Acids, Bases and Salts Exercise Solutions | KalviGavani

Step-by-step procedure:

1. Dip the red litmus paper into each of the three test tubes and observe the colour change. 2. The test tube in which red litmus turns blue contains the basic solution (bases turn red litmus blue). 3. The other two test tubes do not change the colour of red litmus — one contains the acidic solution and the other contains distilled water (both are either acidic or neutral, so red litmus stays red). 4. Now take the piece of litmus that turned blue (from the basic solution test tube). Dip this blue litmus into each of the remaining two test tubes. 5. The test tube in which blue litmus turns red contains the acidic solution (acids turn blue litmus red). 6. The test tube in which blue litmus remains blue contains distilled water (neutral solution does not change litmus colour).

Key reasoning: We first identify the basic solution because only a base can turn red litmus blue. The resulting blue litmus then acts as a fresh indicator to distinguish acid from water.

Curd and other sour substances contain acids (lactic acid in curd, acetic acid in vinegar, citric acid in lemon juice). Brass is an alloy of copper and zinc, and copper vessels contain copper metal.

When acidic food is stored in such vessels, the acid reacts with the metal:

• With copper: Cu(s) + 2CH₃COOH(aq) → Cu(CH₃COO)₂(aq) + H₂(g)
• With zinc: Zn(s) + 2CH₃COOH(aq) → Zn(CH₃COO)₂(aq) + H₂(g)

These reactions produce toxic metal salts that dissolve into the food, making it unfit for consumption and potentially causing food poisoning. The food also develops a metallic taste and may get discoloured.

Hydrogen gas (H₂) is liberated when an acid reacts with a metal.

Example: Zinc reacts with dilute sulphuric acid:

Word equation: Zinc + Sulphuric acid → Zinc sulphate + Hydrogen

Balanced chemical equation: Zn(s) + H₂SO₄(aq) → ZnSO₄(aq) + H₂(g)

Test for hydrogen gas: Pass the gas through a soap solution to form soap bubbles. Bring a burning candle near a bubble. If the gas is hydrogen, it burns with a 'pop' sound. This is the standard test for hydrogen gas.

The gas that extinguishes a burning candle is carbon dioxide (CO₂). Since effervescence is produced and CO₂ is evolved, compound A must be a metal carbonate or metal hydrogen carbonate that contains calcium (because calcium chloride is one of the products).

Therefore, compound A is calcium carbonate (CaCO₃) — commonly found as limestone or marble.

Balanced chemical equation:

CaCO₃(s) + 2HCl(aq) → CaCl₂(aq) + H₂O(l) + CO₂(g)

Acids like HCl and HNO₃ ionise in aqueous solution to produce H⁺ (hydronium, H₃O⁺) ions, which are responsible for their acidic character:

• HCl(aq) → H⁺(aq) + Cl⁻(aq), or more accurately: HCl(aq) + H₂O(l) → H₃O⁺(aq) + Cl⁻(aq)
• HNO₃(aq) + H₂O(l) → H₃O⁺(aq) + NO₃⁻(aq)

Alcohols (e.g. C₂H₅OH) and glucose (C₆H₁₂O₆) contain hydrogen atoms, but they do not ionise in water to release H⁺ ions. Their hydrogen is bonded covalently and does not dissociate in aqueous solution. Since no free H⁺ ions are produced, these solutions do not conduct electricity and do not show acidic behaviour.

Activity to demonstrate: Set up an electric circuit with a bulb and two electrodes dipped in a glucose solution. The bulb does not glow, proving that glucose solution does not conduct electricity (no ions present). Now replace the glucose solution with HCl solution — the bulb glows brightly, confirming the presence of free ions.

An aqueous solution of an acid conducts electricity because acids ionise in water to produce free ions (H⁺/H₃O⁺ and the corresponding anion). These ions are mobile charge carriers that move towards the oppositely charged electrodes, thereby completing the electric circuit and allowing current to flow.

For example, HCl(aq) produces H⁺(aq) and Cl⁻(aq) ions which carry charge through the solution.

(a) 1 (b) 4 (c) 5 (d) 10

(a) NaCl (b) HCl (c) LiCl (d) KCl

(a) 4 mL (b) 8 mL (c) 16 mL (d) 12 mL

(a) Antibiotic (b) Analgesic (c) Antacid (d) Antiseptic

(a) Dilute sulphuric acid reacts with zinc granules.

Word equation: Zinc + Sulphuric acid → Zinc sulphate + Hydrogen

Balanced equation: Zn(s) + H₂SO₄(aq) → ZnSO₄(aq) + H₂(g)

(b) Dilute hydrochloric acid reacts with magnesium ribbon.

Word equation: Magnesium + Hydrochloric acid → Magnesium chloride + Hydrogen

Balanced equation: Mg(s) + 2HCl(aq) → MgCl₂(aq) + H₂(g)

(c) Dilute sulphuric acid reacts with aluminium powder.

