NCERT Exemplar Solutions Class 10 Science Chapter 1

Reverse-the-change test. I always ask one question: can I undo this and get the same substance back? If yes, it is physical; if no, it is chemical.

Concept used. In a physical change the chemical formula of the substance is unchanged. In a chemical change at least one bond is broken and a new bond is made, so a new formula appears among the products.

1. Option (a) boiling water. H₂O(l) → H₂O(g). Same molecule, only the state changes. Condense the vapour and you get water back. Physical.
2. Option (b) melting ice. H₂O(s) → H₂O(l). Freeze it again and the ice returns. Physical.
3. Option (c) dissolving salt. NaCl ions spread through water; evaporate the water and the salt crystals come back. Physical.
4. Option (d) burning LPG. The fuel is destroyed and carbon dioxide plus water vapour are created. There is no way to cool the smoke and recover LPG. Chemical.

Why this matters. Telling physical from chemical change is the gate to the whole chapter. Combustion, rusting, cooking and digestion are all chemical because they make new substances.

Final Answer: Option (d): combustion of LPG, the only chemical change in the list.

Track the oxygen and hydrogen. For a Class 10 reaction the quickest redox test is to follow where oxygen and hydrogen go. If oxygen is added somewhere and removed somewhere, it is redox.

Concept used. Oxidation is gain of oxygen or loss of hydrogen; reduction is the reverse. Displacement means one element takes the place of another in a compound.

1. Oxidation half. NH₃ has nitrogen tied to hydrogen. In NO that nitrogen is tied to oxygen instead. Nitrogen has gained oxygen and lost hydrogen, so ammonia is oxidised.
2. Reduction half. Free O₂ is electrically neutral. In H₂O that oxygen now carries a share of hydrogen's electrons, so molecular oxygen is reduced.
3. Why displacement too. Oxygen kicks hydrogen off nitrogen and bonds to nitrogen in its place. One element replacing another inside a compound is the displacement pattern.
4. Eliminate the rest. Combination needs a single product; we have two. Neutralisation needs acid and base; neither is present. So (ii) and (iv) are out.

Why this matters. Many reactions wear two labels at once. Reading the question as "tick all that apply" stops students from settling for the first match they see.

Final Answer: Option (c): statements (i) displacement and (iii) redox are both correct.

Name who-does-what to who. In every redox line, first decide which species is oxidised and which is reduced, then read off the agents from those two findings.

Concept used. A reducing agent is itself oxidised; an oxidising agent is itself reduced. Oxidation here means gaining oxygen and reduction means losing oxygen.

1. Iron's fate. Fe (no oxygen) becomes Fe₃O₄ (full of oxygen). Iron gains oxygen, so iron is oxidised. That makes iron the reducing agent.
2. Water's fate. H₂O (has oxygen) becomes H₂ (no oxygen). Water loses oxygen, so water is reduced. That makes water the oxidising agent.
3. Score the statements. (i) iron oxidised: yes. (ii) water reduced: yes. (iii) water as reducing agent: no, water is the oxidising agent. (iv) water as oxidising agent: yes.
4. So the correct set is (i), (ii) and (iv), which is option (c).

Why this matters. This very reaction is how blacksmiths once made iron react with steam. Getting the agent labels right is what separates a careful answer from a guess.

Final Answer: Option (c): statements (i), (ii) and (iv) are correct.

Feel the beaker. The simplest classroom test for these is whether the beaker gets hot (exothermic) or cold (endothermic) to touch.

Concept used. Heat flows out in an exothermic change and into the system in an endothermic change. State changes from liquid to gas or solid to gas always absorb heat.

1. Quick lime plus water. This slaking reaction is famous for boiling its own water. Strongly exothermic. Mark (i).
2. Diluting acid. Concentrated acid mixing with water releases hydration energy as heat. Exothermic. Mark (ii).
3. Evaporation. Liquid water must absorb the latent heat of vaporisation to leave as vapour. Endothermic, so reject (iii).
4. Sublimation of camphor. Solid to gas needs energy to break the solid's forces. Endothermic, so reject (iv).

Why this matters. Cooks, chemists and even cement workers rely on the heat of slaking lime, while sweat cooling our skin is everyday endothermic evaporation.

Final Answer: Option (a): processes (i) and (ii) are exothermic.

Read the thermometer. The whole question reduces to one rule: hotter solution means heat released (exothermic), colder solution means heat absorbed (endothermic).

Concept used. Dissolving a substance breaks its lattice (needs energy) and hydrates the ions (releases energy). If the release wins, the solution warms; if the lattice cost wins, it cools.

1. Beaker A, NaOH. Sodium hydroxide releases strong hydration energy, so the water heats up. Exothermic.
2. Beaker B, anhydrous CuSO₄. The dry salt grabs water of crystallisation, releasing heat. The solution heats up. Exothermic.
3. Beaker C, NaCl. For common salt the energy to break the lattice slightly exceeds the hydration energy, so the solution cools a little. Endothermic.
4. Putting these together, exothermic in A and B (statement i) and endothermic in C (statement iv): option (c).

Why this matters. This experiment shows that not all dissolving is the same. Whether a salt heats or cools its water is a balance of two hidden energy terms.

Final Answer: Option (c): statements (i) and (iv) are correct.

Colour change is the clue. The fading of permanganate purple is the textbook signature of a redox titration, so the explanation must be a reaction, not dilution.

