Two physical properties used to distinguish metals from non-metals are lustre (metals have a shiny surface) and electrical conductivity (metals are good conductors, non-metals are not).

Source: Chapter 3, Section 3.1 Physical Properties

The examiner expects any two valid physical properties from: lustre, malleability, ductility, conductivity of heat/electricity, or sonority. Lustre and electrical conductivity are the most commonly cited pair. Stating just the property name without a brief qualifier may cost half a mark, so include a one-phrase distinction for each.

(B) They can be beaten into thin sheets.

Metals possess the property of malleability — they can be beaten into thin sheets. Gold and silver are the most malleable metals. Metals are also good conductors of heat and electricity, and form basic (not acidic) oxides.

Source: Metals and Non-metals, Section 3.1.1

• The question asks for a physical property of metals.
• Option B describes malleability, which is a key physical property of metals confirmed in Activity 3.3.
• Option A is wrong — metals form basic oxides (non-metals form acidic oxides).
• Option C is wrong — metals are good conductors.
• Option D describes non-metals (brittle, dull).
• Remember: malleability and ductility are the two most commonly tested physical properties of metals in MCQs.

When sodium reacts with cold water, the two products formed are sodium hydroxide (NaOH) and hydrogen gas (H₂).

This is a standard 1-mark factual recall question. Examiners expect both products named clearly. The reaction is: 2Na + 2H₂O → 2NaOH + H₂↑. Note: the source passages do not explicitly state this reaction, but it is standard CBSE Class 10 Chapter 1 content on displacement reactions involving metals and water.

Potassium and sodium are highly reactive metals. When they react with cold water, the reaction is so vigorous and exothermic that the hydrogen gas evolved immediately catches fire.

In contrast, calcium's reaction with water is less violent and the heat evolved is not sufficient to ignite the hydrogen gas produced. Hence calcium does not catch fire.

$$2\text{Na(s)} + 2\text{H}_2\text{O(l)} \rightarrow 2\text{NaOH(aq)} + \text{H}_2\text{(g)} + \text{heat (large)}$$
$$\text{Ca(s)} + 2\text{H}_2\text{O(l)} \rightarrow \text{Ca(OH)}_2\text{(aq)} + \text{H}_2\text{(g)} + \text{heat (less)}$$

Source: Chapter 3, Section 3.2.2

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The key distinction examiners look for is reactivity/heat evolved, not the gas produced (all three produce H₂). Award-winning answers state that K and Na react violently/exothermically enough to ignite H₂, while Ca's reaction is less exothermic so H₂ does not catch fire. Including the equations earns full marks for a 2-mark question.

(B) No reaction takes place.

Copper is less reactive than hydrogen in the activity series, so it cannot displace hydrogen from dilute HCl; hence no reaction occurs.

Examiners expect students to apply the reactivity series: only metals above hydrogen (e.g., Fe, Zn, Mg) can displace hydrogen from dilute acids. Copper lies below hydrogen in the series, so it does not react with dilute HCl. This is a standard application question from Chapter 1 (Chemical Reactions and Equations).

HNO₃ is a strong oxidising agent. When a metal reacts with nitric acid, the hydrogen gas produced is immediately oxidised to water by HNO₃, which itself gets reduced to nitrogen oxides (N₂O, NO, NO₂). HCl is not an oxidising agent, so hydrogen gas is released freely.

(Exception: Mg and Mn react with very dilute HNO₃ to produce H₂.)

Source: Chapter 3, Section 3.2.3

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• The key phrase examiners look for: "HNO₃ is a strong oxidising agent" — this must appear.
• Then explain the consequence: H₂ is oxidised to water, and HNO₃ is reduced to nitrogen oxides.
• Contrast with HCl: HCl has no oxidising property, so H₂ escapes.
• Mentioning the exception (Mg/Mn with very dilute HNO₃) is a bonus point but not compulsory for 2 marks.

Observation 1: When iron filings are added to copper sulphate solution, the blue colour of the solution fades and a reddish-brown deposit of copper is formed on the iron. The reaction is:
$$\text{Fe}(s) + \text{CuSO}_4(aq) \rightarrow \text{FeSO}_4(aq) + \text{Cu}(s)$$

Observation 2: When copper pieces are added to iron sulphate solution, no reaction occurs — there is no change in the colour of the solution and no deposit is formed.

Conclusion: Iron displaces copper from copper sulphate solution, but copper cannot displace iron from iron sulphate solution. This shows that iron is more reactive than copper.

Source: Chapter 1, Section 1.2.3 – Displacement Reaction

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• The key point examiners look for is both observations (one reaction, one no reaction) plus the equation for the first case, and the conclusion about relative reactivity.
• Many students forget to state that copper + FeSO₄ gives no reaction — this is worth marks.
• The logic: a more reactive metal displaces a less reactive one from its salt solution. Since Fe displaces Cu but not vice versa, Fe > Cu in reactivity.

