Introduction to Acids & Bases

Acids and bases are fundamental chemical substances that play crucial roles in chemistry and everyday life. Acids taste sour, while bases taste bitter and feel slippery.

Common Organic Acids

Citric Acid

Found in citrus fruits like lemons, oranges

Lactic Acid

Found in sour milk, yogurt

Formic Acid

Found in ant bites, bee stings

Essential Water Role:

HCl(g) + H₂O(l) → H₃O⁺(aq) + Cl⁻(aq)

Water is essential for acid ionization – HCl only forms H⁺ ions in water!

H⁺
OH⁻
H₂O
H₂O
H₂O

Memory Trick: “ACID vs BASE”

Acids: Sour taste, Corrosive, Indicators turn Red, Donate H⁺

Bases: Bitter taste, Slippery feel, Indicators turn Blue, Accept H⁺ (or donate OH⁻)

Acid-Base Concepts

Conjugate Acid-Base Pairs

When an acid donates a proton (H⁺), it becomes its conjugate base. When a base accepts a proton, it becomes its conjugate acid.

Key Relationship:

Acid₁ + Base₂ ⇌ Base₁ + Acid₂

Where Acid₁/Base₁ and Base₂/Acid₂ are conjugate pairs

Example: HCl in Water

HCl(aq) + H₂O(l) ⇌ H₃O⁺(aq) + Cl⁻(aq)

Acid: HCl | Base: H₂O

Conjugate Acid: H₃O⁺ | Conjugate Base: Cl⁻

Lewis Acid-Base Concept

Lewis theory focuses on electron pairs rather than protons:

Lewis Acid: Electron pair acceptor (e.g., BF₃, AlCl₃)

Lewis Base: Electron pair donor (e.g., NH₃, H₂O, OH⁻)

Lewis Acid-Base Reaction:

BF₃ + :NH₃ → F₃B:NH₃

BF₃ (acid) accepts electron pair from NH₃ (base)

pH, pOH & Calculations

Ionization of Water

Water auto-ionizes: H₂O + H₂O ⇌ H₃O⁺ + OH⁻

At 25°C: K_w = [H⁺][OH⁻] = 1.0 × 10⁻¹⁴

0 (Strong Acid)
7 (Neutral)
14 (Strong Base)

pH Formulas

pH = -log₁₀[H⁺] (where [H⁺] is in mol/dm³)

pOH = -log₁₀[OH⁻]

pH + pOH = 14 (at 25°C)

Example 9.2: Household Ammonia

Given: [OH⁻] = 0.005 M = 5 × 10⁻³ M

Step 1: pOH = -log(5 × 10⁻³) = 2.30

Step 2: pH = 14 – pOH = 14 – 2.30 = 11.70

Step 3: [H⁺] = 10⁻¹¹·⁷⁰ = 2.0 × 10⁻¹² M

Applications

Buffer Solutions

Buffer solutions resist pH changes when small amounts of acid or base are added.

Weak Acid
Its Salt
Conjugate Base

Types of Buffers

Acidic Buffer: Weak acid + its salt with strong base (pH < 7)

Example: CH₃COOH + CH₃COONa

Basic Buffer: Weak base + its salt with strong acid (pH > 7)

Example: NH₄OH + NH₄Cl

Henderson-Hasselbalch Equation

For acidic buffers: pH = pKₐ + log([salt]/[acid])

For basic buffers: pOH = pK_b + log([salt]/[base])

Salt Hydrolysis

Salts react with water to produce acidic or basic solutions.

NH₄Cl (Acidic Salt):

NH₄⁺ + H₂O ⇌ NH₄OH + H⁺

Turns blue litmus red

Na₂CO₃ (Basic Salt):

CO₃²⁻ + H₂O ⇌ HCO₃⁻ + OH⁻

Turns red litmus blue

Salt Hydrolysis Summary

Salt Type Example pH Hydrolysis
Strong acid + Strong base NaCl, KNO₃ = 7 (Neutral) No hydrolysis
Strong acid + Weak base NH₄Cl, CuSO₄ < 7 (Acidic) Cation hydrolyzes
Weak acid + Strong base CH₃COONa, NaCN > 7 (Basic) Anion hydrolyzes

MCQ Exercises

1. Water cannot act as:

A Lewis acid
B Lewis base
C Conjugate acid
D Conjugate base

Correct Answer: C – Water CAN act as conjugate acid

Actually, water CAN act as all: Lewis acid (accepts e⁻ pair), Lewis base (donates e⁻ pair), conjugate acid (H₃O⁺/H₂O), conjugate base (H₂O/OH⁻).

