Class 10 Science Chapter 3 Previous Year Questions - Metals and Non-metals

Ans: (b) Calcium and magnesium
Explanation:

• Highly reactive metals like sodium, calcium, magnesium, and aluminium are extracted by electrolytic reduction of their molten chlorides.
• Gold and silver are less reactive, often found native or extracted by other methods (e.g., cyanide process).
• Iron is extracted by reduction with carbon in a blast furnace.
• Calcium and magnesium fit the description as they require electrolysis (e.g., NaCl or CaCl₂).

Ans: (a) Mg · O → Mg²⁺[O²⁻]
Explanation:

• Magnesium (Mg, atomic number 12: 2,8,2) loses 2 electrons to form Mg²⁺.
• Oxygen (O, atomic number 8: 2,6) gains 2 electrons to form O²⁻.
• The ionic compound is MgO (Mg²⁺[O²⁻]), with a 1:1 ratio.
• Option (a) correctly shows the electron transfer and ionic bond formation.

Ans: (a) Reduction with carbon

Explanation:

• Metals in the middle of the reactivity series (e.g., iron, zinc) are commonly extracted by reducing their oxides with carbon 
  (e.g., Fe₂O₃ + 3C → 2Fe + 3CO).
• Hydrogen reduction is less common, used for specific metals like tungsten.
• Aluminium reduction (thermite) is specific to certain metals (e.g., chromium).
• Electrolytic reduction is for highly reactive metals (e.g., sodium, aluminium).

Ans: (c) It has weak electrostatic forces of attraction between its oppositely charged ions.
Compound C is ionic (A loses electrons, B gains, forming A⁺ and B⁻).
Ionic compounds have:

• High melting points due to strong electrostatic forces (a).
• High solubility in water for many ionic compounds (b).
• Conductivity in molten or aqueous state due to mobile ions (d).
• Weak electrostatic forces (c) are not characteristic, as ionic bonds are strong.

Ans: (b) When it is heated with iron (III) oxide, molten iron is obtained.

• Thermite reaction: 2Al + Fe₂O₃ → 2Fe + Al₂O₃ + Energy.
• Aluminium reduces iron oxide to molten iron, which is used to join railway tracks.
• The reaction is exothermic (A is partially correct but less specific).
• Molten aluminium oxide (c) is a byproduct, not used for joining.
• No alloy is formed (D is incorrect).

Ans: (b) Al₂O₃ and MgO

• Aluminium burns in air: 4Al + 3O₂ → 2Al₂O₃ (aluminium oxide).
• Magnesium burns in air: 2Mg + O₂ → 2MgO (magnesium oxide).
• Al₃O₄ and MgO₂ are not chemically correct formulas.

Ans: (d) Zinc

• Reactivity series: K > Na > Ca > Mg > Al > Zn > Fe > Pb > Cu.
• Metal M displaces Fe from FeSO₄, so M is above Fe (Mg, Al, Zn).
• M does not displace Al from Al₂(SO₄)₃, so M is below Al.
• Zinc (Zn) fits: Zn + FeSO₄ → ZnSO₄ + Fe; no reaction with Al₂(SO₄)₃.

Ans: (b) Pale green

• Reaction: Fe + CuSO₄ → FeSO₄ + Cu.
• CuSO₄ solution is blue (Cu²⁺ ions).
• FeSO₄ solution is pale green (Fe²⁺ ions).
• After 1 hour, Cu²⁺ is replaced by Fe²⁺, changing the color to pale green.

Ans: (b) Ductility

• Ductility allows metals to be drawn into wires.
• Malleability is for hammering into sheets.
• Rigidity refers to stiffness, and resistivity is electrical resistance.

Ans: (a) Aluminium

• Aluminium forms a protective Al₂O₃ layer, preventing further corrosion.
• Copper forms a green patina (Cu₂(OH)₂CO₃), not just oxide.
• Silver and gold are resistant but do not rely on oxide layers.

Ans: (c) Assertion (a) is true, but Reason (R) is false.

• A: Most metals do not produce H₂ with HNO₃ because it’s a strong oxidizing agent, oxidizing H₂ to H₂O.
• R: Nitric acid is an oxidizing agent, not a reducing agent.
• Thus, A is true, R is false.

Ans: (b) Both Assertion (a) and Reason (R) are true, but Reason (R) is not the correct explanation of Assertion (a).

• A: Highly reactive metals (e.g., Na, K, Al) have strong bonds in oxides, requiring electrolysis, not carbon reduction.
• R: Displacement reactions (e.g., thermite) are used for some metals, but this does not explain A.

Ans: (d) Assertion (a) is false, but Reason (R) is true.

• A: Cooking utensils require malleability (shaping into sheets) and thermal conductivity, not ductility (wire drawing).
• R: Copper is both ductile and malleable, so R is true.

Ans: (c) Assertion (a) is true, but Reason (R) is false.

• A: Brass is made by melting copper and adding zinc (not tin).
• R: Copper is the primary metal in brass, but A mentions tin, making it incorrect. Corrected A would make R true, but as stated, R does not explain A.

Ans: (a) Both Assertion (a) and Reason (R) are true and Reason (R) is the correct explanation of the Assertion (a).

• A: AgCl turns grey due to formation of silver metal.
• R: 2AgCl → 2Ag + Cl₂ (photolytic decomposition), explaining the grey color.

