(C) Magnesium oxide

When magnesium ribbon is burned in air, it reacts with oxygen to form magnesium oxide: Magnesium + Oxygen → Magnesium oxide.

Source: Chapter 1, Section 1.1; Chapter 3, Section 3.2.1

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• The textbook explicitly states: "when a magnesium ribbon is burnt in oxygen, it gets converted to magnesium oxide."
• Air contains oxygen, so burning in air produces the same oxide product.
• Magnesium nitride (A) can form in pure nitrogen, but in air the dominant reaction is with oxygen. The textbook does not mention nitride formation here — always go with what the prescribed text states.
• For a 1-mark MCQ, simply identifying the correct option with a one-line justification is sufficient.

(C) Combination reaction

In this reaction, two reactants (CaO and H₂O) combine to form a single product Ca(OH)₂, which is the definition of a combination reaction.

Source: Chapter 1, Section 1.2.1 Combination Reaction

The key indicator here is one product formed from two reactants. The textbook directly uses this equation as the example for combination reactions. Also note it is exothermic (releases heat), but that does not change its classification as a combination reaction. Eliminate other options: no single substance breaks down (not decomposition), no element replaces another (not displacement), and no ion exchange occurs (not double displacement).

(B) Nitrogen dioxide

When lead nitrate is heated, brown fumes of nitrogen dioxide (NO₂) are produced: $2\text{Pb(NO}_3)_2 \xrightarrow{\Delta} 2\text{PbO} + 4\text{NO}_2 + \text{O}_2$.

Source: Chapter 1, Section 1.2.2, Activity 1.6

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The key fact is that the brown fumes specifically identify NO₂. Although PbO and O₂ are also products of the same reaction, they are not brown fumes — PbO is a yellow solid and O₂ is a colourless gas. Examiners test whether students can distinguish between all the products of this reaction and correctly identify which one appears as brown fumes.

(C) Na₂SO₄(aq) + BaCl₂(aq) → BaSO₄(s) + 2NaCl(aq)

This is a precipitation reaction because an insoluble white precipitate of BaSO₄ is formed when two aqueous solutions are mixed.

Source: Chapter 1, Section 1.2.4 Double Displacement Reaction

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• A precipitation reaction produces an insoluble solid (precipitate) from two aqueous solutions. The (s) symbol after BaSO₄ is the key indicator in the equation.
• Option A is a combination reaction; B is a combination/oxidation reaction; D is a displacement reaction.
• Examiners expect you to identify the precipitate (BaSO₄) as the reason for your choice.

(A) Oxidised

In ZnO + C → Zn + CO, carbon gains oxygen (from ZnO) to form CO. Gain of oxygen means oxidation. So carbon is oxidised.

The key concept: oxidation = gain of oxygen; reduction = loss of oxygen. Here, ZnO loses oxygen (ZnO is reduced → Zn), and C gains that oxygen (C is oxidised → CO). Carbon acts as the reducing agent but itself gets oxidised — a common exam confusion point. Remember: the substance that gets oxidised is the reducing agent.

(B) Black and white photography

Silver chloride (AgCl) decomposes into silver metal and chlorine gas in sunlight — this photochemical reaction is the basis of black and white photography, where silver deposits form the image. This is a standard example from Chapter 1/Chapter 3 of NCERT Science Class 10.

(B) 3Fe + 4H₂O → Fe₃O₄ + 4H₂

Balancing by hit-and-trial: oxygen needs coefficient 4 on H₂O (to match 4 O in Fe₃O₄), giving 8 H atoms, so 4H₂ on RHS. Then Fe must be 3 on LHS to match Fe₃O₄. All atoms balance: Fe=3, H=8, O=4 on both sides. Source: Chapter 1, Section 1.1.2.

(C) It releases energy in the form of heat

During respiration, glucose (organic compound) is broken down, releasing energy. Since energy is released in the process, it is classified as an exothermic reaction.

