Chemical Changes and Reactions Class 9 ICSE — Chemistry Ch 2
Chemical changes and reactions class 9 ICSE is Chapter 2 of your Chemistry textbook. It explains what happens when substances change into new ones. The chapter covers the conditions that start a reaction, the signs that one has occurred, and the four main reaction types. It also covers energy changes, reversible reactions, and real-world uses like neutralization. This guide walks through each section in the book’s order.
Download Chemical Changes and Reactions Class 9 ICSE Chemistry Chapter 2 PDF
Download the Chemical Changes and Reactions Class 9 ICSE Chemistry chapter PDF to revise offline or print for practice. It is the full chapter as printed in the book — 11 pages, 2 sections.
What is inside Chapter 2
11 pages · 2 sections · 7 worked examples
| Section | Title |
|---|---|
| 2.1 | CHEMICAL REACTION |
| 2.2 | CHARACTERISTICS OF CHEMICAL REACTIONS |
2.1 CHEMICAL REACTION
A chemical reaction breaks the chemical bonds of the reactants and forms new bonds to create products. A chemical bond is the force that holds the atoms of a molecule together. Reactions begin when particles collide. Collisions happen when reactants are in close contact or when energy is supplied.
Seven conditions can set off a chemical change:
- Close contact — two solids react when mixed. White lead nitrate and white potassium iodide form yellow lead iodide: \mathrm{Pb(NO_3)_2(s)} + 2\mathrm{KI(s)} \rightarrow 2\mathrm{KNO_3(s)} + \mathrm{PbI_2(s)}.
- Solution — reactants mix in molten or aqueous form. Sodium chloride and silver nitrate give a white silver chloride precipitate: \mathrm{NaCl(aq)} + \mathrm{AgNO_3(aq)} \rightarrow \mathrm{AgCl\downarrow} + \mathrm{NaNO_3(aq)}.
- Heat — some reactions need heating. The symbol Δ means heat. Copper carbonate decomposes: \mathrm{CuCO_3(s)} \xrightarrow{\Delta} \mathrm{CuO(s)} + \mathrm{CO_2(g)}.
- Light — photochemical reactions absorb light. Silver nitrate decomposes in sunlight: 2\mathrm{AgNO_3} \xrightarrow{\text{sunlight}} 2\mathrm{Ag} + 2\mathrm{NO_2} + \mathrm{O_2}.
- Electricity — electrolysis drives decomposition. Water splits: 2\mathrm{H_2O} \xrightarrow{\text{electricity}} 2\mathrm{H_2} + \mathrm{O_2}.
- Pressure — high pressure forces a reaction. Nitrogen and hydrogen form ammonia above 200 atm: \mathrm{N_2} + 3\mathrm{H_2} \rightleftharpoons 2\mathrm{NH_3}.
- Catalyst — a substance that changes the rate without being consumed. Potassium chlorate decomposes at 700°C alone, but at 300°C with manganese dioxide: 2\mathrm{KClO_3} \xrightarrow{\mathrm{MnO_2}} 2\mathrm{KCl} + 3\mathrm{O_2}.
Catalysts come in three kinds. A positive catalyst speeds a reaction, like MnO₂ speeding hydrogen peroxide decomposition. A negative catalyst slows one, like phosphoric acid retarding hydrogen peroxide. A promoter improves a catalyst’s efficiency, like molybdenum with iron in ammonia manufacture.
Notice what a catalyst does to temperature. Without one, potassium chlorate needs 700°C to decompose. With MnO₂, it starts at 300°C. That is a drop of 400°C, and the catalyst stays unchanged at the end. Sounds can also drive a reaction. Acetylene absorbs sound energy and breaks into carbon and hydrogen: \mathrm{C_2H_2} \xrightarrow{\text{Sound energy}} 2\mathrm{C} + \mathrm{H_2}.
This checklist appears throughout the Class 9 ICSE Chemistry book, so learn each condition with its example.
2.2 CHARACTERISTICS OF CHEMICAL REACTIONS
Four observable signs confirm a chemical reaction. These signs appear in many “what do you observe” questions, so the details matter.