Word equation: Aluminium + Sulphuric acid → Aluminium sulphate + Hydrogen

Balanced equation: 2Al(s) + 3H₂SO₄(aq) → Al₂(SO₄)₃(aq) + 3H₂(g)

(d) Dilute hydrochloric acid reacts with iron filings.

Word equation: Iron + Hydrochloric acid → Iron(II) chloride + Hydrogen

Balanced equation: Fe(s) + 2HCl(aq) → FeCl₂(aq) + H₂(g)

Note: Iron forms Fe²⁺ (ferrous/iron(II)) ions with dilute HCl, not Fe³⁺.

Activity:

1. Take a beaker containing a glucose solution (or alcohol). 2. Set up an electric circuit with a 6V battery, a bulb, connecting wires, and two graphite electrodes. 3. Dip the electrodes into the glucose solution and observe the bulb. 4. Observation: The bulb does not glow. 5. Now replace the glucose solution with dilute HCl solution and repeat. 6. Observation: The bulb glows brightly.

Conclusion: Glucose and alcohol do not ionise in aqueous solution to produce H⁺ ions. Although they contain hydrogen atoms, these are bonded covalently and do not dissociate as H⁺ ions in water. Without free ions, the solution cannot conduct electricity and hence does not show acidic behaviour. HCl, on the other hand, ionises completely to give H⁺ and Cl⁻ ions, which conduct electricity.

This proves that the presence of hydrogen alone does not make a compound an acid — the ability to release H⁺ ions in aqueous solution is essential.

Distilled water is pure water containing almost no dissolved ions. For electricity to be conducted through a solution, free ions must be present as charge carriers. Since distilled water lacks ions, it does not conduct electricity.

Rainwater, however, dissolves carbon dioxide (CO₂) from the atmosphere as it falls through the air. This dissolved CO₂ reacts with water to form carbonic acid:

CO₂(g) + H₂O(l) → H₂CO₃(aq)

Carbonic acid partially ionises in water:

H₂CO₃(aq) ⇌ H⁺(aq) + HCO₃⁻(aq)

These free H⁺ and HCO₃⁻ ions make rainwater a weak electrolyte, enabling it to conduct electricity. Additionally, rainwater may dissolve other atmospheric gases and dust particles that contribute ions.

Acids show acidic behaviour by releasing H⁺ (hydronium) ions in solution. This ionisation requires water as the medium:

HCl(g) + H₂O(l) → H₃O⁺(aq) + Cl⁻(aq)

In the absence of water, the acid molecules remain unionised (no dissociation occurs), so there are no free H⁺ ions available. Without H⁺ ions, the characteristic acidic properties — sour taste, turning blue litmus red, reacting with metals, conducting electricity — cannot be exhibited.

For example, dry HCl gas does not turn blue litmus red, but HCl dissolved in water (hydrochloric acid) does. This is because only in aqueous solution does HCl dissociate to produce H⁺ ions.

(a) neutral? Solution D (pH = 7). A neutral solution has equal concentrations of H⁺ and OH⁻ ions, giving pH 7.

(b) strongly alkaline? Solution C (pH = 11). pH much greater than 7 indicates a strongly basic solution with high OH⁻ ion concentration.

(c) strongly acidic? Solution B (pH = 1). pH close to 0 indicates a very strong acid with very high H⁺ ion concentration.

(d) weakly acidic? Solution A (pH = 4). pH slightly below 7 indicates a weak acid with moderate H⁺ ion concentration.

(e) weakly alkaline? Solution E (pH = 9). pH slightly above 7 indicates a weak base with moderate OH⁻ ion concentration.

Summary table:

When an acid solution is diluted by adding water, the total number of H₃O⁺ ions remains the same (no new acid molecules ionise and no existing ions are consumed), but these ions are now distributed throughout a larger volume of the solution.

Therefore, the concentration of H₃O⁺ ions decreases (concentration = number of ions / volume).

As a result:

• The solution becomes less acidic.
• The pH increases (moves closer to 7, but does not exceed 7 for an acid).

For example, diluting 1M HCl to 0.1M HCl reduces the H₃O⁺ concentration from 1 × 10⁻¹ mol/L to 1 × 10⁻² mol/L, and the pH rises from 1 to 2.

When excess base is dissolved in a NaOH solution, more OH⁻ ions are released into the solution (from the additional base dissolving and ionising). The volume increases only slightly compared to the increase in the number of OH⁻ ions.

Therefore, the concentration of OH⁻ ions increases significantly.

As a result:

• The solution becomes more basic (stronger alkali).
• The pH increases (moves further from 7 towards 14).
• The H⁺ ion concentration decreases (since pH + pOH = 14, higher OH⁻ means lower H⁺).