Ans: (c) (i) and (ii)

• (i) Mg burns with a dazzling white flame, correct.
• (ii) Forms white MgO powder, correct.
• (iii) Mg does not significantly vaporize; it reacts with O₂.
• (iv) MgO forms Mg(OH)₂ in water, which is basic and turns red litmus blue, not blue to red.

Ans: (c) It is an endothermic reaction.

• 2Mg + O₂ → 2MgO is exothermic, releasing heat and light.
• (a), (b), and (d) are true: white flame, white MgO powder, combination reaction.

Ans: (a) Only (i)

• Reactivity series: Mg > Al > Zn > Fe > Pb > Cu.
• (i) Mg displaces Cu: Mg + CuSO₄ → MgSO₄ + Cu.
• (ii) Pb is below Fe, no reaction.
• (iii) Al is below Ca, no reaction.
• (iv) Ca is below Zn, no reaction.

Ans:
Safety Measures:

1. Wear safety goggles to protect eyes from the dazzling flame.
2. Use tongs to hold the ribbon to avoid burns.

Observations:

1. Burns with a dazzling white flame.
2. Forms a white powder (MgO).

Explanation:

• The bright flame can harm eyes, and the hot ribbon can cause burns.
• Reaction: 2Mg + O₂ → 2MgO, producing intense light and white ash.

Ans:
(a) Type: Displacement reaction.
Equation: Cu + 2AgNO₃ → Cu(NO₃)₂ + 2Ag
(b) Method: Electrolytic refining.

Explanation:

• Cu, more reactive than Ag, displaces Ag from AgNO₃.
• Electrolytic refining purifies silver using a silver anode and cathode in an electrolyte.

Ans:
Metal X: Mercury (Hg).
Equations:

1. 2HgS + 3O₂ → 2HgO + 2SO₂
2. 2HgO → 2Hg + O₂

Explanation:

• Cinnabar (HgS) is roasted to form mercury(II) oxide (HgO).
• HgO decomposes on heating to yield mercury metal.

Ans:
(i) Free State: Gold (Au), low in reactivity series (below Cu).
(ii) Compound: Aluminium (Al, as bauxite Al₂O₃), high in reactivity series (above Zn).

Explanation:

• Gold, being unreactive, exists native.
• Aluminium, reactive, forms stable compounds like oxides.

Ans: (a)

Formation:

• Mg (2,8,2) → Mg²⁺ (2,8) + 2e⁻
• Cl (2,8,7) + e⁻ → Cl⁻ (2,8,8) [×2 for 2Cl]
• Mg + 2Cl → MgCl₂ (Mg²⁺[Cl⁻]₂)

Cation: Mg²⁺; Anion: Cl⁻

Explanation:

• Mg donates 2 electrons to two Cl atoms, forming ionic MgCl₂.

Ans: (b)
Statement: Aluminium oxide (Al₂O₃) is amphoteric, reacting with both acids and bases.
Equations:

• With acid: Al₂O₃ + 6HCl → 2AlCl₃ + 3H₂O
• With base: Al₂O₃ + 2NaOH → 2NaAlO₂ + H₂O

Explanation:

• Al₂O₃ reacts with HCl to form a salt (like a base) and with NaOH to form sodium aluminate (like an acid), proving amphoteric nature.

Ans: (a)

Activity: Place three iron nails in test tubes:

• Tube A: Iron nail in water with air.
• Tube B: Iron nail in boiled water (no oxygen) with oil layer.
• Tube C: Iron nail in dry air (no moisture).

Observations:

• Tube A: Rust (Fe₂O₃·nH₂O) forms; requires water and oxygen.
• Tube B: No rust; oxygen absent.
• Tube C: No rust; moisture absent.

Conclusion: Iron rusts only with both water and oxygen.

Explanation:

• Rusting is a chemical reaction: 4Fe + 3O₂ + 2nH₂O → 2Fe₂O₃·nH₂O.
• Moisture and oxygen are essential for corrosion.

Ans:
(a) Justification:
Displacement reactions reduce metal oxides using more reactive metals.
Examples:

1. Fe₂O₃ + 2Al → 2Fe + Al₂O₃ (thermite, extracts iron).
2. Cr₂O₃ + 2Al → 2Cr + Al₂O₃ (extracts chromium).

(b) Reason: Highly reactive metals (e.g., Na, K, Al) form very stable oxides due to strong metal-oxygen bonds, which carbon cannot break. Electrolysis is required.

Explanation:

• Middle-series metals (Fe, Zn) have less stable oxides, reducible by Al.
• High-series metals need stronger reducing conditions (electrolysis).

Ans:
Electron-dot structures:

• Sodium (Na): 2,8,1 → Na·
• Oxygen (O): 2,6 → :Ö:

Formation of Na₂O:

• 2Na· → 2Na⁺ + 2e⁻
• :Ö: + 2e⁻ → [Ö]²⁻
• 2Na⁺ + [Ö]²⁻ → Na₂O

Cation: Na⁺; Anion: O²⁻

• Each Na donates 1 electron to O, forming ionic Na₂O.

Ans:
Activity for Conductivity:

• Take a metal rod (e.g., copper) and fix thumbtacks with wax at equal intervals.
• Heat one end with a burner.

Observation: Wax melts sequentially from the heated end, showing heat conduction.