The key link is between two chapters: Ch. 1 defines exothermic reactions as those in which heat/energy is given out, and Ch. 5 states that during respiration, organic compounds like glucose are broken down to provide energy. Option C correctly combines both ideas. Options A and D describe products/reactants of respiration, not why it is exothermic. Option B describes photosynthesis, not respiration.

(B) Iron displaces copper from copper sulphate, forming iron sulphate: $\text{Fe}(s) + \text{CuSO}_4(aq) \rightarrow \text{FeSO}_4(aq) + \text{Cu}(s)$. Since FeSO₄ is pale green, the blue colour fades.

Source: Chapter 1, Section 1.2.3 Displacement Reaction

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This is a displacement reaction — iron (more reactive) displaces copper from CuSO₄. The blue colour is due to Cu²⁺ ions; as they are replaced by Fe²⁺ (iron sulphate, which is light green/colourless), the blue fades. Options C and D are wrong reaction types; Option A is incorrect because no blue precipitate forms.

(C) Corrosion

The green coating on copper is basic copper carbonate, formed when copper reacts with moist CO₂ in air. This is an example of corrosion.

Source: Chapter 1, Section 1.3.1; Chapter 3, Section 3.5

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• Corrosion is the attack on a metal by substances like moisture or acids in its surroundings. The green coat on copper and black coat on silver are classic CBSE examples of corrosion.
• Rancidity refers to spoilage of fats/oils — eliminate it immediately.
• Reduction is loss of oxygen — the opposite of what happens here.
• Precipitation is formation of an insoluble salt in solution — not applicable here.
• Remember: rusting of iron, blackening of silver, and greening of copper are all examples of corrosion.

(B) Prevent the chips from getting oxidised

Nitrogen is an inert gas. It prevents the chips from reacting with oxygen, which would cause oxidation (rancidity) of fats and spoil the chips.

Nitrogen is used because it is chemically inert and does not react with the oils/fats in chips. Oxygen causes oxidation, making chips rancid and stale. This concept is linked to Chapter 1 (prevention of oxidation/rancidity). Option D is wrong — CO₂ is not the problem gas; Option C is the opposite of what nitrogen does; Option A is incorrect as nitrogen has no effect on crunchiness.

(D) CuO

In the reaction, CuO loses oxygen to form Cu, so CuO is reduced. (Reduction = loss of oxygen.)

Source: Chapter 1, Section 1.2.5 Oxidation and Reduction

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Examiners expect you to recall the definition: reduction = loss of oxygen. In this reaction, CuO → Cu (oxygen is removed), so CuO is reduced. H₂ gains oxygen to form H₂O, so H₂ is oxidised. This is a classic redox (oxidation-reduction) reaction example directly from the textbook (equation 1.29/1.30).

(C) Photochemical decomposition

In this reaction, AgCl decomposes into Ag and Cl₂ using sunlight (light energy), making it a photochemical decomposition reaction.

Source: Chemical Reactions and Equations, Chapter 1

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• Key criterion: Identify the energy source causing decomposition. Here, sunlight (light energy) drives the reaction → photochemical decomposition.
• Thermal decomposition uses heat; electrolytic decomposition uses electricity; this reaction uses light.
• This is a classic NCERT example often cited in Exercise Q12 (decomposition using light energy).
• For MCQs, just state the correct option and give a one-line reason — no elaboration needed.

(B) It indicates the substance is dissolved in water.

The notation (aq) stands for aqueous, meaning the substance is present as a solution in water.

Source: Chapter 1, Section 1.1.2 (Writing Symbols of Physical States)

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The textbook explicitly states: "The word aqueous (aq) is written if the reactant or product is present as a solution in water." The four state symbols to remember are: (s) – solid, (l) – liquid, (g) – gas, (aq) – dissolved in water. 'aq' has nothing to do with acidity or insolubility — those are common traps in this MCQ.

(A) Both A and R are true and R is the correct explanation of A.

FeSO₄·7H₂O crystals first lose water of crystallisation on heating, then thermally decompose to give Fe₂O₃, SO₂ and SO₃. R correctly explains the two-step process described in A.