Concept used. Permanganate is reduced from MnO₄^- (oxidation state + 7, purple) to Mn²⁺ (oxidation state + 2, colourless) while it oxidises ferrous to ferric ions.

1. What gets oxidised. Fe²⁺ loses one electron to become Fe³⁺. So FeSO₄ is the reducing agent and is oxidised.
2. What gets reduced. The manganese drops from + 7 to + 2, gaining electrons, so KMnO₄ is the oxidising agent and is reduced.
3. Why the colour goes. The purple belongs to MnO₄^-. Once it becomes Mn²⁺, that purple is gone, which is exactly what is seen.
4. Reject the others. Dilution would only lighten the purple, never wipe it out completely, so (c) is wrong; and KMnO₄ does not self-decompose here, so (d) is wrong.

Why this matters. This is the basis of permanganate titrations used in labs to measure iron content, where the end point is the first lasting pink.

Final Answer: Option (a): permanganate oxidises ferrous sulphate, and being reduced to Mn²⁺ its purple colour disappears.

Count the swaps. Single displacement swaps one element; double displacement swaps two ion pairs at once. Counting swaps sorts these fast.

Concept used. Double displacement follows AB + CD → AD + CB, with both cation and anion finding new partners and usually a precipitate, gas or water forming.

1. Reaction (ii). Na₂SO₄ gives Na+ and SO₄^2-; BaCl₂ gives Ba²⁺ and Cl-. They recombine as BaSO₄ (insoluble) and NaCl. Two swaps, so double displacement.
2. Reaction (i). Free lead pushes copper out of CuCl₂. Only one element is replaced, so it is single displacement.
3. Reactions (iii) and (iv). Carbon plus oxygen, and methane plus oxygen, are combination and combustion. They build oxides; nothing is exchanged.
4. Hence only (ii) qualifies, giving option (b).

Why this matters. Precipitation tests in the lab, like the formation of white BaSO₄, rely on double displacement to detect sulphate ions in an unknown sample.

Final Answer: Option (b): reaction (ii) alone is a double displacement reaction.

Trace where the grey comes from. The grey solid is the leftover metal, so the cause must be a decomposition that frees silver.

Concept used. Light supplies the energy for the decomposition 2AgCl → 2Ag + Cl₂. The freed silver atoms cluster into grey metal, while chlorine escapes as a gas.

1. Identify the grey. Silver metal is grey. Its appearance tells us silver was set free from the white AgCl.
2. Name the process. Breaking one compound into simpler pieces using light is photochemical decomposition. So (i) is correct.
3. Reject sublimation. Sublimation would carry AgCl away as vapour, not leave grey metal, so (ii) is wrong.
4. Reject (iii) and (iv). The grey is not chlorine (a pale green gas) and not an oxidation product; it is plain silver. Both are wrong.

Why this matters. The same reason silver chloride is stored in dark bottles, to stop this light-driven breakdown, links straight to Question 34 later in this chapter.

Final Answer: Option (a): only statement (i) is true, the grey is silver from photo-decomposition.

Two facts settle it. One: does the reaction heat up? Two: is the product acidic or basic? Answer both and the option falls out.

Concept used. An exothermic reaction warms its surroundings. The product calcium hydroxide is a base, and bases give pH greater than 7 in water.

1. Energy direction. Quick lime plus water famously gives off so much heat the water can boil. Heat out means exothermic, confirming (ii).
2. Nature of the product. Ca(OH)₂ is a hydroxide, so it is a base. Lime water turns red litmus blue.
3. pH of a base. A basic solution sits above 7 on the pH scale, so (iii) is correct and (iv) wrong.
4. Combining the exothermic finding with the basic pH gives (ii) and (iii): option (b).

Why this matters. Lime water is used to test for carbon dioxide gas, which turns it milky. Knowing it is basic explains why it reacts with the acidic gas CO₂.

Final Answer: Option (b): statements (ii) exothermic and (iii) pH above 7 are correct.

Two labels, one reaction. Whenever a swap of ions throws down an insoluble solid, the reaction earns both names at once.

Concept used. Double displacement is the exchange AB + CD → AD + CB. If AD or CB is insoluble, that solid is a precipitate and the reaction is also called precipitation.

1. Write and check. BaCl₂ + (NH₄)₂SO₄ → BaSO₄ + 2NH₄Cl. Barium leaves chloride for sulphate, ammonium leaves sulphate for chloride. Both pairs swap, so double displacement.
2. Spot the precipitate. BaSO₄ is famously insoluble in water and drops out as a white solid, so this is also precipitation.
3. Eliminate the rest. No free element is set free, ruling out displacement (i); two products form, ruling out combination (iii).
4. So both (ii) and (iv) apply, giving option (d).

Why this matters. The insoluble white BaSO₄ is the same "barium meal" used in medical X-rays of the digestive tract, because it shows up clearly and does not dissolve.

Final Answer: Option (d): statements (ii) precipitation and (iv) double displacement are both correct.

Read it off the balanced line. For any gas ratio in a reaction, the coefficients in the balanced equation are the whole answer.

Concept used. In a balanced equation the mole ratio of any two species equals the ratio of their coefficients. Water is H₂O, with two H per one O.

1. Balance the split. 2H₂O → 2H₂ + O₂. Four H and two O on each side, so it is balanced.
2. Pick the coefficients. H₂ carries a 2; O₂ carries a 1.
3. Form the ratio. Hydrogen : oxygen = 2 : 1. This matches the 2-to-1 atom count in water itself.
4. So the volume of hydrogen collected is double that of oxygen, option (b).

