Common criterion: A chemical reaction has taken place when there is a change in the chemical nature of the substance, resulting in the formation of one or more new substances with different properties. Observable indicators include change in colour, state, temperature, or evolution of gas.

Both rusting of iron and burning of magnesium produce new substances (iron oxide and magnesium oxide respectively), confirming both are chemical reactions.

Type of reaction:

• Rusting of iron — It is a corrosion (slow oxidation) reaction:

$4\text{Fe} + 3\text{O}_2 + x\text{H}_2\text{O} \rightarrow 2\text{Fe}_2\text{O}_3 \cdot x\text{H}_2\text{O}$

• Burning of magnesium — It is a rapid oxidation (combination) reaction:

$2\text{Mg} + \text{O}_2 \rightarrow 2\text{MgO}$

Both are oxidation reactions, but differ greatly in rate.

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

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• Key criterion (1 mark): The examiner wants you to state that formation of a new substance with changed properties = chemical reaction. The textbook lists indicators: change in colour, state, temperature, gas evolution.
• Rusting (1 mark): Identify it as corrosion/slow oxidation. Mentioning it occurs in humid atmosphere matches the textbook ("The surface of some metals, such as iron, is corroded when they are exposed to moist air for a long period of time").
• Burning Mg (1 mark): Identify it as rapid oxidation or combination reaction — textbook explicitly calls it burning in oxygen forming magnesium oxide.
• Both are oxidation/combination reactions, but the rate difference is what the question highlights — always note this contrast.
• Writing balanced equations earns full credit and demonstrates understanding; don't skip them in a 3-mark answer.

Evolution of bubbles (gas produced) and the container becoming warm (heat released/exothermic reaction) confirm that a chemical reaction has taken place.

CBSE expects students to name the two observable indicators: production of gas (bubbles of hydrogen) and change in temperature (heat released). Both are standard signs of a chemical reaction from Chapter 1. No equation is needed for 1 mark.

The student's argument is incorrect. Cooking food is a chemical change, not physical. During cooking, new substances are formed due to breaking and making of bonds between atoms. Indicators like change in colour, change in smell, and irreversibility confirm a chemical reaction has occurred. The change cannot be reversed, unlike physical changes.

Source: Chapter 1 – Chemical Reactions and Equations, Chapter Introduction

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• The textbook explicitly lists cooking food as an example where "the nature and identity of the initial substance have somewhat changed" — i.e., a chemical reaction occurs.
• Key examiner expectation: state that new substances are formed (bonds broken and made) and give evidence of chemical change (irreversibility, change in colour/smell).
• Don't just say "it is a chemical change" — briefly justify why with at least one indicator. That earns both marks.

(i) This approach is chemically incorrect because we cannot change the formula of any reactant or product to balance an equation. H₂O₄ is not a real compound; altering a formula violates the law of conservation of mass and misrepresents the actual substances involved. Balancing is done only by placing coefficients (whole numbers) in front of the formulae.

(ii) Using the hit-and-trial method (placing coefficients only):

$$3\text{Fe}(s) + 4\text{H}_2\text{O}(g) \rightarrow \text{Fe}_3\text{O}_4(s) + 4\text{H}_2(g)$$

Verification:

| Element | LHS | RHS |
|---------|-----|-----|
| Fe | 3 | 3 ✓ |
| H | 8 | 8 ✓ |
| O | 4 | 4 ✓ |

Source: Chapter 1, Section 1.1.2 Balanced Chemical Equations

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• Part (i) carries ~1 mark: state the rule clearly — formulae must never be changed; only coefficients are adjusted. The textbook explicitly says "we cannot alter the formulae of the compounds."
• Part (ii) carries ~2 marks: write the balanced equation with correct coefficients (3, 4, 1, 4). The verification table is good practice and shows the examiner you have checked your answer, but keep it brief. State symbols are a bonus if asked; include them if the question says "with state symbols."
• The key term to use: hit-and-trial method (as named in the NCERT text).

The equation Mg + O₂ → MgO is called a skeletal chemical equation because it only shows the reactants and products using chemical formulae, but the number of atoms of each element is not equal on both sides. On the LHS, there are 2 oxygen atoms (O₂), but only 1 on the RHS (MgO). This makes it unbalanced.

To make it scientifically valid, it must be balanced by adding appropriate coefficients:

$$2\text{Mg} + \text{O}_2 \rightarrow 2\text{MgO}$$

This step is necessary because of the Law of Conservation of Mass, which states that mass can neither be created nor destroyed in a chemical reaction. Therefore, the number of atoms of each element must be equal on both sides.

Source: Chapter 1, Sections 1.1.1 and 1.1.2

Examiners look for three things here (1 mark each): (1) defining "skeletal equation" and identifying the imbalance, (2) writing the correctly balanced equation, and (3) naming the Law of Conservation of Mass and stating it. Don't just say "atoms are not equal" — explain which atom is unequal as evidence. Always write the balanced equation explicitly; it earns a dedicated mark.

(i) The equation Mg + O₂ → MgO is called a skeletal chemical equation because it only represents the reactants and products using chemical formulae, without indicating the correct number of atoms/molecules of each substance.

(ii) It is unbalanced because the number of atoms of each element is not equal on both sides — there are 2 oxygen atoms on LHS but only 1 on RHS, violating the law of conservation of mass.

Balanced equation:

$$2\text{Mg} + \text{O}_2 \rightarrow 2\text{MgO}$$

Source: Chapter 1, Section 1.1.1 Writing a Chemical Equation

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• Examiners expect two distinct points: definition of skeletal equation + reason it is unbalanced. Both parts carry marks separately.
• The balanced equation must be shown explicitly — simply stating "add coefficient 2" is not enough.
• Key phrase to use: "law of conservation of mass" — it signals you understand why balancing is necessary.
• Do not change the formula of any compound (e.g., do not write MgO₂); only add coefficients.