(i) Evolution of gas — one product is a gas. Zinc with dilute sulphuric acid gives hydrogen with effervescence. Effervescence is the formation of gas bubbles in a liquid when one reactant is a liquid. \mathrm{Zn} + \mathrm{H_2SO_4} \rightarrow \mathrm{ZnSO_4} + \mathrm{H_2}.
(ii) Change of colour — the reactants change colour. Iron dropped into blue copper sulphate solution turns it light green as ferrous sulphate forms. Red brown copper deposits at the same time. \mathrm{Fe} + \mathrm{CuSO_4(aq)} \rightarrow \mathrm{FeSO_4(aq)} + \mathrm{Cu}.
(iii) Formation of a precipitate — an insoluble solid forms. A precipitate is an insoluble solid produced in a reaction. Silver nitrate and sodium chloride give white silver chloride: \mathrm{AgNO_3(aq)} + \mathrm{NaCl(aq)} \rightarrow \mathrm{AgCl\downarrow} + \mathrm{NaNO_3(aq)}. Ferrous sulphate with sodium hydroxide gives dirty green ferrous hydroxide: \mathrm{FeSO_4(aq)} + 2\mathrm{NaOH(aq)} \rightarrow \mathrm{Fe(OH)_2\downarrow} + \mathrm{Na_2SO_4(aq)}.
(iv) Change of state — reactants and products have different states. Ammonia gas and hydrogen chloride gas form solid ammonium chloride: \mathrm{NH_3(g)} + \mathrm{HCl(g)} \rightarrow \mathrm{NH_4Cl(s)}. These four signs are the basis of most observation questions in this ICSE chapter.
1. Direct combination or synthesis
A combination reaction joins two or more substances into one product. The general form is \mathrm{A} + \mathrm{B} \rightarrow \mathrm{AB}. When two elements combine, the book calls it synthesis. Three patterns appear:
- Two elements — carbon burns in oxygen: \mathrm{C} + \mathrm{O_2} \xrightarrow{\text{heat}} \mathrm{CO_2} + \text{Heat}. Magnesium burns to 2\mathrm{MgO}.
- Element + compound — carbon monoxide burns in oxygen: 2\mathrm{CO(g)} + \mathrm{O_2(g)} \xrightarrow{\text{heat}} 2\mathrm{CO_2(g)}.
- Compound + compound — ammonia and hydrogen chloride form ammonium chloride: \mathrm{NH_3(g)} + \mathrm{HCl(g)} \rightarrow \mathrm{NH_4Cl(s)}.
The book gives two demonstrations. Black lead sulphide heated in a test tube reacts with oxygen to form white lead sulphate: \mathrm{PbS} + 2\mathrm{O_2} \rightarrow \mathrm{PbSO_4}. A magnesium ribbon burns with a dazzling light to give magnesium oxide.
2. Decomposition
Decomposition breaks one compound into elements or simpler compounds. The products do not recombine to form the original compound. It can happen by heat, light, or an electric current. Decomposition by heat is thermal decomposition.
Three product patterns appear. A compound can break into elements: 2\mathrm{HgO(s)} \xrightarrow{\Delta} 2\mathrm{Hg(l)} + \mathrm{O_2(g)}. It can break into an element and a compound: potassium nitrate on heating gives potassium nitrite and oxygen. It can break into two compounds: \mathrm{CaCO_3(s)} \xrightarrow{\text{heat}} \mathrm{CaO(s)} + \mathrm{CO_2(g)}.
The chapter lists thermal decomposition families. Metal hydroxides decompose to oxide and water. Metal carbonates give oxide and carbon dioxide. Metal bicarbonates or hydrogen carbonates give carbonate, water, and carbon dioxide: 2\mathrm{NaHCO_3} \xrightarrow{\Delta} \mathrm{Na_2CO_3} + \mathrm{H_2O} + \mathrm{CO_2}. Metal nitrates give oxide, nitrogen dioxide, and oxygen.