Activity for Melting Point:

• Heat small pieces of copper and plastic in separate crucibles.

Observation: Plastic melts quickly; copper requires much higher heat, indicating a high melting point.

Explanation:

• Metals conduct heat due to free electrons.
• High melting points result from strong metallic bonds.

Ans:
(i) Least Reactive: D
(ii) Observation: C displaces Cu from CuSO₄ (CuSO₄ + C → CSO₄ + Cu).
(iii) Order: B > C > A > D

Explanation:

• Reactivity series: More reactive metal displaces less reactive one.
• D: No reactions, least reactive.
• A: Displaces Cu, below Zn and Al.
• C: Displaces Fe, below Al, likely displaces Cu.
• B: Displaces Fe, Cu, Zn, most reactive.

Ans:
Method: Electrolytic refining.
Description:

• Impure metal is the anode, pure metal is the cathode.
• Electrolyte is a solution of the metal’s salt (e.g., CuSO₄ for copper).
• On passing current, metal ions from the anode dissolve, deposit as pure metal on the cathode.
• Impurities (sludge) settle at the bottom.

Explanation:

• Used for metals like Cu, Al, ensuring high purity.

Ans: (b)
(a) Formation of AlN:

• Al (2,8,3) → Al³⁺ (2,8) + 3e⁻
• N (2,5) + 3e⁻ → N³⁻ (2,8)
• Al³⁺ + N³⁻ → AlN

(b) Reason: Ionic compounds have strong electrostatic forces forming a rigid lattice, but pressure displaces ions, causing like charges to repel, leading to brittleness.
Explanation:

• Al donates 3 electrons to N, forming ionic AlN.
• Ionic lattices break when lattice alignment is disrupted.

Ans:
(a) Reactivity Series: A list of metals arranged in order of decreasing reactivity.

• Development: Based on displacement reactions and tendency to lose electrons.

• Order: Ca > Al > Pb > Cu

(b) Equation: Fe₂O₃ + 2Al → 2Fe + Al₂O₃

Explanation:

• Reactivity is tested by observing which metals displace others from solutions.
• Thermite reaction reduces Fe₂O₃ to iron.

Ans: (a) (i)

• I: Ag (silver, unreactive).
• II: Al (forms protective Al₂O₃).
• III: K (catches fire spontaneously).
• IV: Cu (forms black CuO on heating).

(ii) Amphoteric Oxides: React with both acids and bases.

• Al₂O₃ + 6HCl → 2AlCl₃ + 3H₂O
• Al₂O₃ + 2NaOH → 2NaAlO₂ + H₂O

(iii) Alkalis: Water-soluble bases. Example: NaOH.

Explanation:

• Al₂O₃ shows dual reactivity, confirming amphoteric nature.
• Alkalis dissolve in water, producing OH⁻ ions.

Ans:
(I) Changes:

• Electrical conductivity decreases due to disruption of metallic lattice by alloying element.
• Melting point usually decreases as alloys have less ordered structures.
  (II) Alloy: Solder; Constituents: Lead, Tin.
  (III)(a) Alloys: Mixtures of two or more metals (or metal and non-metal) to enhance properties.
• Brass Preparation: Melt copper, add zinc in specific proportions, and cool.

Explanation:

• Alloys improve strength and corrosion resistance.
• Solder’s low melting point aids welding.
• Brass (Cu + Zn) is harder than pure copper.

Ans: (c)

• Mercury is a metal that is liquid at room temperature.
• Bromine is a non-metal that is also liquid at room temperature.
• Thus, the correct answer is (c) Mercury and Bromine.

Ans: (c)

Oxides of aluminum and zinc are classified as amphoteric. 
This means they can react with both acids and bases, displaying the following characteristics:

• They can act like acids in some reactions.
• They can behave like bases in other reactions.
• This versatility makes them important in various chemical processes.

Ans: (c)
The correct answer is (c) Copper and silver. These metals can be found in their pure, uncombined form (free state) in nature, as well as in compounds with other elements (combined state). This means they can exist both as standalone metals and as part of various minerals.

Ans: (b)

 Metal ‘X’ is aluminium, used in the thermit process. When aluminium is heated with oxygen, it produces aluminium oxide (Y), which is amphoteric. This means it can react with both acids and bases.

• Metal X: Aluminium (Al)
• Oxide Y: Aluminium oxide (Al₂O₃)
• Nature: Amphoteric

Ans: (i) The cathode is made of pure copper and the anode is made of impure copper.
(ii)) The solution used in this process is acidified copper sulphate, with the formula CuSO₄.
(iii) (A)  When electric current is passed through the electrolytic cell:

• Pure copper from the anode dissolves into the electrolyte.
• An equivalent amount of pure copper is deposited on the cathode.
• Soluble impurities enter the solution.
• Insoluble impurities settle at the bottom of the anode, forming anode mud.

• Beaker A: The blue colour of the solution fades or becomes colourless.
• Reason: Zinc is more reactive than copper, displacing it from the solution.
• Beaker B: No change in colour is observed.
• Reason: Silver is less reactive than copper, so no reaction occurs.

Ans: (b)
Assertion (A) is correct because different metals react differently with water and acids. However, while Reason (R) is also true, it explains how we extract metals rather than directly explaining why their reactivities vary.

