Chemical Reactions and Equations Exercise Solutions | KalviGavani

Magnesium is a highly reactive metal. When exposed to air, it slowly reacts with oxygen and forms a thin, dull layer of magnesium oxide (MgO) on its surface. This oxide layer acts as a protective coating that prevents the magnesium underneath from coming into direct contact with oxygen.

If we try to burn the ribbon without cleaning, the oxide layer interferes — the ribbon does not ignite easily and does not burn with its characteristic brilliant white flame. Cleaning the ribbon with sandpaper removes this MgO coating, exposing fresh magnesium metal so that it burns readily and brightly in air.

Balanced equation for the burning:

2Mg(s) + O₂(g) → 2MgO(s)

Observations: The magnesium ribbon burns with a dazzling white flame and leaves behind a white powdery ash of magnesium oxide.

(i) Hydrogen + Chlorine → Hydrogen chloride

Unbalanced: H₂ + Cl₂ → HCl

On the left: H = 2, Cl = 2. On the right: H = 1, Cl = 1. To balance, place coefficient 2 before HCl:

H₂(g) + Cl₂(g) → 2HCl(g)

This is a combination reaction — two elements combine to form one compound.

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

Unbalanced: BaCl₂ + Al₂(SO₄)₃ → BaSO₄ + AlCl₃

Step-by-step balancing:

• Ba: left = 1, right = 1. Place 3 before BaCl₂ and 3 before BaSO₄ → Ba = 3 on both sides.
• Al: left = 2 (in Al₂(SO₄)₃), right = 1 (in AlCl₃). Place 2 before AlCl₃ → Al = 2 on both sides.
• SO₄: left = 3 (in Al₂(SO₄)₃), right = 3 (in 3BaSO₄). Balanced.
• Cl: left = 3×2 = 6 (in 3BaCl₂), right = 2×3 = 6 (in 2AlCl₃). Balanced.

3BaCl₂(aq) + Al₂(SO₄)₃(aq) → 3BaSO₄(s)↓ + 2AlCl₃(aq)

This is a double displacement (precipitation) reaction — barium sulphate is a white insoluble precipitate.

(iii) Sodium + Water → Sodium hydroxide + Hydrogen

Unbalanced: Na + H₂O → NaOH + H₂

Step-by-step:

• Na: place 2 before Na and 2 before NaOH → Na = 2 on both sides.
• H: left = 2×2 = 4 (in 2H₂O), right = 2×1 + 2 = 4 (in 2NaOH + H₂). Balanced.
• O: left = 2 (in 2H₂O), right = 2 (in 2NaOH). Balanced.

2Na(s) + 2H₂O(l) → 2NaOH(aq) + H₂(g)

This is a displacement reaction — sodium displaces hydrogen from water. It is also highly exothermic; the hydrogen gas may catch fire.

(i) Solutions of barium chloride and sodium sulphate react to give insoluble barium sulphate and sodium chloride solution.

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

Verification of balance: Ba = 1, Cl = 2, Na = 2, S = 1, O = 4 on both sides.

Observations: A white precipitate of barium sulphate forms immediately when the two clear solutions are mixed. The precipitate is insoluble in water and settles at the bottom. This is a double displacement (precipitation) reaction.

(ii) Sodium hydroxide solution reacts with hydrochloric acid to produce sodium chloride solution and water.

NaOH(aq) + HCl(aq) → NaCl(aq) + H₂O(l)

Verification of balance: Na = 1, O = 1, H = 2, Cl = 1 on both sides.

This is a neutralisation reaction (a type of double displacement). It is exothermic — the temperature of the solution rises slightly. No visible precipitate forms; both products remain dissolved.

(i) Substance X is calcium oxide, commonly called quicklime. Its formula is CaO.

Quicklime is used for whitewashing because when it reacts with water, it forms calcium hydroxide (slaked lime), which is a cheap, white, and adhesive substance ideal for coating walls.