Four experiment observations matter. Heating lead nitrate leaves yellow lead monoxide and gives brown nitrogen dioxide and oxygen. Heating zinc carbonate gives carbon dioxide that turns lime water milky; zinc oxide is yellow hot but white cool. Orange ammonium dichromate swells and leaves green chromium oxide with nitrogen and water vapour. Blue hydrated copper sulphate turns white as water is driven off.
Digestion is a decomposition reaction in the body. Starch is decomposed by enzymes into glucose, which is then oxidized into carbon dioxide and water: \text{Starch} \xrightarrow{\text{Enzymes}} \text{Glucose} \xrightarrow{[O]} \mathrm{CO_2} + \mathrm{H_2O}.
Reversible reactions and thermal dissociation
Some reactions can run both ways. The sign ⇌ shows this. The direction of chemical change reverses when conditions change. The reaction ends with a mixture of reactants and products.
The chapter gives four cases. Hydrated copper sulphate (blue) loses water to become white anhydrous copper sulphate. Adding water restores the blue: \mathrm{CuSO_4 \cdot 5H_2O(s)} \rightleftharpoons \mathrm{CuSO_4(s)} + 5\mathrm{H_2O(g)}. This blue-to-white-to-blue change is the test to detect the presence of water.
Steam over red hot iron gives magnetic oxide and hydrogen: 3\mathrm{Fe} + 4\mathrm{H_2O} \rightleftharpoons \mathrm{Fe_3O_4} + 4\mathrm{H_2}. Ammonium chloride heats to ammonia and hydrogen chloride, which recombine as a white sublimate on the cooler part of the tube: \mathrm{NH_4Cl} \rightleftharpoons \mathrm{NH_3\uparrow} + \mathrm{HCl\uparrow}. Nitrogen tetraoxide (colourless) heats to reddish brown nitrogen dioxide and cools back: \mathrm{N_2O_4} \rightleftharpoons 2\mathrm{NO_2}.
Thermal dissociation is a reversible decomposition brought about only by heat. One line separates it from thermal decomposition. Thermal decomposition is not reversible. Thermal dissociation is reversible. The arrow tells you which is which.
3. Displacement
Displacement is a chemical change in which a more active element displaces a less active element from its salt solution. The chapter shows three experiments.
- Zinc in copper sulphate solution — the blue colour fades to colourless, and reddish brown copper settles. \mathrm{CuSO_4} + \mathrm{Zn} \rightarrow \mathrm{ZnSO_4} + \mathrm{Cu\downarrow}.
- Magnesium in dilute sulphuric acid — brisk effervescence, and hydrogen burns with a pop. \mathrm{Mg} + \mathrm{H_2SO_4} \rightarrow \mathrm{MgSO_4} + \mathrm{H_2}.
- Chlorine gas through potassium iodide — the colourless solution turns yellow brown as iodine appears. 2\mathrm{KI} + \mathrm{Cl_2} \rightarrow 2\mathrm{KCl} + \mathrm{I_2}.
These observations build an activity series. Metals run K, Na, Ca, Mg, Al, Zn, Fe, Pb, [H], Cu, Hg, Ag, Au, Pt with reactivity decreasing downwards. Non-metals run F, Cl, Br, I with reactivity decreasing downwards. The rule: the more reactive element displaces the less reactive one from its salt solution.
Here is the rule one step further. Chlorine sits above bromine in the non-metal order. So chlorine displaces bromine from a bromide solution. Pass chlorine water into sodium bromide solution: 2\mathrm{NaBr} + \mathrm{Cl_2} \rightarrow 2\mathrm{NaCl} + \mathrm{Br_2}. The colourless solution turns orange-brown as bromine forms.
4. Double decomposition
Double decomposition exchanges radicals between two compounds in solution. It is also called double displacement. The general form is \mathrm{AB} + \mathrm{CD} \rightarrow \mathrm{AD} + \mathrm{CB}. Two types appear: precipitation and neutralization.