Ans: (a) Amphoteric oxide (zinc oxide) reacts with acids as well as bases to produce salt and water.  

• Reaction with acids: Zinc oxide reacts with acids to form a salt and water. For example:

  ZnO+2HCl→ZnCl₂+H₂O

• Reaction with bases: Zinc oxide reacts with strong bases to form a salt (zincate). For example:

  ZnO+2NaOH+H₂O→Na₂[Zn(OH)₄]
  (b) If kept in open, sodium metal reacts vigorously with air and catches fire / kerosene oil does not allow sodium to come in contact with air and catch fire.
  (c) Nitric acid is a strong oxidising agent. It oxidises the hydrogen produced in the reaction to water. This is due to the strong oxidizing nature of nitric acid, which reduces itself to nitrogen oxides (like NO, NO₂) while oxidizing the hydrogen.

Ans: (a) (i) Reduction Process- Roasting
Reason: Mercury has low reactivity.
(ii) Reduction Process- Roasting
Reason: Copper has low reactivity.
(iii) Reduction Process- Electrolytic Reduction.
Reason: Sodium has high reactivity.
OR
(b) (i) Change in appearance - White to black colour.
Reason: Silver sulphide is formed.
(ii) Change in appearance – Reddish brown to green colour.
Reason: Basic Copper Carbonate is formed.
(iii) Change in appearance- Grey to brown colour.
Reason: Rust (iron oxide) is formed.

Ans: Ore of mercury: Cinnabar

Form in nature: Sulphide ore

The extraction of mercury from cinnabar involves the following chemical reactions:

• When heated in air, cinnabar (HgS) converts to mercuric oxide (HgO):
• 2HgS(s) 3O₂(g) → 2HgO(s) 2SO₂(g)
• Further heating of mercuric oxide produces mercury:
• 2HgO(s) → 2Hg(l) O₂(g)

Ans: Zn + H₂SO₄ → ZnSO₄ + H₂ (g)

To test for the presence of hydrogen gas:

• Bring a burning matchstick near the gas.
• If hydrogen is present, it will burn with a pop sound.

Ans: (c)
Assertion (A): A piece of zinc metal gets a reddish-brown coating when kept in copper sulfate solution for some time. This is true. When zinc is placed in a copper sulfate solution, a displacement reaction occurs where zinc displaces copper from copper sulfate. The copper then deposits on the zinc surface as a reddish-brown coating.
Reason (R): Copper is more reactive than zinc. This is false. In the reactivity series, zinc is more reactive than copper, which is why zinc can displace copper from copper sulfate.
Since Assertion (A) is true but Reason (R) is false, the correct answer is (c) (A) is true, but (R) is false

Ans: (d) Na₂CO₃
Na₂CO₃ ⋅ 10H₂O (hydrated sodium carbonate) has the largest concentration of water molecules in its crystalline structure, with 10 molecules per formula unit.

Ans: (i) Transfer of electrons during the creation of magnesium chloride:
(ii) The two properties of ions Compounds are:
(a) Solubility: Electrovalent compounds are generally soluble in water and insoluble in solvents such as kerosene, petrol, etc.
(b) Conduction of Electricity: 

• The conduction of electricity through a solution involves the movement of charged particles.
• A solution of an ionic compound in water contains ions, which move to the opposite electrodes when electricity is passed through the solution.
• Ionic compounds in the solid state do not conduct electricity because movement of ions in the solid is not possible due to their rigid structure.
• But ionic compounds conduct electricity in the molten state.
• This is possible in the molten state since the electrostatic forces of attraction between the oppositely charged ions are overcome due to the heat.
• Thus, the ions move freely and conduct electricity.

• Sodium atom has one electron in its outermost shell. 
•  If it loses the electron from its M shell then its L shell now becomes the outermost shell and that has a stable octet. 
•  The nucleus of this atom still has 11 protons but the number of electrons has become 10, so there is a net positive charge giving us a sodium cation Na⁺. 
•  On the other hand chlorine has seven electrons in its outermost shell and it requires one more electron to complete its octet. 
•  If sodium and chlorine were to react, the electron lost by sodium could be taken up by chlorine. 
•  After gaining an electron, the chlorine atom gets a unit negative charge, because its nucleus has 17 protons and there are 18 electrons in its K, L and M shells. 
•  This gives us a chloride anion Cl^–. 

• The conduction of electricity through a solution involves the movement of charged particles.
• A solution of an ionic compound in water contains ions, which move to the opposite electrodes when electricity is passed through the solution.
• Ionic compounds in the solid state do not conduct electricity because the movement of ions in the solid is not possible due to their rigid structure.
• But ionic compounds conduct electricity in the molten state.
• This is possible in the molten state since the electrostatic forces of attraction between the oppositely charged ions are overcome due to the heat.
• Thus, the ions move freely and conduct electricity.

Ans: (a) An ore of mercury is cinnabar, and mercury is present in it in the form of mercury sulfide (HgS).
(b) When zinc carbonate (ZnCO₃) is heated strongly in a limited supply of air, it undergoes thermal decomposition to produce zinc oxide (ZnO), carbon dioxide (CO₂), and water (H₂O):
ZnCO₃(s) → ZnO(s) + CO₂(g)
(c) (I)  The metal $\mathrm{ A\; }$is aluminium (Al), and the reaction is called the thermite reaction.
(II) The chemical equation for the reaction of iron (A) with iron(III) oxide (Fe₂O₃) is:
2Al(s) + Fe₂O₃(s) → 2Fe(s) + Al₂O₃(s)
This reaction is highly exothermic and is used in various industrial applications, including joining railway tracks due to its high heat generation and the ability to melt and fuse metals.