Since mass is conserved (Law of Conservation of Mass), the number of atoms of each element must be equal on both sides, making it compulsory to balance a chemical equation.

The examiner expects you to link mass conservation directly to the need for balancing. Name the law clearly. One crisp sentence is sufficient for 1 mark.

Changing subscripts in a chemical formula changes the substance itself. For example, H₂O₄ is a completely different compound from H₂O. We can only add coefficients (e.g., 4 H₂O) to balance an equation, because the law of conservation of mass requires that the number of atoms of each element remains the same — not that we alter the identities of the reactants or products.

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

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Examiners look for two key points (1 mark each):

1. Changing subscripts changes the identity/formula of the compound — it becomes a different substance entirely.
2. Only coefficients may be changed; this follows from the law of conservation of mass, which governs why we balance equations in the first place.

The textbook explicitly states: "we cannot alter the formulae of the compounds or elements involved in the reactions… we can put coefficient '4' as 4 H₂O and not H₂O₄." Quote or paraphrase this directly if possible.

The notation (aq) means the substance is dissolved in water (aqueous solution). For example, NaCl(aq) indicates sodium chloride dissolved in water.

(l) denotes that the substance is in pure liquid form, not dissolved in any solvent. For example, H₂O(l) means liquid water itself.

Difference: (aq) = dissolved in water; (l) = pure liquid state.

Source: Chapter 1, Writing a Chemical Equation (physical states notation); Chapter 2, Acids-Bases

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• Examiners expect both notations to be defined clearly and a direct comparison made — that earns both marks.
• A common mistake is saying (l) means "liquid solution" — it strictly means the pure substance in liquid phase, not a solution.
• You can support with examples (e.g., HCl(aq) vs Br₂(l)) to make the distinction crisp.
• The source passages mention H⁺(aq) and OH⁻(aq) ions in solution, confirming (aq) = dissolved in water.

The symbol (g) is used with H₂O because in this reaction, water is not in liquid form but is used as steam (water vapour). Since the reaction requires high temperature, water exists in the gaseous state. The notation (g) indicates this physical state accurately, making the equation more informative.

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

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• The textbook explicitly states: "the symbol (g) is used with H₂O to indicate that in this reaction water is used in the form of steam."
• Examiners expect two key points: (1) water is in the form of steam/vapour due to high temperature, and (2) (g) denotes gaseous state, not liquid.
• Do not just say "water is heated" — say it is used as steam, which is the gaseous state.

The arrow in a chemical equation represents the direction of the reaction, showing that reactants (LHS) are converted into products (RHS).

Reaction conditions like heat or a catalyst are written above or below the arrow because they are not reactants or products — they only influence how the reaction proceeds. Placing them on the arrow keeps the equation accurate and uncluttered.

Source: Chapter 1, Section 1.1.2 Writing a Chemical Equation

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• 1 mark for stating what the arrow represents (direction of reaction / conversion of reactants to products).
• 1 mark for explaining why conditions go on the arrow — they are neither reactants nor products, only facilitating factors.
• Use examples from the textbook (e.g., sunlight/chlorophyll written above/below the arrow in Eq. 1.12) to support if needed, but keep it brief for 2 marks.
• Avoid writing a long paragraph; examiners look for these two specific points clearly stated.

(B) 2Al + 3Cl₂ → 2AlCl₃

LHS: 2 Al, 6 Cl atoms. RHS: 2 Al, 6 Cl atoms — both sides are equal, so the equation is balanced.

To balance Al + Cl₂ → AlCl₃, apply the hit-and-trial method. AlCl₃ needs 3 Cl atoms, but Cl₂ provides 2 — so use 3 Cl₂ (6 Cl atoms) and 2 AlCl₃ (6 Cl atoms), then balance Al with coefficient 2. Options C and D use incorrect formulae/products, violating the rule that formulae cannot be altered during balancing.

When water (H₂O) can exist as liquid or steam, omitting its physical state makes the equation misleading. For example, in the reaction of iron with water:

$$3\text{Fe}(s) + 4\text{H}_2\text{O}(g) \rightarrow \text{Fe}_3\text{O}_4(s) + 4\text{H}_2(g)$$

Here, the symbol (g) is essential to indicate that water is used as steam, not liquid water. Without it, the reader cannot determine the actual reaction conditions, making the equation incomplete and potentially misleading about how the reaction proceeds.

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

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• Examiners want: (1) a specific, relevant example, (2) the physical state symbol shown, and (3) a clear reason why omission is misleading.
• The textbook explicitly uses the Fe/H₂O example to illustrate exactly this point — always use it here rather than inventing another.
• Key phrase to include: "water is used in the form of steam" — this is the textbook's own wording and signals to examiners you know the source.
• Don't just say "physical states are important in general"; the question asks for ONE specific situation with an explanation.

Word Equation vs. Balanced Chemical Equation

A word equation names the reactants and products in words, showing the direction of the reaction. It only tells us what substances are involved.

Example:
Zinc + Sulphuric acid → Zinc sulphate + Hydrogen

A balanced chemical equation uses chemical formulae and coefficients, ensuring the number of atoms of each element is equal on both sides (obeying the Law of Conservation of Mass).

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

Additional information provided by a balanced chemical equation:

1. Chemical formulae of all reactants and products.
2. Relative number of atoms/molecules taking part in the reaction.
3. Physical states — (s), (l), (g), (aq) — can be added.
4. Reaction conditions (temperature, pressure, catalyst) can be shown above/below the arrow.
5. Confirms the Law of Conservation of Mass — atoms are neither created nor destroyed.

