Chemical Reactions Class 10 Notes, Formulas, Questions | SSC

This chapter explains what happens when substances change into new substances and how such changes are written as chemical equations. It sets out the rules for writing equations and a step by step trial and error method for balancing them. It then classifies reactions as combination, decomposition, displacement and double displacement, and as endothermic or exothermic, and studies the factors that decide the rate of a reaction. The last part deals with oxidation, reduction, redox reactions, corrosion of iron and rancidity.

Summary

A physical change alters only state or form, caused by changes in temperature or pressure; it is often reversible and the composition stays the same, as when ice melts. A chemical change alters the composition permanently, for example cooking food, ripening of fruit, milk turning into curd and digestion. In a chemical reaction bonds in the reactants break and new bonds form to give products with different properties. For example, carbon and oxygen react to form carbon dioxide.

A reaction can be written as a word equation, and more briefly as a chemical equation using formulae, e.g. CuSO4 + Zn gives ZnSO4 + Cu. Rules: reactants on the left, products on the right, an arrow pointing to the products; plus signs between two or more reactants or products; physical states shown as (s), (l), (g) and (aq); an upward arrow for a gas given off and a downward arrow for a precipitate; a delta sign above the arrow when heat is supplied, and 'heat' written on the product side when heat is given out; conditions such as temperature, pressure and catalyst above or below the arrow (vegetable oil and hydrogen at 60 degrees C with nickel catalyst give vanaspati ghee); and special names or information below the formulae (copper with concentrated nitric acid gives NO2, with dilute nitric acid gives NO).

An equation in which the number of atoms of each element is the same on both sides is a balanced equation; this follows the law of conservation of mass. Balancing is done by trial and error: write the equation, count atoms of each element on both sides, start with the compound having the most atoms and the element that is unequal, apply whole number factors in front of formulae without changing any formula, recheck, and write the final balanced equation. For example NaOH + H2SO4 becomes 2NaOH + H2SO4 gives Na2SO4 + 2H2O.

Reactions are of four types. In a combination reaction two or more reactants form a single product (NH3 + HCl gives NH4Cl; 2Mg + O2 gives 2MgO; CaO + H2O gives Ca(OH)2 with much heat). In a decomposition reaction one reactant gives two or more products (sugar on heating gives carbon and water; CaCO3 on heating gives CaO and CO2, which turns lime water milky; H2O2 slowly gives water and oxygen; electrolysis of acidulated water gives hydrogen and oxygen). In a displacement reaction a more reactive element takes the place of the ion of a less reactive element (Zn, Fe or Pb displace Cu from CuSO4). In a double displacement reaction ions of the reactants are exchanged to form a precipitate (AgNO3 + NaCl gives a white precipitate of AgCl).

Processes and reactions that absorb heat are endothermic; those that give out heat are exothermic. Melting of ice and dissolving KNO3 in water are endothermic processes; freezing of water and dissolving NaOH in water are exothermic processes. Dilution of concentrated sulphuric acid releases a lot of heat, so the acid is always added slowly to water with stirring, never water to acid. Decomposition of CaCO3 on heating is an endothermic reaction, and CaO with water is an exothermic reaction.

Different reactions proceed at different rates, and the same reaction changes rate with conditions (curd sets faster in summer). The rate depends on the nature of the reactants (Al reacts with dilute HCl faster than Zn), the size of particles (powdered Shahabad tile reacts faster than pieces), concentration (concentrated HCl reacts faster with CaCO3), temperature (limestone decomposes only on heating) and a catalyst. A catalyst increases the rate of a reaction without itself undergoing any chemical change; MnO2 speeds up the decomposition of KClO3 and of H2O2. Enzymes act as catalysts in the body, and refrigeration slows the spoiling of food.

Oxidation is a reaction in which a reactant gains oxygen or loses hydrogen (2Mg + O2 gives 2MgO; MgH2 gives Mg + H2). Substances that supply oxygen are oxidants, such as acidified K2Cr2O7, acidified KMnO4, H2O2 (a mild oxidant) and O3; they produce nascent oxygen [O], a reactive form of oxygen before O2 is formed. Acidified dichromate oxidises ethyl alcohol to acetic acid. In terms of electrons, oxidation is the loss of electrons (Fe2+ to Fe3+), that is an increase in positive charge or a decrease in negative charge. Reduction is the gain of hydrogen, loss of oxygen or gain of electrons; the substance that causes it is a reductant. Since oxidation and reduction happen together, as when CuO is reduced to Cu by hydrogen, which is oxidised to water, such reactions are called redox reactions.

Corrosion is the damage to metals caused by their oxidation by components of the atmosphere. Rusting of iron needs both air and water and is faster in the presence of salt. It is electrochemical: at anode regions Fe is oxidised to Fe2+, at cathode regions O2 is reduced to water, and Fe2+ is further oxidised to Fe3+, which forms reddish hydrated iron oxide, rust (Fe2O3.xH2O). Rancidity is the foul smell and taste of oil, ghee or fried food left for a long time, caused by air oxidation; it is prevented by adding antioxidants and storing food in airtight containers.