Why this matters. In the classic electrolysis experiment the test tube over the negative electrode collects twice the gas of the positive electrode, a visible proof of this 2:1 ratio.

Final Answer: Option (b): the mole ratio of hydrogen to oxygen is 2:1.

Energy goes which way? If a step needs heat fed in to happen, it is endothermic; if it gives heat out, it is exothermic.

Concept used. Breaking the forces in a solid or liquid to make a gas costs energy, so it is endothermic. Building those forces back up on condensing or freezing releases energy.

1. Sublimation of dry ice. Solid CO₂ must absorb heat to jump to gas, leaving a cold fog. Endothermic, so (ii).
2. Evaporation of water. Liquid takes in latent heat to become vapour, which is why sweat cools us. Endothermic, so (iv).
3. Dilution of acid. Concentrated H₂SO₄ mixing with water releases hydration heat. Exothermic, reject (i).
4. Condensation. Vapour to liquid gives heat back. Exothermic, reject (iii).

Why this matters. Dry ice keeps ice cream cold for hours because its endothermic sublimation steadily pulls heat from its surroundings.

Final Answer: Option (d): statements (ii) and (iv) are endothermic.

Find the free lead ion. The yellow solid is PbI₂, so the substitute must release Pb²⁺ into water freely.

Concept used. A precipitate forms when free cations and anions meet. So both reactants must dissolve to set their ions loose, and the new lead salt must be soluble.

1. What ion do we need. Lead nitrate supplies Pb²⁺. The stand-in must do the same.
2. Test lead acetate (b). It dissolves well in water, giving Pb²⁺ ions that join I- to form yellow PbI₂. It works.
3. Test lead sulphate (a). It is insoluble, so it never gives free Pb²⁺. Rejected.
4. Test (c) and (d). Ammonium nitrate and potassium sulphate have no lead, so no yellow precipitate is possible. Rejected.

Why this matters. Choosing the right reagent is a real lab skill. Picking an insoluble salt by mistake is a common reason an experiment shows no reaction.

Final Answer: Option (b): lead acetate, a soluble lead salt, can stand in for lead nitrate.

Keep oxygen out. Rancidity is air-oxidation of fats, so the storage gas must be one that will not react with the oil.

Concept used. Antioxidant packaging replaces reactive oxygen with an inert gas. Noble gases like helium and unreactive nitrogen are the standard choices.

1. Name the spoiler. Oxygen oxidises fats and oils, giving the stale smell and taste of rancidity.
2. Reject oxygen options. Options (a) and (b) both include oxygen, which would speed up spoilage. Rejected.
3. Choose the inert pair. Helium (a noble gas) and nitrogen (unreactive in normal conditions) both refuse to react with the oil, so (d) protects it.
4. Hence option (d) is the safe storage choice.

Why this matters. Nitrogen flushing is a routine industrial trick for extending the shelf life of packaged foods, dried fruits and medicines.

Final Answer: Option (d): helium or nitrogen, both inert, are used to store oils and prevent rancidity.

Two questions to answer. What type of reaction is it, and does it take in or give out heat? Settle both and the option is clear.

Concept used. Thermal decomposition breaks a compound apart using heat. Since the products carry more stored energy than the single reactant, energy must be put in, making the reaction endothermic.

1. Classify by counting. One reactant KClO₃ becomes two products, KCl and O₂. One into many is decomposition.
2. Check the energy. The mixture must be kept hot for oxygen to keep coming off. Remove the flame and it stops, so heat is being absorbed: endothermic.
3. Reject the wrong labels. It is not combination (b), not heat-releasing (c), and the trigger is heat not light, so not photochemical (d).
4. Hence option (a): endothermic decomposition.

Why this matters. This is the standard lab way to make oxygen, with a little manganese dioxide as catalyst to lower the temperature needed.

Final Answer: Option (a): a decomposition reaction that is endothermic in nature.

Hunt for a new substance. A process is a chemical reaction only if something new, with new properties, appears at the end.

Concept used. New product equals chemical change; same substance in a new state equals physical change.

1. Option (d). Shiny copper turns into black copper oxide on strong heating in air. The colour and properties change, so a new compound CuO has formed. Chemical.
2. Option (a). Compressed oxygen is still oxygen, only at higher pressure. Physical.
3. Option (b). Liquefied air is still the same gases, now as a cold liquid. Physical.
4. Option (c). Petrol evaporating is liquid turning to vapour, the same petrol. Physical. So only (d) is chemical.

Why this matters. The black coat on heated copper is the start of corrosion. Recognising it as a real reaction connects this chapter to metallurgy and rusting.

Final Answer: Option (d): only heating copper wire in air is a chemical reaction.

Match each symbol to reality. Go species by species and ask: at the reaction temperature, is this a solid, liquid or gas?

Concept used. State symbols must describe the genuine state at the conditions of the reaction. Hydrogen and oxygen are gases, and the heat of combustion turns the product water into vapour.

1. Hydrogen. H₂ is a gas; writing H₂(l) as in (a) is wrong.
2. Oxygen. O₂ is a gas; O₂(l) in (a) and (b) is wrong.
3. Water at reaction temperature. The flame is very hot, so water leaves as steam, H₂O(g), not liquid. That rules out (b) and (c).
4. Survivor. Only (d) lists gas, gas and gas correctly, so (d) is right.