Source: Chapter 1, Sections 1.1 and 1.1.2

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• Examiners expect a clear comparison (1–2 marks), a named example with both equations (1–2 marks), and at least 3 additional points a balanced equation provides (remaining marks).
• The Zn + H₂SO₄ example is directly from the textbook — always prefer textbook examples.
• Don't forget to mention physical states and reaction conditions as extra information; students often miss these.
• The Law of Conservation of Mass is the reason equations must be balanced — mentioning it earns a mark.

(i) Observations and Equation:

1. The iron nail acquires a brownish (reddish-brown) coating of copper on its surface.
2. The blue colour of copper sulphate solution fades (becomes pale green due to formation of FeSO₄).

Balanced equation:

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

(ii) Type of Reaction and Property Determining Displacement:
This is a displacement reaction — iron displaces copper from copper sulphate solution. The property that determines whether one metal can displace another is its reactivity (position in the activity series). A more reactive metal displaces a less reactive metal from its salt solution. Since iron is more reactive than copper (iron is above copper in the reactivity series), it displaces copper.

(iii) Oxidised and Reduced Substances:

• Iron (Fe) is oxidised — it loses electrons, changing from Fe (neutral metal) to Fe²⁺ in FeSO₄.
• Copper (Cu²⁺) is reduced — it gains electrons, changing from Cu²⁺ in CuSO₄ to Cu (neutral metal).

Source: Chapter 1, Section 1.2.3 (Displacement Reaction); Chapter 3, Section 3.2.4 & 3.2.5

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• Part (i): Examiners expect both observations clearly — the colour change of the nail AND the fading of the blue colour. The equation must be balanced with state symbols.
• Part (ii): The key term is "reactivity series" / "activity series." Just saying "iron is more reactive" without linking it to the series/concept of reactivity will lose a mark.
• Part (iii): Always use the electron language — "loses electrons = oxidised," "gains electrons = reduced." Link it back to the specific species (Fe → Fe²⁺ and Cu²⁺ → Cu). This is the OIL RIG concept applied to the reaction.

This is a combination reaction because two reactants (CaO and H₂O) combine to form a single product (calcium hydroxide).

$$\text{CaO}(s) + \text{H}_2\text{O}(l) \rightarrow \text{Ca(OH)}_2(aq) + \text{Heat}$$

Experimental observation: When water is added to calcium oxide, the beaker becomes hot (rise in temperature is felt on touching the beaker), showing that a large amount of heat is released. This makes it an exothermic reaction.

Source: Chapter 1, Section 1.2.1 Combination Reaction

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• 1 mark for correctly identifying it as a combination reaction (and optionally exothermic).
• 1 mark for the observation — touching the beaker and feeling the rise in temperature is the key experimental evidence examiners expect. Saying "heat is released" without linking it to the observation may cost you the mark.
• Always write the balanced equation when the reaction is mentioned — it strengthens your answer and takes only one line.

The reaction responsible for the hard, shiny finish is:

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

Calcium hydroxide (slaked lime) reacts with carbon dioxide present in air to form calcium carbonate (CaCO₃), which gives the wall a hard, shiny finish. This is a combination reaction.

The finish takes two to three days to appear because calcium hydroxide reacts slowly with CO₂ in air. The calcium carbonate layer builds up gradually over this period before it becomes visible.

Source: Chapter 1, Section 1.2.1 – Combination Reaction

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• Examiners expect the balanced chemical equation — award it 1 mark.
• The reaction type (combination reaction) should be named — 1 mark.
• The reason for the delay ("reacts slowly with CO₂") must be stated explicitly — 1 mark.
• Do not confuse this with the preparation of slaked lime (CaO + H₂O); the whitewashing reaction is Ca(OH)₂ with CO₂.
• Note: CaCO₃ is also the chemical formula of marble — a nice fact but not required for marks here.

Category: Exothermic reactions.

Defining feature: In both reactions, heat (energy) is released to the surroundings during the reaction.

• Burning of coal: $\text{C} + \text{O}_2 \rightarrow \text{CO}_2 + \text{heat and light}$
• Formation of water: $\text{H}_2 + \text{O}_2 \rightarrow \text{H}_2\text{O} + \text{heat}$

Both release energy, making them exothermic.

Source: Chapter 1, Types of Chemical Reactions (What you have learnt)

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• Examiners expect the category name (1 mark) and the defining feature — energy/heat released (1 mark).
• Writing a supporting equation is good practice but not compulsory for 2 marks; it strengthens the answer.
• Do not confuse with combustion (combustion is a type of exothermic reaction, not the broader category asked here). The question asks for the common classification, which is exothermic reactions.
• Key phrase to remember: "Reactions in which heat is given out along with the products are called exothermic reactions." (Chapter 1 summary)

The student is not correct. Not every combination reaction is exothermic.

A combination reaction simply involves two or more substances combining to form a single product. While many combination reactions do release heat — for example:

$$\text{CaO}(s) + \text{H}_2\text{O}(l) \rightarrow \text{Ca(OH)}_2(aq) + \text{Heat}$$

— this is not a defining feature of all combination reactions.

The formation of water ($2\text{H}_2 + \text{O}_2 \rightarrow 2\text{H}_2\text{O}$) and burning of coal ($\text{C} + \text{O}_2 \rightarrow \text{CO}_2$) are also exothermic combination reactions, but there exist combination reactions that require energy input (endothermic). Exothermic nature is a property of some combination reactions, not all.

