

Controls

60 %
10 %

Live result

Relative rust rate
72%
Verdict
Rapid rusting
Rust rate72
Rusting — iron + O₂ + waterRelative rust rate: 72 %
e⁻ transfersalt solution
Click to start animation

Galvanisation coats iron with zinc, which corrodes sacrificially instead of the iron.

Board-style practice questions

Q1.(a) What is corrosion and rancidity? (b) Give two methods to prevent each. (c) Why are chips packets flushed with nitrogen?

5 marks