Precipitation forms an insoluble salt as one product. Barium chloride with sodium sulphate gives white barium sulphate: \mathrm{BaCl_2(aq)} + \mathrm{Na_2SO_4(aq)} \rightarrow \mathrm{BaSO_4(s)} + \mathrm{NaCl(aq)}. Copper sulphate with hydrogen sulphide gives black copper sulphide: \mathrm{CuSO_4(aq)} + \mathrm{H_2S(g)} \rightarrow \mathrm{CuS(s)} + \mathrm{H_2SO_4(aq)}. Silver nitrate with hydrochloric acid gives white silver chloride: \mathrm{AgNO_3} + \mathrm{HCl} \rightarrow \mathrm{AgCl\downarrow} + \mathrm{HNO_3}.
Neutralization is the reaction between an acid and a base that forms salt and water only. Sodium hydroxide with hydrochloric acid: \mathrm{NaOH} + \mathrm{HCl} \rightarrow \mathrm{NaCl} + \mathrm{H_2O}. In ionic form, \mathrm{Na^+OH^-} + \mathrm{H^+Cl^-} \rightarrow \mathrm{Na^+Cl^-} + \mathrm{H^+OH^-}. Cancel the common ions and the only change is \mathrm{H^+} + \mathrm{OH^-} \rightarrow \mathrm{H_2O}.
Double decomposition can also evolve gas. Iron sulphide with sulphuric acid gives hydrogen sulphide: \mathrm{FeS(s)} + \mathrm{H_2SO_4(aq)} \rightarrow \mathrm{FeSO_4} + \mathrm{H_2S\uparrow}. Sodium hydrogen carbonate with hydrochloric acid gives carbon dioxide: \mathrm{NaHCO_3} + \mathrm{HCl} \rightarrow \mathrm{NaCl} + \mathrm{H_2O} + \mathrm{CO_2}.
| Reaction type | General form | Products | Chapter example |
|---|---|---|---|
| Combination | \mathrm{A} + \mathrm{B} \rightarrow \mathrm{AB} | One | \mathrm{C} + \mathrm{O_2} \rightarrow \mathrm{CO_2} |
| Decomposition | \mathrm{AB} \rightarrow \mathrm{A} + \mathrm{B} | Two or more | 2\mathrm{HgO} \rightarrow 2\mathrm{Hg} + \mathrm{O_2} |
| Displacement | \mathrm{A} + \mathrm{BC} \rightarrow \mathrm{AC} + \mathrm{B} | Two | \mathrm{CuSO_4} + \mathrm{Zn} \rightarrow \mathrm{ZnSO_4} + \mathrm{Cu} |
| Double decomposition | \mathrm{AB} + \mathrm{CD} \rightarrow \mathrm{AD} + \mathrm{CB} | Two | \mathrm{AgNO_3} + \mathrm{NaCl} \rightarrow \mathrm{AgCl} + \mathrm{NaNO_3} |
As a chemical process, neutralization has many uses
Neutralization appears in daily life. A bee sting puts formic acid into the skin. Rubbing slaked lime or baking soda, both bases, relieves the pain. Excess hydrochloric acid from the stomach glands is relieved by milk of magnesia or sodium hydrogen carbonate. Acid spilled on clothes or body is neutralized with ammonia solution. Acidic soil is treated with slaked lime.
Hydrolysis is the process in which a salt and water react to form an acidic or a basic solution. Only certain salts hydrolyse. A salt of a strong base and weak acid makes a basic solution, which turns red litmus blue. Examples: \mathrm{Na_2CO_3} and \mathrm{K_3PO_4}. A salt of a strong acid and weak base makes an acidic solution, which turns blue litmus red. Examples: \mathrm{FeCl_3} and \mathrm{ZnSO_4}.
Strong acids and strong bases ionize largely in water. Weak ones ionize only a little. Salts formed by the weak partners usually do not hydrolyse.
Exothermic and endothermic reactions
Every reaction involves an energy change. Bond breaking absorbs heat. Bond formation releases heat. The reaction ends with either a surplus or a deficit. That decides the type.