Source: Chapter 1, Section 1.2.1 Combination Reaction

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• CBSE expects you to clearly state the student is wrong, then justify with a definition of combination reaction and at least one example.
• The key distinction: combination reaction is defined by structure (many → one), not by energy change. Exothermic reactions are separately defined as those releasing heat.
• Mention both the definition and at least one example reaction with equation for full 3-mark credit.
• Avoid writing a long essay — the answer above is appropriately concise for 3 marks.

Product: When magnesium ribbon burns in air, the product formed is magnesium oxide (MgO).

$$2\text{Mg} + \text{O}_2 \rightarrow 2\text{MgO}$$

Two types of reactions:

1. Combination reaction – Two reactants (magnesium and oxygen) combine to form a single product (magnesium oxide).

1. Oxidation reaction (Redox) – Magnesium gains oxygen during the reaction, so it is oxidised. Hence, this is also an oxidation reaction.

Thus, the same reaction is classified under both combination and oxidation reactions.

Source: Chapter 1, Section 1.1 & 1.2; Chapter 3, Section 3.2.1

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• Examiners expect the student to name MgO as the product and write the balanced equation (1 mark).
• Two reaction types must be named and justified, not just listed — each needs a one-line reason (1 mark each).
• Common accepted pairs: Combination + Oxidation. Some answers also accept Exothermic as the second type if the question allows it, but Combination + Oxidation is the most textbook-aligned response here.
• Avoid writing lengthy explanations; crisp labelled points score better in board exams.

When ferrous sulphate (FeSO₄) crystals are heated strongly, they decompose to form three products:

1. Ferric oxide (Fe₂O₃) — a reddish-brown solid
2. Sulphur dioxide (SO₂) — a gas
3. Sulphur trioxide (SO₃) — a gas

$$2\text{FeSO}_4 \xrightarrow{\Delta} \text{Fe}_2\text{O}_3 + \text{SO}_2 + \text{SO}_3$$

Evidence of chemical change (from observation alone):

• The colour changes from green (FeSO₄ crystals) to reddish-brown (Fe₂O₃).
• A characteristic odour of burning sulphur is produced (due to SO₂/SO₃ gases).

These observations — colour change and gas evolution with odour — confirm a chemical change, not merely a physical one.

Source: Chapter 1, Section 1.2.2 — Decomposition Reaction

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• Examiners expect you to name all three products with correct formulae — each carries marks.
• The equation (optional but good) shows it's a decomposition/thermal decomposition reaction.
• For "observation alone," stick to what you see (colour change: green → reddish-brown) and smell (sulphur odour) — do not mention formulae or theory here.
• A physical change is reversible and produces no new substance; colour change + new gas with odour proves new substances have formed → chemical change.

(B) Heat energy; thermal decomposition

CaCO₃ is heated strongly to decompose into CaO and CO₂. This is thermal decomposition — a single substance breaks down using heat energy.

• The reaction requires high temperature (heat energy), ruling out options A (electrical) and C (light).
• It is a decomposition reaction (one reactant → two products), not a combination or double displacement, ruling out option D.
• Key term to remember: decomposition using heat = thermal decomposition. Classic CBSE example: CaCO₃ → CaO + CO₂.

Combination Reaction — two or more reactants combine to form a single product:

$$2\text{Mg}(s) + \text{O}_2(g) \rightarrow 2\text{MgO}(s)$$

Decomposition Reaction — a single reactant breaks down into two or more products:

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

Fundamental Difference: In a combination reaction, multiple reactants converge to form one product (many → one). In a decomposition reaction, one substance splits into multiple products (one → many). This shows that the direction of chemical change is opposite in the two reactions — combination builds complexity while decomposition breaks it down.

Source: Chapter 1 — Types of Chemical Reactions, What You Have Learnt

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• The examiner expects one equation each, both balanced — award 1 mark each.
• The third mark is for clearly stating the directional difference (many→one vs one→many).
• Avoid writing long descriptions; use the equations to do the work. State symbols earn you extra credit but are not strictly required for 3 marks here.
• Key phrase from the textbook: "Decomposition reactions are opposite to combination reactions" — use this language.

Water (H₂O) is composed of hydrogen and oxygen in the ratio 2:1 by volume. During electrolysis, water molecules break down as:

$$2H_2O \rightarrow 2H_2 + O_2$$

Since two molecules of hydrogen are produced for every one molecule of oxygen, the volume of hydrogen collected is exactly double that of oxygen.

• Gas at the cathode (negative electrode): Hydrogen (H₂) — larger volume
• Gas at the anode (positive electrode): Oxygen (O₂) — smaller volume

The cathode produces the larger volume of gas.

• The key reason is the 2:1 molar (and volume) ratio of H to O in water — examiners expect you to state this clearly.
• Always identify which electrode produces which gas: cathode → hydrogen, anode → oxygen. This is a common 1-mark direct question.
• You may write the decomposition equation for full marks; it shows the 2:1 ratio directly.
• The source passages provided do not cover this topic directly, so this answer draws on standard CBSE Class 10 Science (Chapter 1 – Chemical Reactions and Equations / Chapter 2 knowledge of water electrolysis as taught in the course).

Silver chloride (AgCl) undergoes photodecomposition when exposed to sunlight (light energy). The reaction is:

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

AgCl decomposes into silver metal and chlorine gas, turning grey. Dark-coloured bottles prevent light from reaching the compound, thus stopping this decomposition.

This is a photochemical decomposition reaction — a type of decomposition triggered by light, not heat. Examiners expect: (1) naming the type of reaction, (2) the balanced equation, and (3) the practical reason for dark bottles. The textbook uses AgCl as a standard example of photodecomposition, so memorise the equation. Note: the source passages mention corrosion of silver (with sulphur), but the photodecomposition of AgCl is the standard NCERT Class 10 Chemistry example from Chapter 1 (decomposition reactions section).

