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(d) (iii) > (i)> (ii)

Acidity is directly proportional to oxidation number. As the oxidation number of S, P and Cl in H₂SO₃ , H₃PO₃ and HCIO₃ is +4, +3, +5 respectively. So decreasing order of acidity will be (iii) > (I) > (ii)

1. Arrhenius 

2. donate, accept 

3. Conjugate acid-base 

4. Lewis 

5. Lewis base 

6. Lewis base, Lewis acid 

7. Lewis acid, Lewis base 

8. Ka value

9. very weak acids 

10. very weak base

(b) Common ion effect

Larger the size of the halogen atom less is the back donation of electrons into empty 2p orbital of B

(a) 0.1 M H₂SO₄

[OH^-] = 1 pOH = 0

pH + pOH = 14 

pH = 14 – 0 = 14

(d) weak acid and strong base

NH₃ has the highest proton affinity.

1. In this case, neither the cations nor the anions undergo hydrolysis. Therefore the solution remains neutral.

2. For example, in the aqueous solution of NaCl, its ions Na⁺ and Cl^- ions have no tendency to react with H⁺ or OH^- ions of water. 

This is because the possible products of such reaction are NaOH and HCI which are completely dissociated. As a result, there is no change in the concentration of W and OH^- ions and hence the solution continues to remain neutral.

(a) Both A and R are correct and R is the correct explanation of A

The value of ionic product of water at 25°C is 1 x 10^-14 

1. Experimentally found that the concentration of H₃O in pure water is 1 x 10^-7 at 25°C.

2. Since the dissociation of water produces equal number of H₃O₄ and OH^- , the concentration of OH^- is also equal to 1 x 10^-7 at 25°C. The ionic product of water at 25°C is

K_w = [H₃O⁺] [OH^-]

= [1 x 10^-7] [1 x 10^-7]

K_w = [1 x 10^-14]

Ag₂ CrO₄ (S) ⇌ 2Ag⁺(aq) + CrO^2-₄ (aq)

K_sp =  [Ag⁺]² [CrO^2-₄]

= (5 x 10^-5)² (4.4 x 10^-4) 

= 1.1 x 10^-12

(c) 0.01 M CaCl₂

0.01 M CaCI₂ gives maximum CI^- ions to keep K_sp of AgCl constant, decrease in [Ag⁺] will be maximum.

The solubility equilibrium in the saturated solution is 

AgCl (s) ⇌ Ag⁺ (aq) + Cl^- (aq) 

The solubility of AgCl is 1.06 x 10^-5 mole per litre.

 [Ag⁺ (aq)] = 1.06 x 10^-5 mol L^-1

[Cl^- (aq)] = 1.06 x 10-⁵ mol L^-1

K_sp = [Ag⁺(aq)] [Cl^- (aq)] 

= (1.06 x 10^-5 mol L^-1) x (1.06 x 10^-5 mol L^-1) 

= 1.12 x 10^-2 moI^-2 L^-2

Ionic product 

1. It is applicable to all types of solutions. 

2. Its value changes with the change in concentration of the ions. 

Solubility product 

1. It is applicable to the saturated solutions. 

2. It has a definite value for an electrolyte at a constant temperature.

1. When the product of molar concentration of the constituent ions i.e., ionic product exceeds the solubility product then the compound gets precipitated. 

2. When the ionic Product > K_sp precipitation will occur and the solution is super saturated. ionic Product < K_sp no precipitation and the solution is unsaturated. ionic Product = K_sp equilibrium exist and the solution is saturated. 

3. By this way, the solubility product finds useful to decide whether an ionic compound gets precipitated when solution that contain the constituent ions are mixed.

(d) Almost neutral

It is a salt of weak acid and weak base.

(a) Both A and R are correct and R is the correct explanation of A

1. The buffering ability of a solution can be measured in terms of buffer capacity.

2. Buffer index, as a quantitative measure of the buffer capacity.

3. It is defined as the number of gram equivalents of acid or base added to 1 litre of the buffer solution to change its pH by unity.

4. β = \(\frac{dB}{d(pH)}\) .dB = number of gram equivalents of acid / base added to one litre of buffer solution. d(pH) = The change in the pH after the addition of acid / base.

1. Buffer is a solution which consists of a mixture of weak acid and its conjugate base (or) a weak base and its conjugate acid. 

2. This buffer solution resists drastic changes in its pH upon addition of a small quantities of acids (or) bases and this ability is called buffer action. 

3. Acidic buffer solution, Solution containing acetic acid and sodium acetate. Basic buffer solution, Solution containing NH₄O and NH₄Cl.