Correction: The question seems flawed. Water can act as conjugate acid (of OH⁻) or conjugate base (of H₃O⁺).

2. pH of 0.01 M HCl solution is:

A 10⁻²
B 10⁺²
C 2.0
D 1.0

Correct Answer: C – 2.0

[H⁺] = 0.01 M = 10⁻² M

pH = -log₁₀(10⁻²) = -(-2) = 2.0

For strong monoprotic acids: pH = -log[Molarity]

3. An aqueous solution of ammonium chloride is:

A Basic
B Acidic
C Neutral
D Amphoteric

Correct Answer: B – Acidic

NH₄Cl is salt of strong acid (HCl) + weak base (NH₄OH)

NH₄⁺ hydrolyzes: NH₄⁺ + H₂O ⇌ NH₄OH + H⁺

Produces H⁺ ions → acidic solution (pH < 7)

4. Which compound gives basic aqueous solution?

A Ammonium nitrate
B Calcium chloride
C Ammonium acetate
D Potassium carbonate

Correct Answer: D – Potassium carbonate

K₂CO₃ is salt of weak acid (H₂CO₃) + strong base (KOH)

CO₃²⁻ hydrolyzes: CO₃²⁻ + H₂O ⇌ HCO₃⁻ + OH⁻

Produces OH⁻ ions → basic solution (pH > 7)

5. Which statement about acids is NOT correct?

A Always contains oxygen
B Contains H⁺ ions in solution
C pH < 7
D Reacts with carbonates to give CO₂

Correct Answer: A – “Always contains oxygen”

This is incorrect! Many acids don’t contain oxygen:

– HCl (hydrochloric acid)

– HBr (hydrobromic acid)

– HI (hydroiodic acid)

– HCN (hydrocyanic acid)

These are “hydro” acids without oxygen.

6. If liquid has pH = 7:

A Must be colorless
B Boiling point = 100°C
C Must be a solution
D Must be neutral

Correct Answer: D – Must be neutral

By definition: pH = 7 means [H⁺] = [OH⁻] = 10⁻⁷ M at 25°C

– Can be colored (e.g., neutral red indicator)

– Can have any boiling point

– Pure water also has pH = 7 (not just solutions)

But MUST be neutral (equal H⁺ and OH⁻ concentrations)

7. Air bubbled through water lowers pH from 7 to 5.6. Which gas causes this?

A Argon
B Carbon dioxide
C Nitrogen
D Oxygen

Correct Answer: B – Carbon dioxide

CO₂ dissolves in water forming carbonic acid:

CO₂ + H₂O ⇌ H₂CO₃ ⇌ H⁺ + HCO₃⁻

This increases [H⁺], lowering pH.

Rainwater pH ~5.6 due to dissolved CO₂ (normal “acid rain” without pollution)

8. 25cm³ of 1M HNO₃ + 50cm³ of 0.5M KOH. Resulting pH?

A 5
B 7
C 9
D 14

Correct Answer: B – 7 (Neutral)

Moles HNO₃ = 0.025 L × 1 M = 0.025 mol

Moles KOH = 0.050 L × 0.5 M = 0.025 mol

Equal moles of strong acid + strong base → complete neutralization

Result: Salt (KNO₃) + water → pH = 7

9. Dry citric acid on dry litmus paper:

A Turns yellow
B Turns green
C Turns red
D No change

Correct Answer: D – No change

Litmus paper works in AQUEOUS medium only!

Dry acid cannot ionize to produce H⁺ ions

Dry litmus cannot detect acidity (needs moisture)

Acids only show acidic properties in solution

10. Which is NOT a base?

A Aqueous ammonia
B Copper oxide
C Potassium chloride
D Sodium carbonate

Correct Answer: C – Potassium chloride

KCl is neutral salt (strong acid + strong base)

Other options are bases:

– NH₄OH (aqueous ammonia): weak base

– CuO: metal oxide (basic oxide)

– Na₂CO₃: basic salt (hydrolyzes to give OH⁻)

11. A strong acid:

A Always partially ionized
B Always fully ionized in solution
C Always decomposes carbonates
D Always contains oxygen

Correct Answer: B – Always fully ionized in solution

Definition of strong acid: 100% dissociation in water

Examples: HCl, HNO₃, H₂SO₄ (first proton), HClO₄

– Weak acids are partially ionized

– All acids (strong/weak) decompose carbonates

– Many strong acids don’t contain oxygen (HCl, HBr)