Source: Chapter 1, Section 1.2.2 Decomposition Reaction

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The textbook explicitly states that ferrous sulphate crystals (FeSO₄·7H₂O) first lose water of crystallisation, then decompose into Fe₂O₃, SO₂, and SO₃. Both steps are caused by heat — first breaking the hydration, then breaking the compound itself. So R (water of crystallisation + heat providing energy for decomposition) directly and correctly explains A. Option (A) is the right choice.

(A) Both A and R are true and R is the correct explanation of A.

A chemical equation must be balanced because the law of conservation of mass states that total mass of reactants equals total mass of products, ensuring equal number of atoms on both sides.

Source: Chapter 1, Section 1.1.2 Balanced Chemical Equations

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• A is true: A balanced equation is essential for stoichiometric calculations so that correct mole ratios are used.
• R is true: The law of conservation of mass is correctly stated.
• R correctly explains A: We balance equations specifically because of this law — atoms (and hence mass) must be conserved. The textbook explicitly states: "the total mass of the elements present in the products has to be equal to the total mass of the elements present in the reactants... Hence, we need to balance a skeletal chemical equation."
• Option (B) would be wrong here because R is not just true but is the direct reason why balancing is necessary.

(A) Both A and R are true and R is the correct explanation of A.

Storing food in airtight containers limits oxygen exposure, which directly prevents oxidation of fats and oils — the cause of rancidity. Hence, R correctly explains A.

Rancidity is caused by oxidation of fats/oils in food. Airtight containers reduce oxygen contact, slowing this oxidation. Since R directly and correctly explains why A is true, option (A) is the right choice. Remember: whenever the reason logically and correctly justifies the assertion, choose option (A).

When copper powder is heated in air, oxygen is added to it, forming black copper(II) oxide.

Chemical equation:

$$2\text{Cu} + \text{O}_2 \xrightarrow{\text{Heat}} 2\text{CuO}$$

Type of reaction: It is an oxidation reaction (copper gains oxygen). It can also be classified as a combination reaction (two substances combine to form one product).

Source: Chapter 1, Section 1.2.5 Oxidation and Reduction

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• Examiners expect the balanced equation with the heat condition arrow — missing this loses a mark.
• Two valid type-labels are acceptable: oxidation (because Cu gains oxygen) and combination (two reactants → one product). Writing either or both is fine; writing both shows depth.
• The key term is copper(II) oxide — "copper oxide" alone may be accepted but the IUPAC name is preferred.

Magnesium ribbon is cleaned with sandpaper before burning because its surface is covered with a thin layer of magnesium oxide (formed due to reaction with atmospheric oxygen). This oxide layer prevents the ribbon from burning properly. Removing it with sandpaper exposes the pure shiny metal, allowing it to burn readily in air.

Source: Chapter 3, Section 3.2.1

The key point examiners look for is: (1) identifying that a pre-existing oxide layer forms on the surface, and (2) stating that this layer must be removed to allow proper burning/reaction. Many students only mention "to clean it" — you must explain why cleaning is necessary (the oxide layer blocks the reaction). The source explicitly states that "surfaces of metals such as magnesium… are covered with a thin layer of oxide" which acts as a protective coat.

When zinc granules are added to dilute sulphuric acid, the following observations indicate a chemical reaction has taken place:

1. Bubbles of gas (hydrogen) are evolved on the surface of the zinc granules.
2. The zinc granules gradually dissolve (decrease in size), showing that a new substance, zinc sulphate, is being formed.

$$\text{Zn}(s) + \text{H}_2\text{SO}_4(aq) \rightarrow \text{ZnSO}_4(aq) + \text{H}_2(g)$$

Source: Chapter 1, Section 1.1.2

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The examiner expects two distinct, observable signs of a chemical reaction — not just the equation. The two most commonly accepted answers are: (1) evolution of hydrogen gas (bubbles) and (2) dissolution/disappearance of zinc granules. A third valid option is that the reaction mixture may become warm (heat is released), indicating an exothermic reaction. Always name the gas produced for full marks. Writing the equation is a bonus but is not mandatory for a 2-mark very short answer.