When lead nitrate is heated, it undergoes thermal decomposition. The balanced chemical equation is:

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

Products identified:

1. Lead monoxide (PbO) — yellow solid residue formed in the test tube.
2. Nitrogen dioxide (NO₂) — brown/reddish-brown fumes observed.
3. Oxygen (O₂) — colourless, odourless gas released.

This is a thermal decomposition reaction (endothermic).

Source: Chapter 1, Chemical Reactions and Equations

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• The examiner expects the correctly balanced equation (1 mark), identification of all three products (1 mark each, or grouped as 1 mark), and the type of reaction (often implied or explicitly asked).
• The most common mistake is missing O₂ as a product — students notice brown fumes (NO₂) but forget oxygen is also released.
• Remember: 2 formula units of Pb(NO₃)₂ give 2 PbO + 4 NO₂ + 1 O₂ — verify atom count on both sides.
• Brown fumes → NO₂; colourless gas → O₂; yellow residue → PbO.

(i) Decomposition of calcium carbonate (thermal decomposition):

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

(ii) Calcium oxide (CaO) / Quick lime is directly used in the construction industry. It reacts vigorously with water to form slaked lime [Ca(OH)₂], which is used for whitewashing walls and in making cement and mortar.

(iii) This reaction is a chemical change because new substances with different properties are formed (CaO and CO₂ from CaCO₃), bonds are broken and new bonds are formed, and the reaction is irreversible under normal conditions. The composition of matter changes permanently, which is the hallmark of a chemical change.

Source: Chapter 1, Section 1.2 (Decomposition Reactions); Section 1.2.1 (Combination Reaction – slaked lime use)

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• (i) Always include the heat symbol (Δ) above the arrow and state symbols — examiners award a mark for a correctly balanced equation with conditions.
• (ii) The textbook specifically mentions CaO (quick lime) reacting with water → slaked lime → whitewashing. Stick to this use.
• (iii) Key indicators of chemical change: new substances formed, different properties, irreversibility, bond breaking/making. Mentioning "no new elements" is not enough — you must explain why it is still chemical (new compounds with new properties are produced).

Iron is more reactive than copper, so it displaces copper from copper sulphate solution. The reaction is:

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

At the ionic level, iron atoms lose electrons and enter the solution as Fe²⁺ ions, forming iron sulphate (FeSO₄). Simultaneously, Cu²⁺ ions from the solution gain electrons and get deposited as copper metal on the nail. This deposited copper gives the nail a brownish coating. As Cu²⁺ ions are consumed and replaced by Fe²⁺ ions, the concentration of copper sulphate decreases, causing the blue colour to fade.

Source: Chapter 1, Section 1.2.3 — Displacement Reaction

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• The examiner expects you to name this as a displacement reaction and write the equation (1 mark).
• You must explain both observations separately: the brownish coating (Cu deposited) and the fading blue colour (Cu²⁺ ions replaced by Fe²⁺) — likely 1 mark each.
• Mentioning electron transfer (Fe loses electrons → Fe²⁺; Cu²⁺ gains electrons → Cu) earns the "atomic/ionic level" mark.
• Do not confuse this with rusting; this is a chemical displacement, not corrosion.

These observations show that zinc and lead are more reactive than copper. A more reactive metal displaces a less reactive metal from its salt solution. Since copper cannot displace zinc or lead from their salt solutions, copper is less reactive than both zinc and lead.

Broader principle: A more reactive metal displaces a less reactive metal from its salt solution. This is the principle of displacement reactions, based on the activity series.

Source: Chapter 1, Section 1.2.3 Displacement Reaction

• The examiner expects you to state the reactivity order (Zn > Cu, Pb > Cu) and explain why no reaction occurs when copper is placed in zinc sulphate or lead nitrate (copper is less reactive, so it cannot displace them).
• Naming the broader principle — "a more reactive metal displaces a less reactive metal from its salt solution" — is essential for full marks.
• Mentioning the activity series by name adds precision and is rewarded.

(i) Types of reaction:

• Displacement reaction: Fe displaces Cu from CuSO₄ because iron is more reactive than copper, so a more reactive metal displaces a less reactive metal from its salt solution.
• Redox reaction: Both oxidation and reduction occur simultaneously in this reaction.

(ii) Oxidation and Reduction:

• Fe is oxidised: It loses electrons/gains no oxygen, going from Fe (0) to Fe²⁺ in FeSO₄. (Iron gains no oxygen here but is oxidised as a species.)
• Cu²⁺ is reduced: Copper ions in CuSO₄ gain electrons to form Cu(s), i.e., Cu²⁺ → Cu(0) — loss of combined state/gain of electrons.

This is determined by tracking changes in the state of the elements: Fe becomes part of a compound (oxidised) and Cu²⁺ becomes free copper (reduced).

Source: Chapter 1, Section 1.2.5 Oxidation and Reduction; What you have learnt, Chapter 1 & 3

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• Examiners expect two reaction types named with reasons: displacement (reactivity basis) and redox (simultaneous oxidation + reduction).
• For Part (ii), at Class 10 level, "oxidised = loses electrons / forms compound" and "reduced = gains electrons / becomes free element" is sufficient. You may also use gain/loss of oxygen language if preferred, but electron transfer is cleaner here since no oxygen transfer occurs directly.
• Always state how you determined it — that's explicitly asked and carries marks.