Why this matters. Correct state symbols are not decoration. They tell a chemist whether to expect bubbles, a liquid or a solid, which is vital when scaling a reaction up.

Final Answer: Option (d): 2H₂(g) + O₂(g) → 2H₂O(g) carries the correct state symbols.

Count the products. A combination reaction always ends with exactly one product. Counting products on the right side sorts these in seconds.

Concept used. Combination: many reactants form one product. Decomposition: one reactant forms many products. Displacement: an element replaces another in a compound.

1. Reaction (ii). Two reactants, one product Mg(OH)₂. Combination.
2. Reaction (iii). Aluminium plus oxygen give the single product Al₂O₃. Combination.
3. Reaction (i). One KClO₃ gives two products, so it is decomposition, not combination.
4. Reaction (iv). Zinc displaces iron, giving two products by a single swap, so it is displacement. Hence only (ii) and (iii), option (d).

Why this matters. Reaction (iii), aluminium burning in oxygen, releases enormous heat. The same combination underlies the thermite reaction met in Question 20.

Final Answer: Option (d): reactions (ii) and (iii) are the combination reactions.

Pattern then balance. Spot the reaction family first, because the family hints at the products, then balance atom by atom.

Concept used. Conservation of mass forces equal atoms on each side. The reaction families seen here are combination (many into one), neutralisation and esterification (both ion-swap double displacements), and combustion (fuel plus oxygen giving CO₂ and water).

1. (a) Nitrogen and hydrogen join into one product. N₂ + 3H₂ → 2NH₃; check: 2 N and 6 H each side. Combination.
2. (b) A base and an acid neutralise. NaOH + CH₃COOH → CH₃COONa + H₂O; salt plus water, double displacement.
3. (c) Alcohol plus acid with conc. acid catalyst give an ester. C₂H₅OH + CH₃COOH → CH₃COOC₂H₅ + H₂O. Esterification, a double displacement.
4. (d) Ethene burns. Balance C (2 CO₂), then H (2 H₂O), then O (need 6 O, so 3 O₂). C₂H₄ + 3O₂ → 2CO₂ + 2H₂O, with heat and light. Combustion.

Why this matters. Reaction (a) is the Haber process that feeds the world's fertiliser industry, while (c) is how fragrant esters used in flavours and perfumes are made.

Final Answer: Balanced equations as above; types: (a) combination, (b) neutralisation, (c) esterification, (d) combustion.

Reactivity decides displacement. For (a) and (c), the more reactive element wins and shoves the weaker one out of its compound.

Concept used. An element higher in the reactivity series displaces one below it from its compound. Combination forms one product; combustion is reaction with oxygen releasing heat.

1. (a) Aluminium is more reactive than iron, so it takes the oxygen and frees molten iron. Fe₂O₃ + 2Al → Al₂O₃ + 2Fe. Displacement and redox.
2. (b) Magnesium and nitrogen combine into one solid. Balance Mg: 3Mg + N₂ → Mg₃N₂. Combination.
3. (c) Chlorine is more reactive than iodine, so it displaces iodine. 2KI + Cl₂ → 2KCl + I₂. Displacement.
4. (d) Balance ethanol burning: 2 C give 2 CO₂, 6 H give 3 H₂O, oxygen balances at 3 O₂. C₂H₅OH + 3O₂ → 2CO₂ + 3H₂O. Combustion.

Why this matters. The halogen displacement in (c) proves the reactivity order of the halogens, a fact used to extract bromine and iodine industrially.

Final Answer: Balanced equations as above; types: (a) displacement/redox, (b) combination, (c) displacement, (d) combustion/redox.

Solubility and balance. Each blank is fixed either by whether a species dissolves or by the need to balance atoms.

Concept used. State symbols record physical state: (s) solid, (aq) dissolved, (g) gas. Coefficients are fixed by conservation of mass.

1. (a) Lead iodide is the famous yellow precipitate, so it is solid: x = (s). Potassium nitrate is soluble: y = (aq).
2. (b) Copper, being more reactive, displaces silver, which deposits as solid metal. Two silver atoms are freed, so x = 2 Ag.
3. (c) Zinc sulphate dissolves, so x = (aq); the hydrogen produced bubbles off as gas, so y = (g).
4. (d) Calcium carbonate needs heat to decompose, so the missing item over the arrow is the condition heat: x = heat.

Why this matters. Filling in state symbols correctly is exactly how chemists predict whether a product will settle, dissolve or bubble away during an actual experiment.

Final Answer: (a) (s), (aq); (b) 2Ag; (c) (aq), (g); (d) Heat.

Hot or cold to touch. The fastest sort is to imagine the beaker: does it warm (exothermic) or cool (endothermic)?

Concept used. Whether a change is exothermic or endothermic depends on the net heat flow. Breaking bonds costs energy; forming bonds and hydrating ions releases energy.

1. (a) Decomposing FeSO₄ demands continuous heating, so it draws heat in. Endothermic.
2. (b) Diluting concentrated acid releases hydration heat, warming the mixture. Exothermic.
3. (c) NaOH dissolving releases strong hydration energy, so the water heats up. Exothermic.
4. (d) NH₄Cl dissolving costs more lattice energy than it gains back, so the water cools. Endothermic.

Why this matters. The cooling of dissolving ammonium chloride is the principle behind instant cold packs, while the heat of dissolving sodium hydroxide is a real safety hazard in the lab.