A skeletal chemical equation is an unbalanced chemical equation where the reactants and products are represented by their chemical formulae but the number of atoms of each element is not equal on both sides. For example: Mg + O₂ → MgO is a skeletal equation.

The Law of Conservation of Mass makes it necessary to balance such an equation. It states that mass can neither be created nor destroyed in a chemical reaction, so the number of atoms of each element must be the same on both sides.

Source: Chapter 1, Sections 1.1.1 and 1.1.2

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• 1 mark for defining skeletal equation (unbalanced, formulae written but atoms unequal on both sides).
• 1 mark for naming and stating the Law of Conservation of Mass.
• Examiners expect the law to be named, not just described. A brief example (like Mg + O₂ → MgO) strengthens the definition but is not compulsory.

Type of reaction: Combination reaction

Balanced chemical equation:

$$\text{Ca(OH)}_2(aq) + \text{CO}_2(g) \rightarrow \text{CaCO}_3(s) + \text{H}_2\text{O}(l)$$

Observable change: A white precipitate of calcium carbonate (CaCO₃) is formed, turning the clear lime water solution milky/cloudy.

Source: Chapter 1, Section 1.2.1 Combination Reaction

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• The examiner expects you to correctly name the reaction type (combination — two reactants form one main product), give the balanced equation with state symbols, and state one visible change.
• The white precipitate/milky appearance is the key observable here — examiners specifically look for this.
• Note: excess CO₂ does not convert CaCO₃ further in the NCERT Class 10 context; the equation given in the textbook (1.14) is the expected answer.
• Always include state symbols (aq), (g), (s), (l) — they can fetch you marks.

Substance oxidised: HCl (hydrochloric acid)
Substance reduced: MnO₂ (manganese dioxide)

Reason: In HCl, hydrogen is oxidised — it loses electrons/hydrogen is removed to form Cl₂. MnO₂ gains hydrogen (is reduced) to form MnCl₂ — it loses oxygen. Thus MnO₂ is the oxidising agent and HCl is the reducing agent.

Examiners expect you to name both substances and give a reason based on gain/loss of oxygen or hydrogen. MnO₂ loses oxygen → gets reduced; HCl loses hydrogen (or Cl⁻ gets oxidised to Cl₂) → HCl is oxidised. Always state which is oxidised and which is reduced — one line each — then one reason line. This question tests Section 1.2.5 (oxidation-reduction) concepts.

Corrosion is the process by which a metal is attacked by substances in its surroundings such as moisture and acids, leading to its deterioration.

Chemical formula of rust: $\text{Fe}_2\text{O}_3 \cdot x\text{H}_2\text{O}$ (hydrated iron oxide)

Conditions required for rusting: Both air (oxygen) and water (moisture) must be present simultaneously. If iron is exposed to only water or only dry air, it does not rust.

Harmful consequence: Corrosion causes severe damage to iron structures such as bridges, railings, car bodies and ships, leading to enormous financial loss every year.

Source: Chapter 1, Section 1.3.1; Chapter 3, Section 3.5

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• Examiners expect the chemical formula of rust — $\text{Fe}_2\text{O}_3 \cdot x\text{H}_2\text{O}$ — to be written correctly. Many students lose marks by omitting it.
• The key experimental finding (Activity 3.14) is that both air and water are needed — state this clearly, not just one condition.
• For the harmful consequence, any one specific example (bridges, ships, etc.) with mention of financial/structural damage is sufficient for 1 mark.

Displacement Reaction: A reaction in which a more reactive element displaces a less reactive element from its compound.

Example: Iron displaces copper from copper sulphate solution.

$$\text{Fe}(s) + \text{CuSO}_4(aq) \rightarrow \text{FeSO}_4(aq) + \text{Cu}(s)$$

Double Displacement Reaction: A reaction in which two different ions are exchanged between the reactants to form two new compounds.