12. Which oxide gives acidic solution in water?

A MgO
B Na₂O
C SO₂
D SiO₂

Correct Answer: C – SO₂

SO₂ is acidic oxide (non-metal oxide)

SO₂ + H₂O → H₂SO₃ (sulfurous acid)

Other oxides:

– MgO, Na₂O: Basic oxides (form bases)

– SiO₂: Amphoteric or neutral (insoluble)

13. Heating blue CuSO₄·5H₂O turns white because:

A Loss of water only
B Loss of water + SO₂
C Reaction with CO₂
D Loss of water, SO₂ and O₂

Correct Answer: A – Loss of water only

Copper sulfate pentahydrate: CuSO₄·5H₂O (blue)

Heating removes water of crystallization:

CuSO₄·5H₂O → CuSO₄ (white) + 5H₂O

The blue color comes from coordinated water molecules. When removed, color changes to white.

Short Questions

2(i) What are conjugate acid-base pairs?

Conjugate acid-base pairs: Two species that differ by a single proton (H⁺).

Definition: When an acid donates H⁺, it becomes its conjugate base. When a base accepts H⁺, it becomes its conjugate acid.

Examples:

  • HCl/Cl⁻ (acid/conjugate base)
  • NH₃/NH₄⁺ (base/conjugate acid)
  • H₂O/OH⁻ (acid/conjugate base)
  • H₃O⁺/H₂O (acid/conjugate base)

Key relationship: Stronger acid → Weaker conjugate base, Weaker acid → Stronger conjugate base

2(ii) Define Lewis acid and base with examples

Lewis Acid: Electron pair acceptor

Characteristics: Electron-deficient, can accept e⁻ pair

Examples: BF₃, AlCl₃, H⁺, Ag⁺, SO₃

Lewis Base: Electron pair donor

Characteristics: Has lone pair, can donate e⁻ pair

Examples: NH₃, H₂O, OH⁻, :NH₃, R-OH

Lewis Acid-Base Reaction: Forms coordinate covalent bond

Example: BF₃ + :NH₃ → F₃B:NH₃

BF₃ (acid) accepts e⁻ pair from NH₃ (base)

2(iii) Ionization of water

Auto-ionization of Water:

H₂O(l) + H₂O(l) ⇌ H₃O⁺(aq) + OH⁻(aq)

Or simply: H₂O(l) ⇌ H⁺(aq) + OH⁻(aq)

Ion Product Constant (K_w):

K_w = [H⁺][OH⁻]

At 25°C: K_w = 1.0 × 10⁻¹⁴

Important Points:

  1. Pure water is neutral: [H⁺] = [OH⁻] = 1.0 × 10⁻⁷ M
  2. In acidic solution: [H⁺] > [OH⁻], [H⁺] > 10⁻⁷ M
  3. In basic solution: [H⁺] < [OH⁻], [H⁺] < 10⁻⁷ M
  4. K_w is temperature dependent (increases with temperature)

At 40°C: K_w = 3.8 × 10⁻¹⁴, so neutral [H⁺] = 1.95 × 10⁻⁷ M

2(iv) Define pH and values for different solutions

Definition: pH = -log₁₀[H⁺] (where [H⁺] in mol/dm³)

pH Scale (at 25°C):

  • Acidic solution: pH < 7, [H⁺] > 10⁻⁷ M
  • Neutral solution: pH = 7, [H⁺] = 10⁻⁷ M
  • Basic/Alkaline solution: pH > 7, [H⁺] < 10⁻⁷ M

pH Range: Typically 0-14, but can be outside this range for very concentrated solutions

Examples:

  • 0.1 M HCl: pH ≈ 1.0
  • Vinegar (5% acetic acid): pH ≈ 2.4
  • Coffee: pH ≈ 5.0
  • Pure water: pH = 7.0
  • Blood: pH = 7.35-7.45
  • Household ammonia: pH ≈ 11-12
  • 0.1 M NaOH: pH ≈ 13.0

Relationship: pH + pOH = 14 (at 25°C)

2(v) What are Kₐ and pKₐ?