Final Answer: (a) endothermic, (b) exothermic, (c) exothermic, (d) endothermic.

Spot the species that grabs oxygen. The reducing agent hands its electrons over, so it ends up oxidised, usually by gaining oxygen.

Concept used. Reducing agent equals the species oxidised. Oxidation means gain of oxygen or loss of hydrogen at this level.

1. (a) Nitrogen in NH₃ ends up bonded to oxygen in NO. Ammonia gained oxygen, so NH₃ is the reducing agent.
2. (b) Fluorine is reduced from F₂ to HF, so the partner water must be oxidised. H₂O is the reducing agent.
3. (c) CO picks up an oxygen to become CO₂, so it is oxidised. CO is the reducing agent that frees iron.
4. (d) H₂ gains oxygen to form H₂O, so hydrogen is the reducing agent.

Why this matters. Carbon monoxide as a reducing agent in (c) is the heart of the blast furnace, where it pulls oxygen off iron ore to give pure iron.

Final Answer: Reducing agents: (a) ammonia; (b) water; (c) carbon monoxide; (d) hydrogen.

Find what loses oxygen. The oxidising agent ends up reduced, which at this level usually means it loses oxygen or gains hydrogen.

Concept used. Oxidising agent equals the species reduced. It hands oxygen to, or takes hydrogen or electrons from, the other reactant.

1. (a) and (d). Pb₃O₄ and V₂O₅ each lose oxygen to the other reactant, so both are reduced and act as oxidising agents.
2. (b). Free O₂ ties up with magnesium and is reduced, so oxygen is the oxidising agent.
3. (c). CuSO₄ is reduced as copper metal is set free, so it is the oxidising agent here.
4. (e) and (f). Water in (e) and CuO in (f) each give up oxygen and are reduced, so each is the oxidising agent in its reaction.

Why this matters. Identifying the oxidant in metal-extraction reactions like (d) and (f) explains how reactive metals are used to pull others out of their oxides.

Final Answer: Oxidising agents: (a) Pb₃O₄, (b) O₂, (c) CuSO₄, (d) V₂O₅, (e) H₂O, (f) CuO.

Balance metals, then ions, then check. Fix the metal atoms first, then the polyatomic ions, then count everything to confirm.

Concept used. Conservation of mass requires equal atoms on both sides. Coefficients in front of formulas are the only allowed adjustment.

1. (a) With equal moles, only one of the two carbonate oxygens reacts, giving NaCl and NaHCO₃. Atoms already balance as written.
2. (b) The bicarbonate then meets more acid: NaHCO₃ + HCl → NaCl + H₂O + CO₂. One Na, one Cl, one C and three O on each side.
3. (c) Start with copper: two Cu needed for Cu₂I₂, so 2CuSO₄. That gives two SO₄, so 2K₂SO₄, which needs four K, so 4KI. The four iodides supply two for Cu₂I₂ and two for one I₂.
4. Final: 2CuSO₄ + 4KI → Cu₂I₂ + 2K₂SO₄ + I₂, fully balanced.

Why this matters. Reactions (a) and (b) together show why baking soda fizzes with acid in stages, which is the chemistry behind antacids and baking.

Final Answer: (a) Na₂CO₃ + HCl → NaCl + NaHCO₃; (b) NaHCO₃ + HCl → NaCl + H₂O + CO₂; (c) 2CuSO₄ + 4KI → Cu₂I₂ + 2K₂SO₄ + I₂.

White solid means precipitation. The insoluble white product is the giveaway that two soluble salts have swapped ions.

Concept used. Double displacement AB + CD → AD + CB; if a product is insoluble it precipitates, so the same reaction is also called precipitation.

1. Swap the ions. K+ pairs with NO₃- and Ag+ pairs with Cl-.
2. Write the equation. KCl + AgNO₃ → AgCl + KNO₃, all atoms balanced as written.
3. Identify the precipitate. AgCl is insoluble and appears as the white solid, so it carries the (s) symbol.
4. Name the type. Two ion pairs swapped and a precipitate formed, so it is a double displacement and precipitation reaction.

Why this matters. This single reaction underpins both qualitative testing for chloride and the early chemistry of photography.

Final Answer: KCl + AgNO₃ → AgCl(whiteppt) + KNO₃; a double displacement / precipitation reaction.

Smell names the gas. The burning-sulphur odour points straight to sulphur oxides, so the green crystals must be breaking apart on heating.

Concept used. Thermal decomposition splits one compound into several using heat. Iron(II) sulphate yields a solid oxide plus two sulphur-oxide gases.

1. Heat the salt. Green FeSO₄ crystals lose their water first, then the anhydrous salt decomposes on stronger heating.
2. Write the products. A reddish-brown solid Fe₂O₃ forms, with SO₂ and SO₃ gases given off.
3. Balance it. Two formula units give the balanced line 2FeSO₄ → Fe₂O₃ + SO₂ + SO₃.
4. Name the type. One reactant, several products, driven by heat: thermal decomposition.

Why this matters. This classic activity shows decomposition in action and also show how the colour change from green to brown signals a chemical reaction.

Final Answer: 2FeSO₄ → Fe₂O₃ + SO₂ + SO₃ on heating; a thermal decomposition reaction.

Light from chemistry, not heat. A firefly's glow is energy from a reaction escaping as light instead of warmth, the reverse of a photochemical reaction.

Concept used. In bioluminescence, an oxidation reaction releases its energy as photons. An enzyme speeds up the oxidation of the protein luciferin by atmospheric oxygen.