(ii) Reaction of calcium oxide with water:

CaO(s) + H₂O(l) → Ca(OH)₂(aq)

Calcium hydroxide is also called slaked lime. This reaction is highly exothermic — a lot of heat is released, and the solution may even boil or produce steam. The calcium hydroxide solution (limewater) is applied to walls; it slowly reacts with CO₂ in air to form a hard, white layer of calcium carbonate:

Ca(OH)₂(aq) + CO₂(g) → CaCO₃(s) + H₂O(l)

This is why whitewashed walls develop a shiny, hard finish over time.

During the electrolysis of water, the balanced chemical equation is:

2H₂O(l) → 2H₂(g) + O₂(g)

From the equation:

• 2 molecules of water produce 2 molecules of hydrogen and 1 molecule of oxygen.
• The ratio of hydrogen to oxygen by moles (and hence by volume at the same temperature and pressure) is 2 : 1.

Therefore, the volume of hydrogen gas collected at the cathode (negative electrode) is exactly double the volume of oxygen gas collected at the anode (positive electrode).

Experimental setup: In the laboratory, acidified water is electrolysed using two platinum or graphite electrodes. Hydrogen collects over the cathode and oxygen over the anode. The test tube over the cathode fills faster and has twice the gas volume compared to the test tube over the anode.

Iron is more reactive than copper (iron is above copper in the reactivity series). When an iron nail is placed in copper sulphate solution, iron displaces copper from the solution:

Fe(s) + CuSO₄(aq) → FeSO₄(aq) + Cu(s)

Observations explained:

• The blue colour of copper sulphate solution fades because Cu²⁺ ions (which give the blue colour) are being consumed as they are displaced from solution.
• The solution gradually turns light green due to the formation of Fe²⁺ ions (ferrous sulphate).
• A reddish-brown coating of copper metal deposits on the surface of the iron nail.
• Over time, the iron nail becomes thinner as it dissolves into the solution.

This is a displacement reaction. It occurs because iron has a higher tendency to lose electrons than copper.

Example:

AgNO₃(aq) + NaCl(aq) → AgCl(s)↓ + NaNO₃(aq)

In this reaction:

• Silver nitrate (AgNO₃) and sodium chloride (NaCl) are both soluble in water.
• When their solutions are mixed, the Ag⁺ ions and Cl⁻ ions combine to form silver chloride (AgCl), which is insoluble and appears as a white curdy precipitate.
• The Na⁺ ions and NO₃⁻ ions remain in solution as sodium nitrate.

The key feature of a double displacement reaction is that the cations and anions of two different compounds exchange partners. This particular example is also a precipitation reaction because an insoluble solid separates from the solution.

Another example:

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

Here, Ba²⁺ and SO₄²⁻ combine to form a white precipitate of barium sulphate.

CuO(s) + H₂(g) → Cu(s) + H₂O(g)

Analysis using oxygen transfer:

• CuO → Cu: Copper oxide loses oxygen to become copper. Loss of oxygen = reduction. So CuO is reduced.
• H₂ → H₂O: Hydrogen gains oxygen to form water. Gain of oxygen = oxidation. So H₂ is oxidised.

Identifying the agents:

• CuO is the oxidising agent (it provides oxygen to H₂).
• H₂ is the reducing agent (it removes oxygen from CuO).

This is a redox reaction because oxidation and reduction occur simultaneously.

2PbO(s) + C(s) → 2Pb(s) + CO₂(g)

(a) Lead is getting reduced. (b) Carbon dioxide is getting oxidised. (c) Carbon is getting oxidised. (d) Lead oxide is getting reduced.

(a) Combination (b) Double displacement (c) Decomposition (d) Displacement

(i) 2H₂(g) + O₂(g) → 2H₂O(l)

Type: Combination reaction (also exothermic).

Two elements (hydrogen and oxygen) combine to form a single product (water). Energy is released when the bonds in H₂O form, so this is also exothermic.

Exothermic reactions are chemical reactions that release energy (usually in the form of heat) to the surroundings. The energy of the products is less than the energy of the reactants.