Example: Sodium sulphate reacts with barium chloride to form a white precipitate of barium sulphate.

$$\text{Na}_2\text{SO}_4(aq) + \text{BaCl}_2(aq) \rightarrow \text{BaSO}_4(s) + 2\text{NaCl}(aq)$$

Source: Chapter 1, Sections 1.2.3 and 1.2.4

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• 3 marks split: ~1 mark for each definition + example + equation (examiners often award ½ mark for definition, ½ for equation per reaction type).
• The key distinction: in displacement, one element pushes out another; in double displacement, two ions exchange between compounds.
• Always balance the equations — unbalanced equations lose marks.
• The BaSO₄ precipitate (white, insoluble) is a classic double displacement / precipitation example straight from the textbook Activity 1.10.

(i) $\text{HNO}_3 + \text{Ca(OH)}_2 \rightarrow \text{Ca(NO}_3)_2 + \text{H}_2\text{O}$

Count: N=1 LHS, N=2 RHS; balance by placing 2 before HNO₃ and 2 before H₂O:

$$2\text{HNO}_3 + \text{Ca(OH)}_2 \rightarrow \text{Ca(NO}_3)_2 + 2\text{H}_2\text{O}$$

(ii) $\text{NaOH} + \text{H}_2\text{SO}_4 \rightarrow \text{Na}_2\text{SO}_4 + \text{H}_2\text{O}$

Na=1 LHS, Na=2 RHS; balance by placing 2 before NaOH and 2 before H₂O:

$$2\text{NaOH} + \text{H}_2\text{SO}_4 \rightarrow \text{Na}_2\text{SO}_4 + 2\text{H}_2\text{O}$$

(iii) $\text{BaCl}_2 + \text{H}_2\text{SO}_4 \rightarrow \text{BaSO}_4 + \text{HCl}$

Cl=2 LHS, Cl=1 RHS; balance by placing 2 before HCl:

$$\text{BaCl}_2 + \text{H}_2\text{SO}_4 \rightarrow \text{BaSO}_4 + 2\text{HCl}$$

Source: Chapter 1, Section 1.1.2 Balanced Chemical Equations

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• Key rule: Use coefficients only — never change the subscripts inside a formula.
• For each equation, count atoms of every element on both sides and adjust coefficients to make them equal (hit-and-trial method).
• A common mistake is changing H₂O to H₂O₂ to balance oxygen — this is wrong.
• Each sub-question is 1 mark; examiners check that all atoms are balanced on both sides. Writing unbalanced equations scores zero even if the formulae are correct.

Balanced Chemical Equation:

$$\text{CaCO}_3(s) \xrightarrow{\Delta} \text{CaO}(s) + \text{CO}_2(g)$$

Type of Decomposition: It is a thermal decomposition reaction, as the compound is broken down by the application of heat. Since energy is absorbed, it is also an endothermic reaction.

Industrial Use of Product (CaO / Quick Lime): Calcium oxide (quick lime) is used in the manufacturing of cement and glass, and also for whitewashing — it reacts with water to form slaked lime [Ca(OH)₂], which is used to coat walls.

Source: Chapter 1, Section 1.2 & What You Have Learnt

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• Equation (1 mark): Must be balanced with correct state symbols and a heat arrow (Δ). CaCO₃ → CaO + CO₂ is already balanced as written.
• Type (1 mark): "Thermal decomposition" is the specific type; adding "endothermic" strengthens the answer. Examiners expect both terms.
• Industrial use (1 mark): The textbook links CaO (quick lime) to slaked lime and whitewashing (Section 1.2.1). Mentioning cement/glass manufacturing also earns the mark. Any one valid industrial use is sufficient.

(i) The yellow precipitate formed is lead(II) iodide (PbI₂).

(ii) Balanced chemical equation with state symbols:

$$\text{Pb(NO}_3)_2\text{(aq)} + 2\text{KI(aq)} \rightarrow \text{PbI}_2\text{(s)} + 2\text{KNO}_3\text{(aq)}$$

(iii) This is a double displacement reaction (also called a precipitation reaction), because the ions of the two reactants exchange places and an insoluble precipitate (PbI₂) is formed.