1. The reactants. The insect's body holds luciferin and an enzyme that controls the reaction.
2. The reaction. Oxygen from air oxidises the luciferin while the enzyme makes the process happen smoothly.
3. The output. The energy of this oxidation comes out as visible light rather than heat, so the firefly glows.

Why this matters. It shows that chemical reactions can release energy as light, just as some absorb light to proceed, tying chemistry to the living world.

Final Answer: The glow is light energy released when luciferin is oxidised by air through an enzyme, a chemical (bioluminescent) reaction.

Living versus plucked. The key difference is that a grape on the vine defends itself, while a plucked grape is open to microbes.

Concept used. Fermentation is the anaerobic breakdown of sugar by microbes into ethanol and carbon dioxide, a chemical change.

1. On the plant. The grape is alive and its immune system stops microbes from fermenting it, so no change occurs.
2. After plucking. The grape can no longer defend itself, so yeasts and bacteria grow on it.
3. The condition. In the absence of air (anaerobic conditions), these microbes ferment the sugar into alcohol.
4. The verdict. A new substance, alcohol, is produced, so fermentation is a chemical change.

Why this matters. This same anaerobic fermentation is how wine, bread and curd are made, all everyday chemical changes driven by microbes.

Final Answer: Plucked grapes ferment under anaerobic conditions through microbial action; it is a chemical change.

New substance or not. Decide each case by one test: at the end, do we have a brand new substance or the original one in a new form?

Concept used. Physical change keeps the formula; chemical change makes new products with new properties.

1. (a) and (e). Evaporation of petrol and sublimation of NH₄Cl are state changes; the substance is unchanged. Physical.
2. (c). A red-hot iron rod is just hot iron; cool it and it is the same metal. Physical.
3. (b). Burning LPG makes carbon dioxide and water, new substances. Chemical.
4. (d). Curdling turns milk protein into curd, a new substance that cannot become milk again. Chemical.

Why this matters. Sorting everyday events into physical and chemical is the practical payoff of this chapter, useful from the kitchen to the laboratory.

Final Answer: Physical: (a), (c), (e); chemical: (b), (d).

Rank by reactivity. All four observations line up perfectly with each metal's place in the reactivity series, so rank them first.

Concept used. Only metals above hydrogen displace it from acid, releasing hydrogen gas. Higher reactivity means a faster, hotter reaction.

1. Silver (least reactive). Below hydrogen, so no displacement and no change.
2. Lead (mildly reactive). Just above hydrogen, so it reacts slowly, giving a few bubbles of H₂: Pb + 2HCl → PbCl₂ + H₂.
3. Aluminium (reactive). Reacts readily and the reaction is exothermic, so the temperature rises.
4. Sodium (very reactive). Reacts so fast and releases so much heat that the reaction is explosive.

Why this matters. The same reactivity order decides which metals can be safely stored, which need an inert cover, and which are extracted by displacement.

Final Answer: Behaviour matches reactivity: Ag no reaction, Pb slow with bubbles, Al exothermic warming, Na explosive.

Decode the clues. Three hints, group 2, cement, bones, all converge on calcium, and "turns red litmus blue" confirms a basic oxide.

Concept used. Metal oxides of group 2 are basic. Calcium oxide plus water gives calcium hydroxide, a base that turns red litmus blue.

1. Identify the element. A group 2 element in bones and cement is calcium.
2. Identify X. Its oxide is CaO, quick lime, widely used in cement.
3. Write the reaction. CaO + H₂O → Ca(OH)₂, the slaking of lime.
4. Confirm basicity. Ca(OH)₂ in water is lime water, a base, so it turns red litmus blue, matching the clue.

Why this matters. Recognising basic metal oxides by the litmus test is a skill carried straight into the next chapter on acids, bases and salts.

Final Answer: X is CaO; CaO + H₂O → Ca(OH)₂; the basic product turns red litmus blue.

Balance, then follow the oxygen. After balancing, track where oxygen and electrons move to pin down the reaction type.

Concept used. Double displacement swaps ion pairs; displacement frees an element; redox involves both oxidation and reduction together.

1. (a) Lead and hydrogen swap their partners, giving PbCl₂ and acetic acid. Pb(CH₃COO)₂ + 2HCl → PbCl₂ + 2CH₃COOH. Double displacement.
2. (b) Reactive sodium pushes hydrogen out of ethanol. 2Na + 2C₂H₅OH → 2C₂H₅ONa + H₂. Displacement.
3. (c) CO gains oxygen (oxidised) while Fe₂O₃ loses it (reduced). Fe₂O₃ + 3CO → 2Fe + 3CO₂. Redox.
4. (d) Sulphur loses hydrogen and oxygen is reduced to water. 2H₂S + O₂ → 2S + 2H₂O. Redox.

Why this matters. Reaction (c) is the core of iron extraction in a blast furnace, where carbon monoxide reduces ore to metal on an industrial scale.

Final Answer: (a) double displacement, (b) displacement, (c) redox, (d) redox; balanced equations as above.

Block the trigger. The decomposition needs light, so removing light stops the reaction; a dark bottle does exactly that.

Concept used. Photochemical decomposition uses light energy to split a compound. AgCl breaks into silver and chlorine when light falls on it.