Examples: 1. Burning of methane (combustion): CH₄(g) + 2O₂(g) → CO₂(g) + 2H₂O(l) + Heat 2. Respiration: C₆H₁₂O₆(aq) + 6O₂(g) → 6CO₂(g) + 6H₂O(l) + Energy (ATP) 3. Quicklime in water: CaO(s) + H₂O(l) → Ca(OH)₂(aq) + Heat 4. Rusting of iron: 4Fe(s) + 3O₂(g) + 2xH₂O(l) → 2Fe₂O₃·xH₂O(s) + Heat

Endothermic reactions are chemical reactions that absorb energy from the surroundings. The energy of the products is greater than the energy of the reactants.

Examples: 1. Thermal decomposition of calcium carbonate: CaCO₃(s) → CaO(s) + CO₂(g) (requires continuous heating) 2. Photosynthesis: 6CO₂(g) + 6H₂O(l) → C₆H₁₂O₆(aq) + 6O₂(g) (requires sunlight) 3. Electrolysis of water: 2H₂O(l) → 2H₂(g) + O₂(g) (requires electrical energy) 4. Decomposition of silver chloride: 2AgCl(s) → 2Ag(s) + Cl₂(g) (requires sunlight)

During respiration, glucose (from digested food) is slowly oxidised by oxygen in the cells of our body:

C₆H₁₂O₆(aq) + 6O₂(g) → 6CO₂(g) + 6H₂O(l) + Energy

This reaction releases a large amount of energy stored in the bonds of glucose molecules. The energy released is stored as ATP (adenosine triphosphate), which powers all life processes — muscle contraction, nerve impulses, cell division, and maintaining body temperature.

Since energy (heat) is released to the surroundings during this process, respiration is classified as an exothermic reaction.

Note: Respiration is a slow combustion — it happens at body temperature with the help of enzymes, unlike the rapid burning of fuel which happens at high temperatures.

In a combination reaction, two or more reactants combine to form a single product:

A + B → AB

Example: 2Mg(s) + O₂(g) → 2MgO(s)

In a decomposition reaction, a single reactant breaks down into two or more products:

AB → A + B

Example: 2MgO(s) → 2Mg(s) + O₂(g) (theoretical reverse)

The two types are exact opposites in direction: combination builds up a complex substance from simpler ones, while decomposition breaks down a complex substance into simpler ones. If you reverse a combination equation, you get a decomposition equation, and vice versa. That is why decomposition reactions are called the opposite of combination reactions.

(i) Decomposition by heat (Thermal decomposition):

2Pb(NO₃)₂(s) → 2PbO(s) + 4NO₂(g) + O₂(g)

Lead nitrate decomposes on heating. Observations: The white crystals turn yellow on heating and release brown fumes of nitrogen dioxide. A crackling sound (decrepitation) may also be heard.

Other example: CaCO₃(s) → CaO(s) + CO₂(g)

(ii) Decomposition by light (Photolytic decomposition / Photodecomposition):

2AgCl(s) → 2Ag(s) + Cl₂(g)

Silver chloride decomposes in sunlight. Observations: White silver chloride turns grey as metallic silver is formed. This is why silver chloride is stored in dark bottles — to prevent photodecomposition.

Other example: 2AgBr(s) → 2Ag(s) + Br₂(g) — this reaction is the basis of black-and-white photography.

(iii) Decomposition by electricity (Electrolytic decomposition / Electrolysis):

2H₂O(l) → 2H₂(g) + O₂(g)

Water decomposes when electric current is passed through acidified water. Observations: Bubbles of hydrogen form at the cathode and oxygen at the anode; the volume of hydrogen is twice that of oxygen.

Other example: 2NaCl(l) → 2Na(l) + Cl₂(g) — molten sodium chloride is electrolysed to obtain sodium metal.

Copper is more reactive than silver (copper is above silver in the reactivity series). When copper is placed in silver nitrate solution, it displaces silver:

Cu(s) + 2AgNO₃(aq) → Cu(NO₃)₂(aq) + 2Ag(s)

Type: Displacement reaction.

Observations:

• The colourless silver nitrate solution gradually turns blue due to the formation of copper(II) nitrate (Cu²⁺ ions are blue).
• Shiny, white deposits of silver form on the copper surface and may appear as needle-like crystals.
• The copper strip gradually dissolves (becomes thinner).

This method is used in the refining (purification) of silver. The impure silver is dissolved in nitric acid to form silver nitrate, and then a more reactive metal (copper) displaces pure silver from the solution.