Source: Chapter 1, Section 1.2.4 Double Displacement Reaction

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• (i) Just naming PbI₂ with its correct IUPAC name earns the mark.
• (ii) Examiners check: correct formulae, coefficient 2 before KI and KNO₃, and all four state symbols — (aq), (aq), (s) for the precipitate, (aq). Missing state symbols or unbalanced equation loses marks.
• (iii) Both terms ("double displacement" and "precipitation") are acceptable; mentioning both is safer. The key reasoning is ion exchange + precipitate formation. The textbook explicitly identifies this lead iodide reaction as a double displacement/precipitation reaction.

(i) Potassium + Water → Potassium hydroxide + Hydrogen

$$2\text{K}(s) + 2\text{H}_2\text{O}(l) \rightarrow 2\text{KOH}(aq) + \text{H}_2(g)$$

(ii) Barium chloride + Aluminium sulphate → Barium sulphate + Aluminium chloride

$$3\text{BaCl}_2(aq) + \text{Al}_2(\text{SO}_4)_3(aq) \rightarrow 3\text{BaSO}_4(s) + 2\text{AlCl}_3(aq)$$

(iii) Hydrogen sulphide + Oxygen → Water + Sulphur dioxide

$$2\text{H}_2\text{S}(g) + 3\text{O}_2(g) \rightarrow 2\text{H}_2\text{O}(l) + 2\text{SO}_2(g)$$

Source: Chapter 1, Section 1.1.2 (Balanced Chemical Equations)

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• Each sub-question is 1 mark: award is for the correctly balanced equation with proper state symbols.
• Key state symbols: metals/solids → (s); gases → (g); solutions in water → (aq); liquids → (l). BaSO₄ is insoluble, so (s); KOH dissolves, so (aq).
• Balancing tip: always check atom count on both sides before writing the final equation.
• Do not forget to write the word equation before the chemical equation — examiners appreciate it, and it helps you set up the formula equation correctly.

In the reaction ZnO + C → Zn + CO, oxidation and reduction happen at the same time (simultaneously) because one substance loses oxygen while another gains it in the same reaction.

• Carbon (C) is oxidised — it gains oxygen (from ZnO) to form CO.
• Zinc oxide (ZnO) is reduced — it loses oxygen to form Zn.

Since one substance cannot gain oxygen unless another loses it, oxidation and reduction always occur together. Such reactions are called redox reactions.

Source: Chapter 1, Oxidation and Reduction (gain/loss of oxygen)

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• Examiners expect you to clearly state the definition: oxidation = gain of oxygen; reduction = loss of oxygen.
• You must name the specific substance oxidised and reduced with a reason (what oxygen did) — this earns separate marks.
• The key phrase to include is "simultaneously" or "at the same time," showing you understand why it's called a redox reaction.
• Do not mix up ZnO and C — C is oxidised, ZnO (or Zn) is reduced.

Chemical Reaction: A chemical reaction is a process in which new substances with new properties are formed due to the breaking and making of bonds between atoms. Whenever a chemical change occurs, a chemical reaction has taken place.

Evidence 1 – Evolution of a Gas (Activity 1.3):
Take a few zinc granules in a conical flask. Add dilute sulphuric acid. Bubbles of hydrogen gas are observed around the zinc granules, indicating a chemical reaction has occurred.

$$\text{Zn} + \text{H}_2\text{SO}_4 \rightarrow \text{ZnSO}_4 + \text{H}_2\uparrow$$

Evidence 2 – Change in Colour (Activity 1.2):
Take lead nitrate solution in a test tube. Add potassium iodide solution. A bright yellow precipitate of lead iodide is immediately formed, showing a change in colour.

$$\text{Pb(NO}_3)_2 + 2\text{KI} \rightarrow \text{PbI}_2 + 2\text{KNO}_3$$

Source: Chapter 1, Chemical Reactions and Equations, Section 1.1 and Activities 1.2 & 1.3

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• Definition (1 mark): State that new substances are formed; mention breaking and making of bonds.
• Activity for gas evolution (2 marks): Describe the zinc + acid activity clearly; write the balanced equation with ↑ for gas.
• Activity for colour change (2 marks): Describe the lead nitrate + potassium iodide activity; write the balanced equation. The yellow precipitate of PbI₂ is the key observation.
• Examiners award marks for: correct activity description, correct observation, and correctly balanced equations. Always balance equations — an unbalanced equation loses a mark.