An exothermic reaction gives out heat and raises temperature. Carbon burning in oxygen releases heat: \mathrm{C} + \mathrm{O_2} \rightarrow \mathrm{CO_2} + \text{Heat}. Water added to quicklime gives slaked lime and heat: \mathrm{CaO} + \mathrm{H_2O} \rightarrow \mathrm{Ca(OH)_2} + \text{Heat}. Hydrogen burning in oxygen forms water with heat: 2\mathrm{H_2} + \mathrm{O_2} \rightarrow 2\mathrm{H_2O} + \text{Heat}. Ammonia forms with an iron catalyst at 450–500°C and releases heat.
Respiration, rusting, and burning coal, petrol, or kerosene are common exothermic reactions.
An endothermic reaction absorbs heat and lowers temperature. Carbon heated with sulphur forms carbon disulphide: \mathrm{C} + 2\mathrm{S} \xrightarrow{\Delta} \mathrm{CS_2}. Nitrogen and oxygen heated to about 3000°C form nitric oxide: \mathrm{N_2} + \mathrm{O_2} \xrightarrow{\text{Heat (3000°C)}} 2\mathrm{NO(g)}. Calcium carbonate decomposes at about 1000°C: \mathrm{CaCO_3(s)} \xrightarrow{\text{Heat}} \mathrm{CaO(s)} + \mathrm{CO_2(g)}.
Two related definitions close the chapter. A photochemical reaction absorbs light energy; photosynthesis is the example. An electrochemical reaction absorbs electrical energy; fused potassium chloride splits into ions: \mathrm{KCl} \xrightarrow{\text{electric current}} \mathrm{K^+} + \mathrm{Cl^-}.
What to practise from Chemical Changes and Reactions Class 9 ICSE
Two exercise blocks close this chapter, and both map directly to the sections above.
- Exercise 2(A) — 7 questions: definitions (chemical reaction, chemical bond, effervescence, precipitate), the conditions for a chemical change with one example each, photochemical and electrochemical reactions, and reasons such as why silver nitrate is kept in coloured bottles and why molybdenum is used in ammonia manufacture.
- Exercise 2(B) — 10 questions: reaction types, completing statements, the difference between thermal decomposition and thermal dissociation, synthesis, neutralization applications, hydrolysis, and complete-and-balance equations.
Every question form has a home in this ICSE chapter. A “define” question pulls from section 2.1. A “name the type” question pulls from the four reaction types. A “what do you observe” question pulls from the four characteristics in 2.2. A “complete and balance” question pulls from the equations across both sections. When you practise, name the question form, then go to the section that answers it.
For more practise material, see the Class 9 ICSE books collection. The official CISCE website carries the Class 9 ICSE Chemistry syllabus that frames these exercises.
Chemical Changes and Reactions Class 9 ICSE frequently asked questions
What are the important topics in chemical changes and reactions class 9 ICSE?
The chapter covers seven conditions for a chemical change, four observable characteristics, and four reaction types. It also covers reversible reactions, thermal dissociation, and exothermic and endothermic changes. Neutralization and hydrolysis close the chapter.
What is the difference between thermal decomposition and thermal dissociation in this ICSE chapter?
Thermal decomposition breaks a compound by heat, and the products do not recombine. Thermal dissociation is a reversible decomposition by heat, so the products can reform the original compound. The ⇌ arrow in the equation tells you a dissociation is reversible.
What is the difference between displacement and double displacement in Class 9 ICSE Chemistry?
Displacement involves one element replacing another from a salt solution. Double displacement involves two compounds exchanging radicals in solution. In displacement one new element appears; in double displacement two new compounds form.
Why is silver nitrate solution kept in coloured bottles in this ICSE chapter?
Silver nitrate decomposes in light, forming silver, nitrogen dioxide, and oxygen. The coloured bottle blocks light and stops the decomposition. This is a photochemical reaction.
How does hydrated copper sulphate show that a reaction is reversible in Class 9 ICSE Chemistry?
Blue hydrated copper sulphate loses water on heating and turns white. Adding water restores the blue colour. The reversible equation uses ⇌, and this colour change is the test for the presence of water.
Reference: ICSE Class 9 Chemistry textbook material, chapter 2: Chemical Changes and Reactions — CISCE-aligned.