(b) 2

[H₃O⁺] = 0.01 M

pH = – 1og₁₀[H₃O⁺] = – log₁₀ [0.01] 

= – 1og₁₀ [10^-3] = 3

pH = 3

(c) pH x pOH = 1 x 10¹⁴

(b) strongly acidic 

pH = – log₁₀ [H⁺] 

[H⁺] =10^-pH

= 10⁰ = 1 

[H⁺] = 1 M 

The, solution is strongly acidic

(d) Calcium oxalate

1. HClO₄ , HCI, H₂SO₄ – are strong acids 

2. NH₂^- , O^2- , H^- – are strong bases 

3. HNO₂ , HF, CH₃COOH are weak acids

(c) CI^- 

CI^- is a conjugate base of strong acid HCl.

If the pH of an aqueous solution is 7, the solution is neutral

(d) 10^-6 M 

The concentration of Barium in the water.

Lewis Acids 

1. Lewis acids are substances that can accept one or more lone pair of electrons. 

2. All metal ions (or) atoms can act as Lewis acids. Examples: Fe²⁺ , Fe³⁺ , Cu²⁺ , Cr³⁺ 

3. Molecules that contain a polar double bond can act as Lewis acids. Examples: SO₂ , CO₂ , SO₃ 

4. Molecules in which the central atom can expand its act due to the availability of empty d-orbitais can act as Lewis acid. Example: SiF₄ , SF₄ , FeCI₃ 

5. Carbonium ion (CH₃)₃C⁺ can act as Lewis acid 

6. Electron deficient molecules such as BF₃ , AlCl₃ , BeF₂ act as Lewis acid (electron pair acceptors).

Lewis Bases 

1. Lewis bases are substances that can donate one or more lone pair of electrons. 

2. All anions can act as Lewis bases. Examples: F^- , Cl^- , CN^- , SO₄^2- 

3. Molecules that contain carbon-carbon multiple bond. Example: CH₂ = CH₂ , CH = CH 

4. All metal oxides can act as Lewis bases. Examples : CaO, MgO, Na₂O 

5. CH₂^- carbanion cari act as Lewis acid 

6. Electron rich molecules such as NH₃ , H₂O, ROH, R – O – R, R – NH₂ act as Lewis base (Electron pair donors)

BF₃  → electron deficient → Lewis acid

PF₃ → electron rich → Lewis base

CO → having lone pair of electron → Lewis base

F^- → unshared pair of electron → lewis base

Cl^- is the conjugate base of HCI

(c) H₂SO₄

Reason: H₂SO₄ is a strong acid whereas others are weak acids.

(b) 10^-8 M Ag⁺ and 10^-8 M I^- 

K_sp of AgI = 1.5 x 10^-16

10^-8 M Ag⁺ and 10^-8 M I^- 

Ionic product = 10^-16 = K_sp

HCI + H₂O ⇌ H₃O⁺ + Cl^- in this case, in addition to auto ionisation of water, HCI molecule also produces H₃O ion by donating a proton to water and hence [H₃O]> [OH]. It means that the aqueous HCI solution is acidic. Similarly in basic solution such as aqueous NH₃ , [OH^-] > [H₃O⁺] and it is basic.

(a) Both A and R are correct and R is the correct explanation of A

(a) OH^- and HSO₄

(d) NH₃

Reason: NH₃ is a Lewis base whereas others are Lewis acid.

(d) HCI

Reason: HCl is a strong acid whereas others are weak acids.

(d) H₂O, OH^-

(c) CH₃COOH 

Reason: CH₃COOH is a weak acid whereas others are strong acids.

BF₃ , H⁺ ions are Lewis acids.

(b) HCI + CI^- 

(d) HCl, Cl^-

(d) HCl + H₂SO₄

(c)  2 3 4 1

HClO₄ ⇌ H⁺ + ClO₄^-

1. According to Lowry – Bronsted concept, a strong acid has weak conjugate base and a weak acid has a strong conjugate base.

2. Let us consider the stabilities of the conjugate bases ClO₄^- , ClO₃^- , CIO₂^- and ClO^- formed from these acid HClO₄ , HClO₃ , HCIO₂ , HOCI respectively. 

These anions are stabilized to greater extent, it has lesser attraction for proton and therefore, will behave as weak base. Consequently the corresponding acid will be strongest because weak conjugate base has strong acid and strong conjugate base has weak acid. 

3. The charge stabilization mercases in the order, ClO^- < ClO₂^- < ClO₃^- < ClO₄^-. 

This means ClO₄^- will have maximum stability and therefore will have minimum attraction for W. Thus CIO₄^- will be weakest base and its conjugate acid HCIO₄ is the strongest acid. 

4. CIO₄^- is the conjugate base of the acid HClO₄

(a) Both A and R are correct and R is the correct explanation of A

(c) H₂CO₃ + NaHCO₃

(b) Sulphuric acid

(d) either (b) or (c)