(a) Reactions in which heat is released along with the products are called exothermic reactions.
Example: Burning of natural gas —
$$\text{CH}_4(g) + 2\text{O}_2(g) \rightarrow \text{CO}_2(g) + 2\text{H}_2\text{O}(g) + \text{Heat}$$

Reactions in which energy is absorbed are called endothermic reactions.
Example: Decomposition of water by electricity —
$$2\text{H}_2\text{O}(l) \xrightarrow{\text{electricity}} 2\text{H}_2(g) + \text{O}_2(g)$$

(b) When quick lime is added to water, heat is released. This is an exothermic reaction. It is also a combination reaction since two substances form a single product.
$$\text{CaO}(s) + \text{H}_2\text{O}(l) \rightarrow \text{Ca(OH)}_2(aq) + \text{Heat}$$

(c) In the decomposition of water by electricity, energy (electrical energy) is absorbed to break water into hydrogen and oxygen. Since energy is absorbed, it is classified as an endothermic reaction. The type of decomposition involved is electrolytic decomposition.

Source: Chapter 1 — Chemical Reactions and Equations, Section 1.2.1 Combination Reaction

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• (a) Always give the definition + a balanced equation for both. Examiners award 1 mark each for correct definitions and 1 mark for each balanced equation.
• (b) Note that this reaction is both a combination AND an exothermic reaction — mentioning both scores full marks.
• (c) The key logic is: energy is supplied/absorbed → endothermic. The specific type is electrolytic decomposition (energy supplied as electricity). Don't write "electrical decomposition."

(a) Oxidation is the gain of oxygen or loss of hydrogen by a substance.
Reduction is the loss of oxygen or gain of hydrogen by a substance.
Example of oxidation: $2\text{Cu} + \text{O}_2 \xrightarrow{\Delta} 2\text{CuO}$ (copper gains oxygen)
Example of reduction: $\text{CuO} + \text{H}_2 \xrightarrow{\Delta} \text{Cu} + \text{H}_2\text{O}$ (CuO loses oxygen)

(b) The balanced chemical equation for recovery of silver is:

$$2\text{AgNO}_3(aq) + \text{Cu}(s) \rightarrow \text{Cu(NO}_3)_2(aq) + 2\text{Ag}(s)$$

Type of reaction: Displacement reaction (copper displaces silver from silver nitrate solution).

(c) Metal X is copper (Cu). The black oxide formed is copper(II) oxide (CuO).

Reaction 1 (heating in air):
$$2\text{Cu} + \text{O}_2 \xrightarrow{\Delta} 2\text{CuO}$$

Reaction 2 (hydrogen passed over CuO):
$$\text{CuO} + \text{H}_2 \xrightarrow{\Delta} \text{Cu} + \text{H}_2\text{O}$$

In Reaction 2: H₂ is oxidised (gains oxygen); CuO is reduced (loses oxygen). This is a redox reaction.

Source: Chapter 1, Section 1.2.5 Oxidation and Reduction

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• (a) Always give both definitions together and use examples directly from the textbook (Activity 1.11). Examiners want the gain/loss language precisely.
• (b) This equation is directly from Exercise Q.14. Remember copper forms Cu(NO₃)₂ and silver is displaced as solid. Naming the reaction type is compulsory.
• (c) The identity of 'X' as copper comes from Activity 1.11. You must write both equations and explicitly state which substance is oxidised and which is reduced in Reaction 2 — that is the scoring point examiners specifically check.

(i)
$$\text{CaO}(s) + \text{H}_2\text{O}(l) \rightarrow \text{Ca(OH)}_2(aq) + \text{Heat}$$
It is an exothermic reaction, as a large amount of heat is released.