A precipitation reaction is a type of double displacement reaction in which two soluble compounds in aqueous solution react to form an insoluble solid called a precipitate. The precipitate settles at the bottom of the container and can be separated by filtration.

Example:

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

When a solution of sodium sulphate is mixed with a solution of barium chloride, the Ba²⁺ ions and SO₄²⁻ ions combine to form barium sulphate (BaSO₄), which is a white, insoluble precipitate. The sodium ions and chloride ions remain dissolved in the solution as sodium chloride.

Observations:

• Both initial solutions are colourless and clear.
• On mixing, a white precipitate appears immediately.
• The precipitate is dense and settles at the bottom.
• The supernatant liquid remains clear (contains dissolved NaCl).

Other examples:

• Pb(NO₃)₂(aq) + 2KI(aq) → PbI₂(s)↓ + 2KNO₃(aq) — bright yellow precipitate of lead iodide
• AgNO₃(aq) + NaCl(aq) → AgCl(s)↓ + NaNO₃(aq) — white curdy precipitate of silver chloride

(a) Oxidation

Oxidation is the process in which a substance gains oxygen or loses hydrogen.

Examples of oxidation by gaining oxygen:

• 2Mg(s) + O₂(g) → 2MgO(s) — magnesium gains oxygen, so magnesium is oxidised.
• C(s) + O₂(g) → CO₂(g) — carbon gains oxygen, so carbon is oxidised.

Examples of oxidation by losing hydrogen:

• CH₃CH₂OH → CH₃CHO + H₂ — ethanol loses hydrogen, so ethanol is oxidised.
• ZnO + C → Zn + CO — in the older definition, consider H₂S + Cl₂ → S + 2HCl — H₂S loses hydrogen, so H₂S is oxidised.

(b) Reduction

Reduction is the process in which a substance loses oxygen or gains hydrogen.

Examples of reduction by losing oxygen:

• CuO(s) + H₂(g) → Cu(s) + H₂O(g) — CuO loses oxygen, so CuO is reduced.
• 2PbO(s) + C(s) → 2Pb(s) + CO₂(g) — PbO loses oxygen, so PbO is reduced.

Examples of reduction by gaining hydrogen:

• Cl₂(g) + H₂(g) → 2HCl(g) — chlorine gains hydrogen, so chlorine is reduced.

Remember: Oxidation and reduction always occur together in a redox reaction. The substance that is oxidised is the reducing agent, and the substance that is reduced is the oxidising agent.

CuO(s) + H₂(g) → Cu(s) + H₂O(g)

Step-by-step analysis:

1. CuO → Cu: Copper oxide has lost oxygen. Loss of oxygen = reduction. Therefore, CuO is reduced. 2. H₂ → H₂O: Hydrogen has gained oxygen. Gain of oxygen = oxidation. Therefore, H₂ is oxidised.

Identifying the agents:

• Since H₂ causes the reduction of CuO (by removing oxygen), H₂ is the reducing agent.
• Since CuO causes the oxidation of H₂ (by providing oxygen), CuO is the oxidising agent.

Summary table:

Iron reacts with oxygen and moisture present in air to form hydrated iron(III) oxide (rust):

4Fe(s) + 3O₂(g) + 2xH₂O(l) → 2Fe₂O₃·xH₂O(s)

Rust is a reddish-brown, porous, and flaky substance that gradually eats away the iron, making it weak and brittle. This process is called corrosion (specifically, rusting in the case of iron).

Painting iron articles creates a protective barrier between the iron surface and the surrounding air (oxygen + moisture). Without contact with oxygen and water, the rusting reaction cannot occur. This is why:

• Iron doors, windows, and grills are painted regularly.
• Ships and bridges are painted with special anti-corrosive paints.
• The paint must be reapplied periodically because it can chip, peel, or crack, exposing the iron underneath.

Other methods to prevent rusting:

• Galvanisation — coating iron with a layer of zinc.
• Alloying — making stainless steel by mixing iron with chromium and nickel.
• Greasing or oiling — applying a film of oil on moving parts.
• Plastic or ceramic coating — on pipes and fittings.