Mind map

Steps to balance a chemical equation (NaOH + H2SO4)

Important definitions

Formulas

Key points to remember

• Composition changes in a chemical change but not in a physical change.
• Never change a formula while balancing; only put coefficients in front.
• A balanced equation obeys the law of conservation of mass.
• Upward arrow: gas evolved; downward arrow: precipitate formed.
• CO2 turns freshly prepared lime water milky.
• Zn, Fe and Pb displace copper from copper sulphate solution.
• Always add concentrated sulphuric acid to water slowly with stirring.
• Dissolving KNO3 in water is endothermic; dissolving NaOH is exothermic.
• Rate of reaction increases with smaller particle size, higher concentration, higher temperature and a catalyst.
• MnO2 catalyses decomposition of KClO3 and H2O2.
• Oxidation and reduction always occur together in a redox reaction.
• Rust is hydrated iron oxide, Fe2O3.xH2O; rusting needs air and water and is faster with salt.
• Antioxidants and airtight containers prevent rancidity.

Solved examples

Example 1. Balance: N2 + H2 gives NH3.

\( \mathrm{N_2 + H_2 \rightarrow NH_3} \): N is 2 on the left and 1 on the right, so put 2 before NH3: \( \mathrm{N_2 + H_2 \rightarrow 2NH_3} \). Now H is 2 on the left and 6 on the right, so put 3 before H2: \( \mathrm{N_2 + 3H_2 \rightarrow 2NH_3} \). Check: N 2 = 2, H 6 = 6. Balanced.

Example 2. Write a balanced equation: calcium chloride + sulphuric acid gives calcium sulphate + hydrogen chloride.

Skeletal: \( \mathrm{CaCl_2 + H_2SO_4 \rightarrow CaSO_4 + HCl} \). Cl is 2 on the left and 1 on the right; put 2 before HCl. Then H is 2 = 2. Balanced: \( \mathrm{CaCl_2 + H_2SO_4 \rightarrow CaSO_4 + 2HCl} \).

Example 3. Balance: H2S2O7(l) + H2O(l) gives H2SO4(l).

S is 2 on the left and 1 on the right; put 2 before H2SO4. Then H is 4 = 4 and O is 8 = 8. Balanced: \( \mathrm{H_2S_2O_7(l) + H_2O(l) \rightarrow 2H_2SO_4(l)} \).

Example 4. Balance: SO2(g) + H2S(aq) gives S(s) + H2O(l).

O is 2 on the left and 1 on the right; put 2 before H2O. Now H is 4 on the right, so put 2 before H2S. S is now 1 + 2 = 3 on the left; put 3 before S. Balanced: \( \mathrm{SO_2(g) + 2H_2S(aq) \rightarrow 3S(s) + 2H_2O(l)} \).

Example 5. Balance: Ag(s) + HCl(aq) gives AgCl + H2.

H is 1 on the left and 2 on the right; put 2 before HCl. Cl becomes 2, so put 2 before AgCl; then Ag becomes 2, so put 2 before Ag. Balanced: \( \mathrm{2Ag(s) + 2HCl(aq) \rightarrow 2AgCl\downarrow + H_2\uparrow} \).

Example 6. Complete and name the products: CuSO4(aq) + Fe(s) gives ? ; CuSO4(aq) + Pb(s) gives ?

\( \mathrm{CuSO_4(aq) + Fe(s) \rightarrow FeSO_4(aq) + Cu(s)} \): ferrous sulphate and copper. \( \mathrm{CuSO_4(aq) + Pb(s) \rightarrow PbSO_4 + Cu(s)} \): lead sulphate and copper. Both are displacement reactions because Fe and Pb are more reactive than Cu.

Example 7. Identify the substances oxidised and reduced: (a) NiO + H2 gives Ni + H2O (b) 2Mg + O2 gives 2MgO (c) Fe + S gives FeS.

(a) NiO loses oxygen, so it is reduced; H2 gains oxygen, so it is oxidised. (b) Mg gains oxygen, so Mg is oxidised; O2 is reduced. (c) Fe loses electrons to form Fe2+, so Fe is oxidised; S gains electrons to form S2-, so S is reduced.

Example 8. Classify as endothermic or exothermic: (a) HCl + NaOH gives NaCl + H2O + heat (b) 2KClO3 on heating gives 2KCl + 3O2 (c) CaO + H2O gives Ca(OH)2 + heat (d) CaCO3 on heating gives CaO + CO2.

(a) Exothermic, heat is given out. (b) Endothermic, heat must be supplied. (c) Exothermic. (d) Endothermic.

Important questions and answers

Q1. Give reason: When the gas formed on heating limestone is passed through freshly prepared lime water, the lime water turns milky.