(i) Sodium loses one electron to form a sodium cation with a +1 charge (Na⁺), and chlorine gains that electron to form a chloride anion with a −1 charge (Cl⁻).

(ii) The oppositely charged Na⁺ and Cl⁻ ions are held together by a strong electrostatic force of attraction.

(iii) Compounds formed by the transfer of electrons from a metal to a non-metal are called ionic compounds (or electrovalent compounds).

Source: Chapter 3, Section 3.3 – How do Metals and Non-metals React?

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• For (i), always state both charges clearly — examiners check both Na⁺ (+1) and Cl⁻ (−1).
• For (ii), the key phrase is electrostatic forces of attraction — avoid vague terms like "magnetic force."
• For (iii), either "ionic compounds" or "electrovalent compounds" is acceptable; writing both is safest.
• This question tests understanding of electron transfer and ionic bond formation — a frequently asked concept in CBSE board exams.

In solid ionic compounds, the ions are held in fixed positions in a rigid crystal lattice and cannot move freely. Since electric current requires the movement of charged particles (ions), ionic solids do not conduct electricity.

When melted or dissolved in water, the ionic lattice breaks down and the ions become free to move. These freely moving ions carry electric charge and thus conduct electricity.

In short: Conductivity requires mobile charged particles. Solid state → ions fixed, no conduction. Molten/dissolved state → ions free to move → conduction occurs.

Source: Chapter 3 (Metals and Non-metals / Ionic compounds), as referenced in Chapter 4, Section 4.1

• Examiners look for two clear points: (1) why solids don't conduct — ions are fixed in lattice, (2) why molten/dissolved state does conduct — ions become free/mobile.
• Use the terms rigid lattice, free ions, and mobile charged particles for full marks.
• This is a 3-mark answer, so cover both states with a brief reason for each, plus a linking idea (conductivity needs mobile ions). Don't write more than 3–4 sentences.

All minerals are naturally occurring compounds in the earth's crust, but only those minerals from which a metal can be profitably extracted are called ores. All ores are minerals, but all minerals are not ores.

Source: Chapter 3, Section 3.4 Occurrence of Metals

The key distinction examiners look for is the profitability/feasibility of extraction — a mineral becomes an ore only when the metal can be extracted from it economically. Always include the phrase "profitably extracted" and the one-way relationship: all ores are minerals, but not all minerals are ores.

Highly reactive metals like sodium and aluminium have a greater affinity for oxygen than carbon does. Therefore, carbon cannot reduce their oxides — it is not a strong enough reducing agent to displace these metals from their oxides.

The method used to extract such metals is electrolytic reduction (electrolysis). For example, sodium is extracted by electrolysis of molten sodium chloride. At the cathode, sodium is deposited:

$$\text{Na}^+ + e^- \rightarrow \text{Na}$$

At the anode, chlorine is liberated:

$$2\text{Cl}^- \rightarrow \text{Cl}_2 + 2e^-$$

Similarly, aluminium is obtained by electrolytic reduction of aluminium oxide.

Source: Chapter 3, Section 3.4.5 — Extracting Metals towards the Top of the Activity Series

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• Why carbon fails: The key reason is affinity for oxygen — Na and Al hold oxygen more strongly than carbon does, so carbon cannot "pull" oxygen away from their oxides. State this clearly for full marks.
• Method name: "Electrolytic reduction" or "electrolysis" — both are accepted. Naming the method earns 1 mark; the reason earns 1–2 marks.
• Equations: Including the cathode/anode half-reactions adds precision and is expected at this level; they demonstrate understanding of the process.
• Avoid vague answers like "carbon is not strong enough" without linking it to oxygen affinity.

(i) Carbon can only reduce metal oxides to their respective metals. It cannot directly reduce sulphides or carbonates efficiently. Therefore, sulphide and carbonate ores must first be converted to metal oxides before reduction with carbon.

(ii)
(a) Sulphide ores are converted to oxides by roasting — heating strongly in the presence of excess air.
Example: $2\text{ZnS(s)} + 3\text{O}_2\text{(g)} \xrightarrow{\Delta} 2\text{ZnO(s)} + 2\text{SO}_2\text{(g)}$

(b) Carbonate ores are converted to oxides by calcination — heating strongly in limited air.
Example: $\text{ZnCO}_3\text{(s)} \xrightarrow{\Delta} \text{ZnO(s)} + \text{CO}_2\text{(g)}$

Source: Chapter 3, Section 3.4.4

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• Part (i) is worth ~1 mark — one clear reason is enough.
• Parts (ii)(a) and (ii)(b) are worth ~1 mark each — name the process AND give the equation; examiners expect both.
• Key terms to spell correctly: roasting (excess air) vs calcination (limited air) — mixing these up is a common error.