1. The unwanted reaction. 2AgCl → 2Ag + Cl₂ happens whenever light reaches the solid.
2. The visible spoilage. Freed silver makes the white salt turn grey, so the reagent is no longer pure.
3. The fix. A dark coloured bottle stops light from reaching the salt, so no decomposition occurs.
4. Hence silver chloride, and other light-sensitive silver salts, are always stored in dark bottles.

Why this matters. The same light sensitivity is harnessed deliberately in photography but must be prevented when the pure reagent is to be stored.

Final Answer: Light decomposes AgCl to silver and chlorine, so dark bottles are used to keep light out and prevent photochemical decomposition.

Balance first, classify second. Get equal atoms on both sides, then count products to name combination versus decomposition, or look for a freed element to spot displacement.

Concept used. Conservation of mass sets the coefficients. Combination ends with one product; decomposition starts with one reactant; displacement frees an element.

1. (a) and (c) and (e). Each ends with a single product (MgCl₂, Na₂S, CaSiO₃). After balancing, all three are combination reactions.
2. (b) and (f). Each starts from one compound and gives two products: 2HgO → 2Hg + O₂ and 2H₂O₂ → 2H₂O + O₂. Both are decompositions.
3. (d). Magnesium displaces titanium from TiCl₄. Balance chlorine: TiCl₄ + 2Mg → Ti + 2MgCl₂. Displacement.
4. Final classification follows directly from how many products and which element is set free.

Why this matters. Reaction (d) is the Kroll process used to extract titanium metal, prized for being strong and light in aircraft and implants.

Final Answer: (a) combination, (b) decomposition, (c) combination, (d) displacement, (e) combination, (f) decomposition.

Two gases, two products. Magnesium is reactive enough to combine with both oxygen and nitrogen, giving two different compounds.

Concept used. Magnesium forms an oxide with O₂ and a nitride with N₂. A basic metal oxide reacts with water to give a hydroxide.

1. Compound X. In oxygen, 2Mg + O₂ → 2MgO. White MgO forms with bright light.
2. Compound Y. In nitrogen, 3Mg + N₂ → Mg₃N₂. The ribbon keeps burning to give magnesium nitride.
3. X in water. MgO is a basic oxide, so it reacts with water: MgO + H₂O → Mg(OH)₂.
4. This gives the formulae X = MgO, Y = Mg₃N₂, and the hydroxide-forming equation.

Why this matters. It shows that a magnesium fire cannot be put out by smothering with nitrogen or carbon dioxide, because the metal reacts with both, an important fire-safety lesson.

Final Answer: X = MgO, Y = Mg₃N₂; MgO + H₂O → Mg(OH)₂.

Above or below hydrogen. The reactivity series settles it: only metals placed above hydrogen can push it out of an acid.

Concept used. A more reactive metal displaces a less reactive one. Hydrogen sits in the series, so metals above it release H₂ from acids, while those below do not.

1. Zinc's position. Zinc is above hydrogen, so it is more reactive and replaces hydrogen from HCl.
2. Zinc's reaction. Zn + 2HCl → ZnCl₂ + H₂, with bubbles of hydrogen gas.
3. Copper's position. Copper is below hydrogen, so it is less reactive and cannot displace hydrogen.
4. Copper's result. No reaction with dilute HCl, so no gas is seen.

Why this matters. This is exactly why copper, being below hydrogen, is safe for water pipes and vessels, while reactive metals would corrode in contact with acids.

Final Answer: Zinc (above hydrogen) reacts: Zn + 2HCl → ZnCl₂ + H₂; copper (below hydrogen) does not react.

Name the attacker. The black layer means the silver has reacted with something in the air, and the usual culprit is hydrogen sulphide.

Concept used. Corrosion is the slow reaction of a metal with air and moisture. Silver reacts with airborne H₂S to make black silver sulphide.

1. Why it blackens. Traces of H₂S in the air react with the silver surface over a few days.
2. Name the phenomenon. This gradual attack on the metal is corrosion, specifically the tarnishing of silver.
3. Name the product. The black coat is silver sulphide, Ag₂S.
4. Why polishing works. Toothpaste, a mild abrasive, scrapes off the thin Ag₂S layer to reveal bright silver again.

Why this matters. Understanding corrosion explains everyday care of metals, from polishing silverware to painting iron gates to prevent rust.

Final Answer: (a) corrosion, silver reacting with H₂S in air; (b) the black substance is silver sulphide, Ag₂S.

Decompose, then read the gas. The four-part answer follows once you write the decomposition and recognise the brown gas as NO₂.

Concept used. Thermal decomposition splits one compound into many on heating. The acidic non-metal oxide NO₂ dissolves to give a solution of pH below 7.

1. Balance the breakdown. Two units of copper nitrate give two CuO, four NO₂ and one O₂: 2Cu(NO₃)₂ → 2CuO + 4NO₂ + O₂.
2. Spot the brown gas. Nitrogen dioxide NO₂ is the familiar reddish-brown gas, so X is NO₂.
3. Classify. One reactant into several products on heating, a clear thermal decomposition.
4. Predict the pH. As a non-metal oxide, NO₂ forms acids in water, so the solution has pH less than 7.

Why this matters. NO₂ dissolving in rain to form acids is a cause of acid rain, linking this lab reaction to a real environmental problem.

Final Answer: (a) 2Cu(NO₃)₂ → 2CuO + 4NO₂ + O₂; (b) NO₂; (c) thermal decomposition; (d) pH below 7.

One test, one observation. Each gas has a signature response; match the gas to its single most reliable test.