When Na₂SO₄ and BaCl₂ solutions are mixed, the ions exchange partners. The reaction is:

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

BaSO₄ forms a precipitate because it is insoluble in water. The SO₄²⁻ and Ba²⁺ ions combine to form this white insoluble solid.

NaCl remains dissolved because it is soluble in water; Na⁺ and Cl⁻ ions stay in solution.

This is a double displacement reaction (also called a precipitation reaction) because the ions of the two reactants exchange places to form two new compounds.

Source: Chapter 1, Section 1.2.4 – Double Displacement Reaction

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• Examiners expect the balanced equation with state symbols — always include (s) for precipitate and (aq) for dissolved species.
• You must name both types: double displacement reaction AND precipitation reaction for full marks.
• The key reason for BaSO₄ precipitating is its insolubility in water — state this explicitly.
• NaCl staying in solution must be explained by its solubility, not just stated.
• Three marks typically split as: equation (1) + reason for precipitate/solution (1) + type of reaction (1).

Yes, I agree with the statement.

Justification:
In a double displacement reaction, ions are exchanged between two compounds. A precipitation reaction is a type of double displacement reaction where one product is insoluble (precipitate). However, double displacement reactions can also produce water (neutralisation) or a gas instead of a precipitate.

Example of a double displacement reaction WITHOUT a precipitate:

$$\text{NaOH}(aq) + \text{HCl}(aq) \rightarrow \text{NaCl}(aq) + \text{H}_2\text{O}(l)$$

Here, OH⁻ and Cl⁻ ions are exchanged (double displacement), but no precipitate is formed — only water and sodium chloride in solution are produced.

Source: Chapter 1, Section 1.2.4 Double Displacement Reaction

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• The examiner awards marks for: (1) agreeing and stating the logic (all precipitation reactions involve ion exchange = double displacement), (2) explaining that double displacement reactions are broader (can produce water/gas too), and (3) a correct example with equation.
• The neutralisation reaction NaOH + HCl is the simplest and most reliable example to use here since it appears in the chapter's exercises.
• Do not write a neutralisation reaction as a precipitation reaction — they are distinct subtypes of double displacement.

The reaction is:

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

Hydrogen is oxidised because it gains oxygen (from CuO) to form water. Oxidation is defined as the gain of oxygen by a substance.

Copper(II) oxide is reduced because it loses oxygen to form copper. Reduction is defined as the loss of oxygen by a substance.

Since one reactant is oxidised and the other is reduced simultaneously, this is called a redox reaction.

Source: Chapter 1, Section 1.2.5 — Oxidation and Reduction

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• 3 marks are split as: 1 mark for identifying the oxidised substance with reason, 1 mark for identifying the reduced substance with reason, 1 mark for the definitions (or the term "redox reaction").
• Always state the definition of oxidation/reduction explicitly — examiners look for "gains oxygen" / "loses oxygen" in the answer.
• Writing the chemical equation at the start is good practice and shows clarity.
• Do not confuse: hydrogen gains oxygen → oxidised; CuO loses oxygen → reduced. A common mistake is flipping these.

In ZnO + C → Zn + CO:

• ZnO is the oxidising agent — it gives oxygen to carbon, thereby getting reduced (loses oxygen to form Zn).
• C is the reducing agent — it gains oxygen from ZnO to form CO, thereby getting oxidised.

Oxidation = gain of oxygen; Reduction = loss of oxygen.

Source: Chapter 1, "What you have learnt"

• Examiners award 1 mark for correctly identifying both agents and 1 mark for the explanation in terms of gain/loss of oxygen.
• A common mistake is confusing the two: always link "oxidising agent" to the substance that provides oxygen (and itself gets reduced), and "reducing agent" to the substance that takes oxygen (and itself gets oxidised).
• Stating the definitions of oxidation and reduction in the last line shows conceptual clarity and helps secure full marks.

The student's claim is incorrect.

Oxidation and reduction always occur simultaneously in a chemical reaction. If a substance gains oxygen (or loses hydrogen), it is oxidised; if a substance loses oxygen (or gains hydrogen), it is reduced. Such reactions are called redox reactions.

For example: $\text{CuO} + \text{H}_2 \xrightarrow{\text{Heat}} \text{Cu} + \text{H}_2\text{O}$

Here, CuO is reduced (loses oxygen) and H₂ is oxidised (gains oxygen) at the same time. One cannot occur without the other — the oxygen lost by one substance must be gained by another.

Source: Chapter 1, Section 1.2.5 — Oxidation and Reduction

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• The examiner expects students to clearly state the claim is wrong and explain why using the definition of oxidation and reduction from the textbook.
• The key concept: oxidation and reduction are complementary — oxygen lost by one reactant is gained by another. This is why such reactions are called redox reactions.
• Including the example reaction (CuO + H₂) with identification of which substance is oxidised and which is reduced scores full marks. Don't just define — apply the concept.

(i) Iron: Corrosion (rusting) — product is rust, $\text{Fe}_2\text{O}_3.\text{xH}_2\text{O}$. Silver: Corrosion (tarnishing) — product is silver sulphide, $\text{Ag}_2\text{S}$.
(ii) Similar: Both are corrosion — metals reacting with substances in air. Different: Iron reacts with oxygen and moisture; silver reacts with sulphur in air.

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

Even though this is marked as 1 mark, it has two parts — examiners expect both to be addressed briefly. Key points: use the word "corrosion" for both, give correct formulae ($\text{Fe}_2\text{O}_3.\text{xH}_2\text{O}$ and $\text{Ag}_2\text{S}$), and note that the difference lies in what the metal reacts with (oxygen + moisture vs. sulphur).