Ans: (b)
In the thermite process, aluminum (Al) is used because of its high reactivity and ability to reduce metal oxides, such as iron oxide, to produce molten iron.
When aluminum is burnt in air, it forms aluminum oxide (Al₂O₃), which is an amphoteric oxide (meaning it can react with both acids and bases).
Therefore, X is aluminum (Al) and Y is aluminum oxide (Al₂O₃), making the correct answer (b) Al and Al₂O₃.

Ans: (A) Copper is a reactive element. When it is heated in the air, it forms black copper oxide (CuO).
(B) Metal oxides are categorised as amphoteric oxides that react with both acids as well as bases to create salts and water. Amphoteric oxides, among many others, include lead oxide and zinc oxide. These oxides are oxygen compounds that show both acidic and basic characteristics. These undergo a neutralisation reaction to form water and salt.
Example: 
(C) 

OR

(C) (i)

(ii) The gas formed is sulphur dioxide which is colourless and poisonous. 
(iii) The nature of the gas is acidic. 
(iv) SO₂ gas has no effect on dry litmus paper because it shows acidic behavior only in the presence of water

Ans: When a clear solution of slaked lime (Ca(OH)₂) is exposed to air, carbon dioxide (CO₂) from the atmosphere dissolves into the solution and reacts with the dissolved calcium hydroxide. This reaction produces calcium carbonate (CaCO₃), which is insoluble and appears as a faint white precipitate (“milky” appearance). The chemical equation is:
Thus, the “milkiness” arises from the formation of solid calcium carbonate.

1. Copper (below H)

   • Does not react with dilute HCl to give hydrogen gas.
   • No visible reaction under normal conditions.

2. Iron (above H)

   • Does react to produce hydrogen gas.
   • Reaction:  Fe(s) + 2HCl(aq) → FeCl₂(aq) + H₂(g)

3. Magnesium (above H)

   • Does react quite vigorously, giving hydrogen gas.
   • Reaction: Mg(s) + 2HCl(aq) → MgCl₂(aq) + H₂(g)

4. Zinc (above H)

   • Does react, evolving hydrogen gas.
   • Reaction: Zn(s) + 2HCl(aq) → ZnCl₂(aq) + H₂(g)

Therefore, corrected table becomes :

So, the correct observations are for  iron and zinc only.

Ans: 

• Potassium is kept immersed in kerosene because it reacts vigorously with air and water.
• The reaction with air forms a layer of potassium oxide on the surface, while the reaction with water produces potassium hydroxide.
• Both of these reactions release hydrogen gas, which is highly flammable.
• By keeping potassium immersed in kerosene, it is protected from air and water, preventing any unwanted reactions.

Ans: (a) Gold and silver are highly malleable metals so, they are used for making jewellery.
(b)  Metals used to make cooking utensils are chosen not just for their good heat conduction but also for their relatively low reactivity. You generally don’t want a metal that reacts with food or with the conditions (high temperature, presence of water vapor, mild acids in food, etc.). Here’s how it ties to reactivity:

1. Metals below hydrogen in the reactivity series (like copper, silver, gold) do not readily react with acids or water under normal cooking conditions.
2. Metals that form stable oxide layers (like aluminum, which is higher in the series) are protected from further corrosion by that oxide coating.
3. Alloys such as stainless steel (predominantly iron but mixed with chromium/nickel) also resist corrosion because chromium forms a very adherent oxide layer that protects the iron underneath.

Thus, metals like iron (in stainless steel), aluminum, and copper are commonly used for utensils because:

• They do not corrode easily or leach into food.
• They are good conductors of heat.
• They maintain a protective surface layer (oxide or alloy layer), preventing further reaction

Ans:

Ans: 

(b) The formation of calcium oxide (CaO) involves the transfer of electrons. Calcium (Ca) donates two electrons to oxygen (O) to form Ca²⁺ cation and O^2- anion. The ionic bond is formed between these ions to create calcium oxide.
(c) In calcium oxide (CaO), the ions present are Ca²⁺ (calcium cation) and O^2- (oxygen anion).
(d) Four important characteristics of calcium oxide (CaO) are:

• It is a white, crystalline solid.
• It has a high melting and boiling point.
• It is an ionic compound.
• It is commonly used as a desiccant and in cement production.

Ans: Three differentiating features between the processes of galvanisation and alloying are as follows:

• Definition: Galvanisation is the process of applying a protective zinc coating to iron or steel to prevent corrosion, while alloying is the process of combining two or more metals or metals with a non-metal to create a new material with enhanced properties.
• Purpose: Galvanisation is primarily used to protect the base metal from corrosion by providing a sacrificial layer of zinc, whereas alloying is done to improve specific properties of the metal, such as strength, hardness, or resistance to corrosion.
• Procedure: In galvanisation, the metal is coated with zinc through processes like hot-dip galvanisation or electroplating. On the other hand, alloying involves melting and mixing different metals or adding non-metallic elements to create an alloy with the desired properties.