• Solution A (pH 6) has more hydrogen ion concentration because lower pH means higher H⁺ concentration. pH 6 is less than pH 8, so H⁺ concentration in A is greater than in B.
• Solution A is acidic (pH < 7 indicates acidic nature).
• Solution B is basic (pH > 7 indicates basic nature).

Quantitatively, the H⁺ ion concentration in solution A is 1 × 10⁻⁶ mol/L and in solution B is 1 × 10⁻⁸ mol/L. Solution A has 100 times more H⁺ ions than solution B (each pH unit difference = 10-fold change).

The concentration of H⁺(aq) ions directly determines the nature (acidic, neutral, or basic) of the solution:

• High H⁺ concentration (greater than 1 × 10⁻⁷ mol/L): The solution is acidic (pH < 7). The higher the H⁺ concentration, the more strongly acidic the solution.
• H⁺ concentration = 1 × 10⁻⁷ mol/L: The solution is neutral (pH = 7). H⁺ and OH⁻ concentrations are equal.
• Low H⁺ concentration (less than 1 × 10⁻⁷ mol/L): The solution is basic (pH > 7). The OH⁻ concentration exceeds the H⁺ concentration.

In summary, as H⁺(aq) concentration increases, the solution becomes more acidic (pH decreases). As H⁺(aq) concentration decreases, the solution becomes more basic (pH increases).

Yes, basic solutions do contain H⁺(aq) ions. This is because water undergoes auto-ionisation (self-ionisation):

H₂O(l) ⇌ H⁺(aq) + OH⁻(aq)

In pure water, the concentrations of H⁺ and OH⁻ are equal (both 1 × 10⁻⁷ mol/L at 25 °C), giving pH 7.

In a basic solution, the concentration of OH⁻ ions is greater than the concentration of H⁺ ions. It is this excess of OH⁻ over H⁺ that makes the solution basic. The H⁺ ions are still present (from water's auto-ionisation), but they are outnumbered by the OH⁻ ions.

For example, in a solution of pH 10, [H⁺] = 1 × 10⁻¹⁰ mol/L and [OH⁻] = 1 × 10⁻⁴ mol/L — OH⁻ ions far exceed H⁺ ions.

A farmer treats the soil with quick lime (CaO) or slaked lime (Ca(OH)₂) when the soil is too acidic (low pH).

When soil becomes excessively acidic (due to acidic rain, excessive use of chemical fertilisers, or decomposition of organic matter), plant growth is adversely affected — nutrients become less available and certain toxic elements (like aluminium) become more soluble.

Quick lime and slaked lime are bases. They neutralise the excess acid in the soil:

CaO(s) + 2H⁺(aq) → Ca²⁺(aq) + H₂O(l)

Ca(OH)₂(s) + 2H⁺(aq) → Ca²⁺(aq) + 2H₂O(l)

This raises the soil pH to a level suitable for healthy crop growth (typically pH 6–7).

The common name of CaOCl₂ is bleaching powder.

Its chemical name is calcium oxychloride (also called calcium hypochlorite chloride). It is used for bleaching cotton and linen in the textile industry, as a disinfectant for drinking water, and as an oxidising agent.

Calcium hydroxide (slaked lime), Ca(OH)₂, on treatment with chlorine yields bleaching powder.

Balanced chemical equation:

Ca(OH)₂(s) + Cl₂(g) → CaOCl₂(s) + H₂O(l)

The reaction is carried out by passing dry chlorine gas over dry slaked lime. The product is bleaching powder (CaOCl₂).

Washing soda — sodium carbonate decahydrate, Na₂CO₃·10H₂O — is used for softening hard water.

Hard water contains dissolved calcium and magnesium salts (Ca²⁺ and Mg²⁺ ions). Washing soda removes these ions by precipitating them as insoluble carbonates:

Na₂CO₃(aq) + CaSO₄(aq) → Na₂SO₄(aq) + CaCO₃(s)↓

Na₂CO₃(aq) + MgCl₂(aq) → 2NaCl(aq) + MgCO₃(s)↓

The precipitated carbonates are removed by filtration, leaving soft water.

When sodium hydrogen carbonate (baking soda, NaHCO₃) is heated, it decomposes to give sodium carbonate, water, and carbon dioxide gas:

Balanced equation: 2NaHCO₃(s) → Na₂CO₃(s) + H₂O(l) + CO₂(g)

Observations: Effervescence is observed due to the release of CO₂ gas. If the gas is passed through lime water (Ca(OH)₂ solution), the lime water turns milky, confirming the gas is CO₂:

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

The solid residue left behind is sodium carbonate (Na₂CO₃), which is alkaline and turns red litmus blue.

Plaster of Paris (CaSO₄·½H₂O) reacts with water to form gypsum (CaSO₄·2H₂O):

Balanced equation: CaSO₄·½H₂O(s) + 1½H₂O(l) → CaSO₄·2H₂O(s)