Chips manufacturers flush packets with nitrogen gas because nitrogen is an inert gas that does not react with food. Chips are oily/fatty snacks that undergo oxidation (rancidity) in the presence of oxygen, making them stale, foul-smelling, and unfit for eating. Nitrogen displaces oxygen inside the packet, preventing this oxidation reaction.

If packets were sealed with air instead, the oxygen present in air would oxidise the fats/oils in the chips, causing rancidity. The chips would become stale quickly, develop a bad smell and taste, and lose their quality before the expiry date.

Source: Chapter 1, Section 1.2.5 – Oxidation and Reduction

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• The key concept tested here is rancidity, which is oxidation of fats/oils by atmospheric oxygen.
• Nitrogen is chosen because it is inert — it does not react with food.
• The examiner expects two parts: (1) why nitrogen is used (prevents oxidation/rancidity), and (2) what happens with air (oxidation occurs → rancidity). Cover both for full marks.
• The word "rancidity" (or "oxidation of fats") is essential — do not omit it.

Oxidation and its harmful effects:

• Rusting of iron: When iron is exposed to moist air, it reacts with oxygen and water to form reddish-brown rust (iron oxide). This damages the metal structurally, causing car bodies, bridges, and iron railings to deteriorate, leading to enormous financial losses every year.
• Rancidity of food: When fats and oils in food are oxidised, they become rancid — their smell and taste change, making food unfit for consumption.

How preventive methods tackle the root cause:

• Painting iron: Rust requires both air (oxygen) and moisture. Painting creates a physical barrier that cuts off contact between iron and moist air, directly eliminating the conditions needed for oxidation.
• Adding antioxidants: These substances preferentially react with oxygen themselves, preventing the fats/oils from getting oxidised and turning rancid.
• Airtight containers / nitrogen flushing: These methods remove or displace oxygen around the food (chips manufacturers flush bags with nitrogen), so oxidation of fats simply cannot occur.

Both strategies share the same logic: deny oxygen access to the substance being protected.

Source: Chapter 1, Section 1.3; Chapter 3, Section 3.5

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• Examiners want you to clearly connect why oxidation is harmful in each case (structural damage vs. taste/smell change) and then analyse — not just list — how each prevention method works against the root cause (oxygen + moisture for rusting; oxygen for rancidity).
• The word "analyse" in the question means you must explain the mechanism of protection, not just name the method.
• Mentioning nitrogen flushing shows you've read the passage carefully and earns extra credit.
• Avoid padding; keep points crisp and targeted at marks.

Burning of Magnesium (Combination Reaction):

$$2\text{Mg}(s) + \text{O}_2(g) \rightarrow 2\text{MgO}(s)$$

Thermal Decomposition of Calcium Carbonate (Decomposition Reaction):

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

Reasoning:

In the first reaction, two reactants (Mg and O₂) combine to form a single product (MgO) — this defines a combination reaction.

In the second reaction, a single reactant (CaCO₃) breaks down into two products (CaO and CO₂) on heating — this defines a decomposition reaction. The two reactions are therefore opposites of each other.

Source: Chapter 1, Section 1.2.1 (Combination Reaction); Chapter 1, Section 1.2.2 (Decomposition Reaction)

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• The examiner awards marks for: (1) correct balanced equation for Mg burning, (2) correct balanced equation for thermal decomposition of CaCO₃, and (3) clear reasoning linking "two reactants → one product" = combination, and "one reactant → two products" = decomposition.
• The key phrase to remember: decomposition is the opposite of combination — the textbook states this explicitly.
• Always balance equations: Mg needs a coefficient of 2, giving 2MgO.
• The heat/delta symbol (Δ) above the arrow for CaCO₃ decomposition shows it requires heat (endothermic); include it for full credit.

When iron nails are placed in copper sulphate solution, iron displaces copper because iron is more reactive than copper.

(i) Colour change: Cu²⁺ ions (which give the blue colour) are removed from the solution and replaced by Fe²⁺ ions (forming light green FeSO₄). Hence, the blue colour fades.

(ii) Brown deposit: Copper metal is deposited on the iron nails, forming the brown coating.

Balanced equation:

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

Redox identification:

• Iron (Fe) is oxidised — it loses electrons (Fe → Fe²⁺).
• Copper (Cu²⁺) is reduced — it gains electrons (Cu²⁺ → Cu).

Source: Chapter 1, Section 1.2.3 Displacement Reaction

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Examiners award marks for: (1) correct balanced equation, (2) explanation of colour change linking Cu²⁺ removal, and (3) correct identification of oxidised and reduced substances. Always state why something is oxidised/reduced (loss/gain of electrons). The term "redox" implies both oxidation and reduction must be addressed — do not skip either.

(B) Decomposition of water by electricity — endothermic; combination of CaO with water — exothermic.

Electrolysis of water requires electrical energy to be absorbed, making it endothermic. The reaction of CaO with water releases a large amount of heat, making it exothermic. The textbook explicitly states CaO + H₂O → Ca(OH)₂ + Heat.

(a) Classification:

• Reaction (i) — Zn with dilute H₂SO₄ — is exothermic. Heat is released during the reaction, which is why the test tube warms up. Reactions that release heat along with products are called exothermic reactions.

• Reaction (ii) — Ba(OH)₂ with NH₄Cl — is endothermic. Energy is absorbed from the surroundings, causing the test tube to cool down. Reactions in which energy is absorbed are called endothermic reactions.

(b) Type of reaction (i) and balanced equation:

Reaction (i) is a displacement reaction — zinc (a more reactive metal) displaces hydrogen from dilute sulphuric acid.