Ans: (a) Complete and balance the following chemical equations:
(i) AI₂O₃ + 6HCl → 2AlCl₃ + 3H₂O
(ii) K₂O + H₂O → 2KOH
(iii) Fe + 2H₂O → Fe(OH)₂ + H₂
(b) Observations:

• When element 'X' is treated with copper sulphate solution, no reaction occurs.
• When element 'X' is treated with zinc sulphate solution, 'X' displaces zinc from the solution, resulting in the formation of 'X' sulphate and zinc metal.
• When element 'X' is treated with silver nitrate solution, 'X' displaces silver from the solution, resulting in the formation of 'X' nitrate and silver metal.

Based on the observations, the increasing order of reactivities is:
Ag<Cu<Zn<X

Ans: 

(ii) Ionic compounds are solids because the particles that make up ionic compounds are held together by strong electrostatic bonds.

Ans: (a) (i) The malleability and ductility properties of gold make it suitable for ornaments. Also, beacause of less reactive natureof gold, it's prefered for making jewellery. 
(ii) Silver and gold. 
(iii) Gallium and caesium have so low melting points that they melt even on keeping them in the palm.

(b) Formation of the ionic compound CaO (calcium oxide) with an electron-dot structure:

Calcium oxide (CaO) is formed through the transfer of electrons between calcium (Ca) and oxygen (O) atoms. Calcium, with an atomic number of 20, loses two electrons to achieve a stable electron configuration similar to that of a noble gas (argon), while oxygen, with an atomic number of 8, gains two electrons to complete its valence shell.

The electron-dot structure of CaO can be represented as follows:

In this structure, the calcium atom loses its two valence electrons to oxygen, resulting in a Ca²⁺cation and an O^2- anion. The electrostatic attraction between these oppositely charged ions forms the ionic bond in calcium oxide (CaO).

Ans: Sodium, magnesium and aluminium have a higher affinity towards oxygen than carbon because these are highly reactive metals. Hence, carbon cannot reduce the oxides of sodium, magnesium and aluminium to their respective metals. These metals are placed at the top of the reactivity series. Highly reactive metals like Na, Mg, Al, etc. are extracted by electrolytic reduction of their molten chlorides or oxides. Electrolytic reduction is brought about by passing electric current through the molten state. Metal gets deposited at the cathode.
NaCl ⇌ Na⁺ + Cl^–
At cathode : Na⁺ + e^– → Na⁺
At anode : 2Cl^– → Cl₂ + 2e^–

Ans: (i) On heating, mercuric oxide decomposes to give mercury and oxygen.

(ii) On heating a mixture of cuprous oxide and cuprous sulphide, copper and sulphur dioxide are produced.

(iii) When aluminium is heated with manganese dioxide, manganese and aluminium oxide are formed.

(iv) Ferric oxide reacts with aluminium to produce aluminium oxide and iron.

(v) On calcination, zinc carbonate produces zinc oxide and carbon dioxide.

Ans: (a) (i) Sodium
(ii) Iodine
(iii) Mercury
(iv) Gold
(v) Silver
(vi) Carbon
(b) Alloys offer several advantages over pure metals:

• They are generally stronger.
• They resist corrosion better.
• They often have lower melting points.
• They usually have lower electrical conductivity. 
  The composition of:
• Solder: An alloy of lead and tin.
• Amalgam: An alloy of mercury with another metal.

Ans: After approximately 30 minutes, the student would observe that the strip of iron would start to show signs of corrosion or rusting. This is because iron is more reactive than copper, aluminium, and zinc and will displace the iron from the iron sulphate solution, forming iron oxide (rust) on its surface.

Explanation: When cleaned metal strips of aluminium (Al), copper (Cu), iron (Fe), and zinc (Zn) are each placed into separate beakers containing freshly prepared iron(II) sulphate (FeSO₄) solution, here’s what the student would observe after about 30 minutes:

1. Aluminium strip (Al)

   • Above Fe in the reactivity series:$Al$
     Al+FeSO₄⟶Al₂(SO₄)₃+Fe
   • Expected change: A brownish deposit of iron on the aluminium strip and possible fading of the greenish FeSO₄ solution.

2. Copper strip (Cu)

   • Below Fe in the reactivity series:
     • Copper cannot displace iron from its salt.
   • Expected change: No reaction; the copper strip remains unchanged, and the solution color stays the same.

3. Iron strip (Fe)

   • Same metal as in the solution:
     • No net displacement reaction occurs because the metal and the dissolved ion are both iron.
   • Expected change: No reaction; no visible change in the metal strip or the solution.

4. Zinc strip (Zn)

   • Above Fe in the reactivity series:$Zn$Zn+FeSO₄⟶ZnSO₄+ Fe
   • Expected change: A brownish deposit of iron on the zinc strip and the FeSO₄ solution may lose its greenish color as Fe²⁺ ions are displaced.

• Al in FeSO₄: Brownish deposit (iron) forms on the strip.
• Cu in FeSO₄: No change.
• Fe in FeSO₄: No change.
• Zn in FeSO₄: Brownish deposit (iron) forms on the strip.