Limestone (CaCO3) decomposes on heating into CaO and CO2. The CO2 reacts with lime water, Ca(OH)2, to form insoluble calcium carbonate, which makes the solution milky.

Q2. Give reason: Pieces of Shahabad tile take time to disappear in HCl, but its powder disappears quickly.

The rate of a reaction depends on particle size. The powder has much smaller particles and so a larger surface in contact with the acid, so it reacts faster with HCl and CO2 effervescence is quicker.

Q3. Give reason: Concentrated sulphuric acid is added slowly to water with constant stirring while diluting.

Dilution of concentrated sulphuric acid releases a very large amount of heat. If water is poured into the acid, the water may evaporate instantly and the acid may splash, causing an accident. Adding small amounts of acid to water with stirring releases only a little heat at a time.

Q4. Give reason: Oil should be stored in airtight containers for a long time.

Oil left in contact with air undergoes air oxidation and becomes rancid, giving a foul smell and taste. An airtight container keeps out air and slows this oxidation.

Q5. What is a redox reaction? Explain with an example.

A reaction in which oxidation and reduction occur simultaneously. When hydrogen is passed over heated copper oxide, CuO + H2 gives Cu + H2O. CuO loses oxygen and is reduced to reddish copper, while H2 gains oxygen and is oxidised to water. CuO is the oxidant and H2 the reductant.

Q6. Explain oxidation and reduction with reference to oxygen and hydrogen, with examples.

Oxidation is gain of oxygen or loss of hydrogen: C + O2 gives CO2; MgH2 gives Mg + H2. Reduction is loss of oxygen or gain of hydrogen: CuO + H2 gives Cu + H2O (CuO loses oxygen); vegetable oil + H2 gives vanaspati ghee (oil gains hydrogen).

Q7. How can the rate of decomposition of hydrogen peroxide be increased?

H2O2 decomposes slowly into water and oxygen at room temperature. Adding manganese dioxide powder, which acts as a catalyst, makes it decompose rapidly without MnO2 itself changing.

Q8. Explain the similarity and difference between adding NaOH to water and adding CaO to water.

Similarity: both give out heat, so both are exothermic. Difference: NaOH simply dissolves in water, which is a physical process with no new substance formed; CaO reacts with water to form a new substance, Ca(OH)2, which is a chemical combination reaction.

Q9. Explain the factors that affect the rate of a chemical reaction.

Nature of reactants: Al reacts faster than Zn with dilute HCl. Particle size: smaller particles react faster. Concentration: concentrated HCl reacts faster with CaCO3 than dilute HCl. Temperature: rate increases with temperature, e.g. limestone decomposes only on heating and curd sets faster in summer. Catalyst: MnO2 increases the rate of decomposition of KClO3.

Q10. Explain the rusting of iron.

Rusting is an electrochemical process needing air and water. Some regions of the iron surface act as anode, where Fe is oxidised to Fe2+ giving electrons; other regions act as cathode, where O2 is reduced with H+ ions to water. Fe2+ ions migrate, are further oxidised to Fe3+ and form reddish hydrated iron oxide, Fe2O3.xH2O, called rust. Salt speeds it up.

Q11. What is a catalyst? Give an example.

A substance in whose presence the rate of a chemical reaction increases without any chemical change in itself. MnO2 catalyses the decomposition of KClO3 into KCl and O2.

Q12. Distinguish between a displacement and a double displacement reaction.

In a displacement reaction a more reactive element displaces the ion of a less reactive element, e.g. Zn + CuSO4 gives ZnSO4 + Cu. In a double displacement reaction ions of two compounds are exchanged and a precipitate forms, e.g. AgNO3 + NaCl gives AgCl (white precipitate) + NaNO3.

Exam tips

• Always show physical states and arrows for gas or precipitate when the question asks for a complete equation.
• Write balancing as numbered steps with an atom count table; examiners give step marks.
• To identify reaction type, count reactants and products first: many to one is combination, one to many is decomposition.
• In redox questions name both the substance oxidised and the one reduced, plus the oxidant and reductant.
• Remember the activity examples (lime water, Shahabad tile, KNO3 and NaOH dissolution); they are common 'give reason' questions.

FAQs

Is dissolving a solute in water a chemical reaction?

No. Dissolution does not form a new substance, so it is a physical process, though it may absorb or release heat.

Why does an aluminium utensil lose its shine after some days?

Aluminium reacts with oxygen in air and a thin layer of aluminium oxide forms on its surface, making it dull. This is a form of corrosion.

What is nascent oxygen?

It is the reactive state of oxygen before O2 molecules form, shown as [O]. Oxidants such as acidified KMnO4, acidified K2Cr2O7, H2O2 and O3 produce it.

Which conditions are needed for rusting?

Both air (oxygen) and water are needed. Rusting is faster in salty water; a nail in dry air with anhydrous calcium chloride or in boiled water under oil does not rust.

These notes were prepared with AI help from the Balbharati Science and Technology Part 1 textbook (2026-27) and checked against it.