Balanced equation:

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

(c) Everyday exothermic process:

Respiration is an exothermic process. Glucose is broken down in the body to release energy, carbon dioxide and water.

$$\text{C}_6\text{H}_{12}\text{O}_6 + \text{O}_2 \rightarrow \text{CO}_2 + \text{H}_2\text{O} + \text{energy}$$

Source: Chapter 1 — Chemical Reactions and Equations

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• Part (a) tests the core definitions of exothermic (heat released → surroundings warm) and endothermic (heat absorbed → surroundings cool). Always link the observation (warming/cooling) to the definition.
• Part (b) the question says "other than exothermic," so write displacement (not single-displacement or double-displacement). The balanced equation is directly given in the textbook as Eq. 1.3 — memorise it with state symbols.
• Part (c) any valid exothermic example scores marks — burning of fuels, combustion of natural gas, or respiration are all safe choices. Respiration is explicitly discussed in the NCERT text as an exothermic reaction and is the most expected answer.

Reaction 1: Decomposition of Silver Chloride

(i) Balanced Chemical Equation:

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

(ii) Identification of Oxidation and Reduction:

• Silver chloride (AgCl) is reduced — Ag⁺ in AgCl loses oxygen/chlorine and silver metal is formed; alternatively, AgCl loses a more electronegative partner. In terms of the textbook definition: Ag gains no oxygen, so it undergoes reduction (loss of the combined state to free metal, i.e., loss of oxygen/non-metal).
• Chloride is oxidised — chlorine is released as Cl₂ (loss of hydrogen / gain of a partner that reduces it is reversed); Cl⁻ loses electrons to form Cl₂.

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Reaction 2: Displacement of Copper by Zinc

(i) Balanced Chemical Equation:

$$\text{Zn}(s) + \text{CuSO}_4(aq) \rightarrow \text{ZnSO}_4(aq) + \text{Cu}(s)$$

(ii) Identification of Oxidation and Reduction:

• Zinc (Zn) is oxidised — Zinc loses electrons and combines with the sulphate, i.e., it goes from free element to a compound (gains oxygen/sulphate group). It is oxidised.
• Copper (Cu) is reduced — Copper in CuSO₄ is in combined form (with oxygen-containing sulphate); it is displaced as free copper metal, meaning it loses oxygen/combined state. Hence, CuSO₄ is reduced.

Source: Chapter 1, Section on Oxidation and Reduction (Gain or Loss of Oxygen)

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• The textbook defines oxidation as gain of oxygen or loss of hydrogen, and reduction as loss of oxygen or gain of hydrogen — always justify using these exact terms.
• For the silver chloride reaction, sunlight supplies energy (photochemical decomposition); remember to show the sunlight condition above the arrow.
• For the zinc-copper reaction, zinc displaces copper because zinc is more reactive — this is a classic displacement + redox example.
• Examiners award marks for: (a) correct balanced equation, (b) correct identification of oxidised substance, and (c) correct identification of reduced substance with brief justification. Missing any one loses a mark.

Precipitation Reaction involves two aqueous solutions reacting through an exchange of ions (double displacement). The insoluble solid forms because the exchanged ions combine to create an insoluble salt.

Example: Na₂SO₄(aq) + BaCl₂(aq) → BaSO₄(s) + 2NaCl(aq)
Here, Ba²⁺ and SO₄²⁻ ions exchange and form insoluble barium sulphate.

Decomposition Reaction involves a single compound breaking down into two or more simpler substances upon supply of energy (heat, light, or electricity).

Example: CaCO₃(s) → CaO(s) + CO₂(g)
The insoluble solid (CaO) is produced from the breakdown of one reactant.

Key distinction: Precipitation requires two reactants and ion exchange; decomposition requires only one reactant and an energy input.

Source: Chapter 1, Section 1.2 / 1.2.4

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• The question specifically says not to distinguish by counting reactants/products — so you must mention ion exchange (for precipitation) and energy input / single reactant breaking down (for decomposition).
• Examiners look for: correct examples with equations, and a clear conceptual distinction about the nature of reactants and the process.
• Precipitation is a sub-type of double displacement; always mention ion exchange.
• For decomposition, mentioning energy supply (heat/light/electricity) strengthens the answer.

(a) Balanced chemical equations:

Step 1 (Heating calcium carbonate):
$$\text{CaCO}_3(s) \xrightarrow{\text{Heat}} \text{CaO}(s) + \text{CO}_2(g)$$

Step 2 (CaO mixed with water):
$$\text{CaO}(s) + \text{H}_2\text{O}(l) \rightarrow \text{Ca(OH)}_2(aq) + \text{Heat}$$

Step 3 (Lime water exposed to air):
$$\text{Ca(OH)}_2(aq) + \text{CO}_2(g) \rightarrow \text{CaCO}_3(s) + \text{H}_2\text{O}(l)$$

(b) Types of reaction:

• Step 1: Decomposition reaction (a single compound breaks down into two simpler substances on heating).
• Step 3: Combination reaction (calcium hydroxide and carbon dioxide combine to form a single product, calcium carbonate).

(c) The second step is an exothermic reaction. Calcium oxide reacts vigorously with water to form calcium hydroxide, releasing a large amount of heat — this is why the container becomes hot. It is a combination reaction, and combination reactions typically release energy as two reactants join to form a single, more stable product.

Source: Chapter 1, Section 1.2.1 (Combination Reaction) and Section 1.2 (Decomposition Reaction)

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• Examiners expect all three equations balanced with correct state symbols.
• "Decomposition" must be named for Step 1 and "combination" for Step 3 — reversing these loses marks.
• For part (c), link the heat release explicitly to the exothermic nature of combination reactions — this is the reasoning the examiner wants, not just a label. The textbook Activity 1.4 is the direct source for this observation.