Ans: (a) The colour of the ferrous sulphate solution in test tube (I) will match the colour of the solution in test tube II when a piece of copper is dropped in it. This is because copper is less reactive than iron and will not displace iron from ferrous sulphate solution, resulting in no change in the colour of the solution.
No reaction: Cu + FeSO₄ → CuSO₄ + Fe
(b) The colour of the ferrous sulphate solution will fade and a black mass will form on the metal surface in the following cases:

1. Zinc displacing iron:

   $Zn$Zn+FeSO₄ ⟶ ZnSO₄+Fe (s)
   • The FeSO₄ solution fades because Fe²⁺ ions are replaced by Zn²⁺.
   • The displaced iron is deposited on the zinc strip, often appearing black or brown in color.

2. Aluminium displacing iron:

   $\hspace{0.17em}Fe (s)2\backslash ,\backslash text\{Al\}\; +\; 3\backslash ,\backslash text\{FeSO\}\_\{4\}\; \backslash ;\backslash longrightarrow\backslash ;\; \backslash text\{Al\}\_\{2\}(\backslash text\{SO\}\_\{4\})\_\{3\}\; +\; 3\backslash ,\backslash text\{Fe\; (s)\}$2Al+3FeSO₄⟶Al₂(SO₄)₃+3Fe (s)
   • Similarly, the pale green color of FeSO₄ solution fades as iron ions are displaced.
   • The black/brown deposit of iron forms on the aluminium strip.

Ans: The ore that produces brisk effervescence when treated with dilute hydrochloric acid is a carbonate ore. One example of a carbonate ore is limestone (CaCO₃).
Steps to obtain metal from the enriched ore:

• Calcination: The ore is heated strongly in the absence of air to convert it into metal oxide.
  CaCO₃(s) → CaO(s) + CO₂(g)
• Reduction: The metal oxide is then reduced using a suitable reducing agent, such as carbon or hydrogen, to obtain the metal.
  CaO(s) + C(s) → Ca(s) + CO(g)

Ans:  

•  The substances present in the air with which silver reacts are hydrogen sulphide (H₂S) and sulphur compounds. Silver reacts with these substances to form a black layer of silver sulphide (Ag₂S) on its surface. 
•  The substances present in the air with which copper reacts are moisture, carbon dioxide (CO₂), and oxygen (O₂). Copper reacts with these substances to form a green layer of copper carbonate (CuCO₃) or copper hydroxide (Cu(OH)₂) on its surface. 

Ans: (a)

(b) (i) Metals conduct electricity due to the flow of free electrons present in them.
(ii) The reaction of iron (III) oxide, Fe₂O₃ with aluminium is highly exothermic and the iron produced melts. This molten iron is used to join cracked iron parts of machines and railway tracks.

Ans: (a)  The reactivity series is a list of metals arranged in order of their decreasing reactivity. It helps predict how metals will react with each other and with other substances.

Key points about the reactivity series:

• Metals higher in the series are more reactive.
• For example, if metal A can displace metal B from a solution, A is more reactive than B.

(b) The metals in the middle of the reactivity series are obtained from their ores by chemical reduction with a suitable reducing agent, e.g.
ZnO + C → Zn + CO
The metals at the top of the series are obtained by electrolytic reduction of their molten orc.
Al₂O₃  → 2Al³⁺ + 3O^2-
At cathode: 2Al³⁺ + 6e^- → 2Al
At anode: O^2- - 2e^- → O
O + O → O₂

Ans: (a) Extraction of metals of medium reactivity:

• The metals in the middle of the reactivity series include zinc, iron, and lead.
• These metals are typically found in the form of carbonates or sulphides.
• To extract the metal, the carbonate ores must first be converted to oxides, as it is easier to obtain metals from their oxides.

(b) Copper glance (Cu₂S) when heated in air gets partially oxidised to copper oxide which further reacts with the remaining copper glance to give copper metal.

Electrolytic refining of copper:

• The process uses an electrolyte of acidified copper sulphate.
• The anode is made of impure copper, while the cathode is a strip of pure copper.
• When electric current is passed, pure copper is deposited on the cathode.
• Soluble impurities dissolve in the solution, and insoluble impurities settle as anode mud.

Ans: To find out the conditions under which iron rusts, you can perform the following activity:

• Take three test tubes and label them as A, B, and C.
• Fill test tube A with water, test tube B with water and oil, and test tube C with water and salt.
• Place a small piece of iron nail in each test tube and allow them to stand undisturbed for a few days.
• Observe the test tubes regularly and note any changes in the appearance of the iron nails.
• After a few days, check for the presence of rust on the iron nails in each test tube.
• Analyze the results and determine the conditions under which iron rusts.

Ans: (i) The type of bond in compound Z is an ionic bond.
(ii) Ionic compounds generally have high melting and boiling points due to the strong electrostatic forces of attraction between the positive and negative ions.
(iii) Ionic compounds like compound Z do not dissolve in non-polar solvents like kerosene or petrol. They are only soluble in polar solvents.
(iv) No, compound Z will not be a good conductor of electricity in a solid state because the ions are held in a fixed position and cannot move. However, it may conduct electricity when dissolved in water or molten state as the ions become free to move and carry electric charge.

Ans: When metal surfaces are exposed to moist air, they can undergo oxidation. This process results in the formation of metal oxides on the surface. The effects on different metals are as follows:

• Copper: The surface develops a greenish coating known as copper(II) oxide (CuO), which gives it a dull appearance.
• Silver: The surface forms a blackish coating called silver sulfide (Ag₂S), also resulting in a dull look.