(ii)
$$\text{Ca(OH)}_2(aq) + \text{CO}_2(g) \rightarrow \text{CaCO}_3(s) + \text{H}_2\text{O}(l)$$
It is a combination reaction, as two substances combine to form a single product.

(iii) The chemical formula of marble is CaCO₃. It is the same as the calcium carbonate layer formed on walls during whitewashing, showing they are chemically identical.

(iv) The formation of slaked lime belongs to the combination reaction category, where calcium oxide and water (two reactants) combine to form a single product, calcium hydroxide.

Source: Chapter 1, Section 1.2.1 – Combination Reaction

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• (i) demands the balanced equation AND the type — don't forget "Heat" on the product side as shown in the textbook.
• (ii) The reaction is a combination reaction (two substances → one product); some students mistakenly call it a "neutralisation" — avoid that here.
• (iii) Simply state CaCO₃ and explain the link — the textbook explicitly says marble and the whitewash layer share the same formula.
• (iv) Even though it is also exothermic, the question asks for the "broad category of chemical reaction from types of reactions," which is combination reaction. Exothermic describes energy change, not the type of reaction.

(i) The brown fumes are nitrogen dioxide (NO₂).

$$2\text{Pb(NO}_3)_2\text{(s)} \xrightarrow{\Delta} 2\text{PbO(s)} + 4\text{NO}_2\text{(g)} + \text{O}_2\text{(g)}$$

(ii) Decomposition of silver chloride in sunlight:

$$2\text{AgCl(s)} \xrightarrow{\text{sunlight}} 2\text{Ag(s)} + \text{Cl}_2\text{(g)}$$

(iii) Light energy (solar/photochemical energy) causes the decomposition of silver chloride.
In contrast, lead nitrate is decomposed by heat energy (thermal energy). Thus, one reaction is photochemical and the other is thermal decomposition.

(iv) All decomposition reactions are endothermic — they require energy to be supplied (in the form of heat, light, or electricity) to break down a single substance into two or more simpler substances.

Source: Chapter 1, Chemical Reactions and Equations

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• (i): The yellow solid left behind is PbO (lead monoxide). Examiners expect the name "nitrogen dioxide" and a correctly balanced equation — check that atoms of Pb, N, and O balance on both sides.
• (ii): The greenish-yellow gas is Cl₂. The equation must show sunlight as the condition above the arrow, not as a reactant. Coefficient of 2 on both sides is essential for balance.
• (iii): The key distinction tested here is type of energy: light (photochemical) vs. heat (thermal). One line for each is enough.
• (iv): This is a direct concept from the chapter — decomposition reactions are endothermic because energy must be absorbed to break bonds. Contrast with combination reactions which are typically exothermic.

(i) The reddish-brown coating on the iron gate is called rust (iron oxide, Fe₂O₃·xH₂O). The chemical process responsible is corrosion, caused by the reaction of iron with moisture and oxygen from air.

(ii) The spoilage of fats and oils due to oxidation is called rancidity. The gas whose contact with oil causes this change is oxygen (O₂).

(iii) The iron gate can be protected by painting or galvanising (coating with zinc). The cooking oil can be protected by storing it in an airtight container or adding antioxidants (nitrogen flushing also slows oxidation).

(iv) Similarity: Both rusting and rancidity are caused by oxidation reactions.
Difference: Rusting affects metals (iron), while rancidity affects fats and oils (organic substances).

Source: Chapter 1, Section 1.3 (Corrosion and Rancidity); Chapter 3, Section 3.5

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• (i) Examiners expect both the name "rust" and the process "corrosion" for full credit.
• (ii) "Rancidity" is the exact textbook term; oxygen is the gas to name (the textbook mentions nitrogen flushing to prevent it, implying oxygen causes it).
• (iii) One method each is enough — painting/galvanising for iron; airtight container/antioxidants for oil.
• (iv) The similarity (both = oxidation) and difference (metal vs. fat/oil) is the key contrast examiners look for. Keep it concise — one line each.