Kₐ (Acid dissociation constant):

For acid HA: HA ⇌ H⁺ + A⁻

Kₐ = [H⁺][A⁻]/[HA]

Measures acid strength: Larger Kₐ = stronger acid

Examples: HCl Kₐ ≈ 10⁷ (very large), CH₃COOH Kₐ = 1.8 × 10⁻⁵

pKₐ: pKₐ = -log₁₀Kₐ

Inverse relationship: Smaller pKₐ = stronger acid

Examples: HCl pKₐ ≈ -7, CH₃COOH pKₐ = 4.76

Significance:

  1. pKₐ < 0: Very strong acid
  2. pKₐ 0-4: Strong acid
  3. pKₐ 4-9: Weak acid
  4. pKₐ > 9: Very weak acid

For buffer solutions: pH = pKₐ + log([salt]/[acid]) (Henderson-Hasselbalch)

2(vi) What are K_b and pK_b?

K_b (Base dissociation constant):

For base B: B + H₂O ⇌ BH⁺ + OH⁻

K_b = [BH⁺][OH⁻]/[B]

Measures base strength: Larger K_b = stronger base

Example: NH₃ K_b = 1.8 × 10⁻⁵

pK_b: pK_b = -log₁₀K_b

Inverse relationship: Smaller pK_b = stronger base

Example: NH₃ pK_b = 4.75

Relationship with Kₐ: For conjugate acid-base pair:

Kₐ × K_b = K_w = 1.0 × 10⁻¹⁴ (at 25°C)

pKₐ + pK_b = 14 (at 25°C)

Example: For NH₄⁺/NH₃ pair:

NH₄⁺: Kₐ = 5.6 × 10⁻¹⁰, pKₐ = 9.25

NH₃: K_b = 1.8 × 10⁻⁵, pK_b = 4.75

Check: 9.25 + 4.75 = 14 ✓

For basic buffers: pOH = pK_b + log([salt]/[base])

2(vii) Relationship between Kₐ and K_b

For conjugate acid-base pair:

Kₐ × K_b = K_w

Where: K_w = ion product of water = 1.0 × 10⁻¹⁴ (at 25°C)

Derivation:

For acid HA: HA ⇌ H⁺ + A⁻, Kₐ = [H⁺][A⁻]/[HA]

For conjugate base A⁻: A⁻ + H₂O ⇌ HA + OH⁻, K_b = [HA][OH⁻]/[A⁻]

Multiply: Kₐ × K_b = ([H⁺][A⁻]/[HA]) × ([HA][OH⁻]/[A⁻]) = [H⁺][OH⁻] = K_w

In logarithmic form:

pKₐ + pK_b = pK_w = 14 (at 25°C)

Examples:

  1. CH₃COOH/CH₃COO⁻ pair:

    Kₐ = 1.8 × 10⁻⁵, pKₐ = 4.74

    K_b = K_w/Kₐ = 10⁻¹⁴/(1.8×10⁻⁵) = 5.6 × 10⁻¹⁰, pK_b = 9.26

    4.74 + 9.26 = 14 ✓

  2. NH₄⁺/NH₃ pair:

    Kₐ(NH₄⁺) = 5.6 × 10⁻¹⁰, pKₐ = 9.25

    K_b(NH₃) = 1.8 × 10⁻⁵, pK_b = 4.75

    9.25 + 4.75 = 14 ✓

Significance: Stronger acid → Weaker conjugate base (and vice versa)

2(viii) Examples of buffer solutions

Buffer Solution: Resists pH change when small amounts of acid/base added

Types and Examples:

  1. Acidic Buffer (pH < 7):
    • Acetic acid-acetate buffer: CH₃COOH + CH₃COONa
    • pH range: 3.7-5.7 (pKₐ = 4.76)
    • Formic acid-formate buffer: HCOOH + HCOONa
    • pH range: 2.8-4.8 (pKₐ = 3.75)
  2. Basic Buffer (pH > 7):
    • Ammonia-ammonium buffer: NH₃ + NH₄Cl
    • pH range: 8.3-10.3 (pK_b = 4.75, so pKₐ = 9.25)
    • Carbonate-bicarbonate buffer: Na₂CO₃ + NaHCO₃
    • pH range: 9.3-11.3 (for HCO₃⁻/CO₃²⁻ pair)

Biological Buffers:

  • Blood buffer: H₂CO₃/HCO₃⁻ (maintains pH 7.35-7.45)
  • Phosphate buffer: H₂PO₄⁻/HPO₄²⁻ (in cells, pH ~6.8-7.2)

Preparation: Mix weak acid with its salt (acidic buffer) or weak base with its salt (basic buffer)