Ans:(a) The process of slowly eating up metals due to their conversion into oxides, carbonates, sulphides, etc., by the action of atmospheric gases and moisture is called corrosion.
(b) The corrosion of iron is called rusting.
(c) Silver articles become black after some time when exposed to air. This is due to the formation of a coating of black silver sulphide (Ag₂S) on its surface by the action of H₂S gas present in the air.
(d) Corrosion of iron is a serious problem. Every year large amount of money is spent to replace damaged iron articles. Corrosion causes damage to car bodies, bridges iron railings, ships and to all objects made of metals especially those of iron.
(e) Corrosion of iron is prevented by coating it with a layer of oil. The reason is that the layer of oil does not allow air and water to reach the surface of iron. Corrosion of iron can also be prevented by painting, greasing, galvanising, anodising, electroplating or making alloys. 

Ans: (i) Most metals do not evolve hydrogen gas when they react with nitric acid because nitric acid is a strong oxidizing agent. It oxidizes hydrogen gas produced during the reaction to water and itself gets reduced to nitrogen oxides. Therefore, instead of hydrogen gas, the products obtained are nitrogen oxides.
(ii) Zinc oxide is considered an amphoteric oxide because it shows both acidic and basic properties. It reacts with acids to form zinc salts and water, exhibiting basic characteristics. Additionally, it reacts with bases to form zincates and water, showing acidic properties. This ability to react with both acids and bases classifies it as an amphoteric oxide.

ZnO + 2HCl →  ZnCl₂   + H₂ O

ZnO + 2NaOH + H₂ O →  Na₂   (Zn (OH)₄ )

(iii) Metals conduct electricity because they have a large number of free or delocalized electrons. These electrons are not bound to any particular atom and can move freely throughout the metal lattice. When a potential difference is applied across a metal, these free electrons can easily move and carry an electric current.

Ans: The reverse of the given chemical reaction is not possible because it violates the principle of conservation of mass. In the forward reaction, zinc (Zn) displaces copper (Cu) from copper sulfate (CuSO₄) to form zinc sulfate (ZnSO₄) and copper. However, in the reverse reaction, copper cannot displace zinc from zinc sulfate, as copper is less reactive than zinc. Hence, the reverse of this reaction is not feasible.

Ans: (a) It is a process in which a metal reacts with substances present in the air to form surface compounds.
(b) Rusting.
(c) Black layer on its surface due to formation of Ag₂S.
(d) It makes the metal weak and brittle, which is a serious problem.
(e) Oiling, painting, greasing, galvanisation, and alloying can prevent iron from corrosion.

Ans: (a) When a metal is attacked by substances around it, such as moisture, acids, etc., it is said to corrode and this process is called corrosion. Corrosion of iron is called rusting.
(b) Coating formed on silver is black and that formed on copper is green.
(c) Damages caused by corrosion:

• It causes damage to car bodies and bridges.
• It damages iron railings and ships.

Prevention of corrosion:

• Corrosion can be prevented by oiling, painting, greasing and galvanising.
• Corrosion can be prevented by galvanising and alloying.

Ans: (i) It is easier to convert metal oxides to metals as compared to carbonates and sulphides. Therefore carbonates are calcinated and sulphides are roasted to oxides.
(ii) There are electrostatic forces of attraction between the cations and anions in ionic compounds which are difficult to break. Therefore ionic compounds have high melting points.
(iii) Metals above hydrogen evolve hydrogen gas and metals below hydrogen do not evolve hydrogen when treated with an acid. That is when hydrogen has been assigned a place in the reactivity series of metals.
(iv) Even if the zinc layer is broken, it will preferentially be oxidised because it is more reactive than iron.
(v) The wires carrying current in homes have a coating of PVC. This is because PVC is an insulating substance and protects from electric shock.

Ans: Carbonic acid is formed.
CO₂ + H₂O → H₂CO₃

Ans: 2Al(s) + 3H₂O (g) → Al₂O₃(s) + 3H₂(g)

Ans:

(b) It is an ionic compound.
Physical properties:
(i) It is hard and solid.
(ii) It has a high melting and boiling point.
(iii) It is soluble in water.

Ans: (A) Sodium and magnesium have a tendency to react with oxygen rather than carbon because these are highly reactive metals. They have a greater affinity for oxygen than for carbon. Hence, their oxides are stable. The reduction of these metallic oxides with carbon requires very high temperature and at that temperature, metals react with carbon to form their corresponding carbides. Hence, carbon cannot reduce the oxides of Na or Mg. 
(B) Galvanisation is a process of formation of thin layer over metal surface. It prevents further contact of metal surface with atmosphere and reduces the corrosion level. So, iron articles are galvanised with a thin layer of zinc over them. Since zinc is more reactive than iron, it undergoes oxidation more readily than iron. As a result, iron articles remain protected. 
(C) Metals such as Na, K, Ca and Mg are highly reactive metals and hence they are not found in their free state in nature. Na, K, Ca and Mg  are alkali and alkaline Earth metals. They are the most reactive metals and readily react with atmospheric oxygen and other gases. Therefore, they are found in nature in the form of their compounds.

Ans: Silver and gold are least reactive metals and are often found in their native or free state. Aluminium is a very reactive metal and is never found in free state, but in combined state.