Concept used. Gas tests rely on a quick chemical or physical response: precipitation, colour change, support of burning, or a sound.

1. (a) CO₂. The milky lime water comes from CaCO₃ forming. Ca(OH)₂ + CO₂ → CaCO₃ + H₂O.
2. (b) SO₂. Being a reducing agent, it decolourises purple acidified KMnO₄.
3. (c) O₂. It does not burn itself but makes a glowing splint flare up, since it supports combustion.
4. (d) H₂. It is flammable, so a lit splint makes it burn with a squeaky pop.

Why this matters. Reliable gas identification is the first step in countless experiments, from testing the products of a reaction to checking the purity of a gas.

Final Answer: (a) lime water turns milky; (b) decolourises acidified KMnO₄; (c) rekindles a glowing splint; (d) pop sound with a burning splint.

Compare reactivities first. For each part, compare the added metal with the metal (or hydrogen) it might displace, then decide.

Concept used. Displacement happens only when the added metal is more reactive than the species it would replace. The reactivity series ranks them.

1. (a) Zinc above copper, so it displaces copper and the blue solution turns colourless. Zn + CuSO₄ → ZnSO₄ + Cu.
2. (b) Aluminium above hydrogen, so it displaces hydrogen gas. 2Al + 6HCl → 2AlCl₃ + 3H₂.
3. (c) Silver below copper, so it cannot displace copper. No reaction is seen.
4. Each result follows directly from the reactivity comparison.

Why this matters. These displacement rules explain how metals are extracted and why iron buckets cannot store copper sulphate solution, a common practical caution.

Final Answer: (a) Zn + CuSO₄ → ZnSO₄ + Cu, blue fades; (b) 2Al + 6HCl → 2AlCl₃ + 3H₂; (c) no reaction.

Most acids give hydrogen, nitric is special. Treat the two common acids and NaOH normally, then handle nitric acid as the oxidising exception.

Concept used. Zinc above hydrogen displaces H₂ from common dilute acids. Nitric acid oxidises that hydrogen, and zinc, being amphoteric, also reacts with alkalis.

1. Sulphuric and hydrochloric. Both give H₂: Zn + H₂SO₄ → ZnSO₄ + H₂ and Zn + 2HCl → ZnCl₂ + H₂.
2. Nitric acid. As an oxidising acid it converts the hydrogen to water and gives N₂O: 4Zn + 10HNO₃ → 4Zn(NO₃)₂ + 5H₂O + N₂O.
3. Sodium chloride. Sodium is more reactive than zinc, so zinc cannot displace it; no reaction.
4. Sodium hydroxide. Amphoteric zinc reacts: Zn + 2NaOH → Na₂ZnO₂ + H₂, releasing hydrogen.

Why this matters. Zinc's amphoteric behaviour, reacting with both acids and alkalis, is a key property used in galvanising and in distinguishing it from many other metals.

Final Answer: H₂ with H₂SO₄, HCl and NaOH; N₂O with HNO₃; no reaction with NaCl.

Precipitate then dissolve. Part (c) is the giveaway: a precipitate that vanishes in acid must be the salt of a weak acid.

Concept used. Precipitation is double displacement giving an insoluble solid. A salt of a weak acid reacts with a strong acid, which sets the weak acid (here SO₂ plus water) free.

1. Form the precipitate. Na₂SO₃ + BaCl₂ → BaSO₃ + 2NaCl. The white solid is BaSO₃.
2. Name the reaction. Two ion pairs swapped, so it is also a double displacement reaction.
3. Add acid. Dilute HCl, a strong acid, attacks the sulphite: BaSO₃ + 2HCl → BaCl₂ + H₂O + SO₂.
4. Explain the disappearance. BaCl₂ is soluble and SO₂ escapes as gas, so the solid BaSO₃ dissolves and the precipitate is gone.

Why this matters. The acid test cleanly distinguishes a sulphite from a sulphate: BaSO₃ dissolves in HCl with a smell of burning sulphur, but BaSO₄ does not dissolve at all.

Final Answer: (a) Na₂SO₃ + BaCl₂ → BaSO₃ + 2NaCl; (b) a double displacement reaction; (c) HCl converts BaSO₃ to soluble BaCl₂ and SO₂, dissolving the precipitate.

Container must not react. For each solution, ask whether the container metal would react with it; if yes, it cannot be stored there.

Concept used. A metal holds a solution safely only when it is less reactive than the dissolved species. The exception is aluminium with nitric acid, which forms a protective oxide coat.

1. Copper with the four solutions. Below hydrogen, copper ignores dilute HCl, ZnCl₂ and water, so all three are safe. Dilute HNO₃ attacks copper, so it is unsafe.
2. Aluminium with HCl. Aluminium displaces hydrogen: 2Al + 6HCl → 2AlCl₃ + 3H₂, so HCl cannot be kept.
3. Aluminium with ZnCl₂. Aluminium is above zinc, so it displaces zinc: 2Al + 3ZnCl₂ → 2AlCl₃ + 3Zn; not safe.
4. Aluminium with HNO₃ and water. Nitric acid forms a protective Al₂O₃ film, and water does not react, so both can be stored in aluminium.

Why this matters. Choosing the right container for a chemical is a practical safety decision in labs and industry, and the reactivity series is the tool that guides it.

Final Answer: Copper: dilute HCl, ZnCl₂, H₂O. Aluminium: dilute HNO₃, H₂O.