Acids, Bases & Salts

Introduction: The pH Scale and Chemical Balance

Imagine a world where stomach acid could dissolve your stomach, or where soapy water burned your skin. In reality, chemistry maintains a delicate balance through acids and bases. Your blood maintains a pH of 7.35–7.45 (slightly basic) through buffer systems. Your stomach produces acid to digest food but protects itself with a mucous layer. Lakes and oceans maintain their pH through chemical equilibria. Understanding acids, bases, and salts isn't just chemistry—it's understanding how life itself functions.

In this chapter, you'll master the definitions, pH calculations, neutralization reactions, and the special behavior of salts. You'll learn why lemon juice is acidic, why baking soda is basic, and how salt solutions can be acidic, basic, or neutral. You'll understand indicators, buffers, and the chemistry of India's water, soil, and industrial processes.

Part 1: Definitions and Concepts

Arrhenius Definition

Acid: A substance that produces H⁺ ions (protons) in aqueous solution.

• Examples: HCl, H₂SO₄, HNO₃, CH₃COOH (acetic acid)

Base: A substance that produces OH⁻ ions in aqueous solution.

• Examples: NaOH, KOH, Ca(OH)₂, NH₃

Limitation: This definition only works in aqueous solutions. It doesn't explain basic behavior of ammonia (NH₃), which doesn't produce OH⁻ directly but accepts protons.

Brønsted-Lowry Definition (Modern, Broader)

Acid: A proton (H⁺) donor

• HCl → H⁺ + Cl⁻ (HCl donates H⁺)
• CH₃COOH → H⁺ + CH₃COO⁻ (acetic acid donates H⁺)

Base: A proton (H⁺) acceptor

• OH⁻ + H⁺ → H₂O (OH⁻ accepts H⁺)
• NH₃ + H⁺ → NH₄⁺ (ammonia accepts H⁺)

Advantage: Works in both aqueous and non-aqueous solutions. More comprehensive.

[Memory Hook] Brønsted-Lowry: Acid = H⁺ donor, Base = H⁺ acceptor

Conjugate Acid-Base Pairs

When an acid donates a proton, it becomes a conjugate base. When a base accepts a proton, it becomes a conjugate acid.

Example: Acetic acid and acetate ion CH₃COOH ⇌ H⁺ + CH₃COO⁻

• CH₃COOH is an acid; CH₃COO⁻ is its conjugate base
• CH₃COO⁻ is a base; CH₃COOH is its conjugate acid

The stronger the acid, the weaker its conjugate base (and vice versa).

Part 2: The pH Scale

Water Ionization

In pure water, molecules occasionally ionize: H₂O ⇌ H⁺ + OH⁻

At 25°C, [H⁺] = [OH⁻] = 10⁻⁷ M (molar)

Ion product of water (Kw): Kw = [H⁺][OH⁻] = 10⁻¹⁴ (at 25°C)

pH Definition

pH = -log₁₀[H⁺]

This measures the concentration of H⁺ ions on a logarithmic scale.

• pH = 7: Neutral ([H⁺] = [OH⁻] = 10⁻⁷)
• pH < 7: Acidic ([H⁺] > 10⁻⁷)
• pH > 7: Basic ([H⁺] < 10⁻⁷)

pH Scale (0–14)

pH	Acidity	[H⁺] (approx)	Example
0–1	Very strong acid	10⁻⁰ to 10⁻¹	Battery acid
2–3	Strong acid	10⁻² to 10⁻³	Lemon juice (pH 2), Vinegar (pH 3)
4–5	Weak acid	10⁻⁴ to 10⁻⁵	Tomato juice (pH 4), Black coffee (pH 5)
6	Weakly acidic	10⁻⁶	Milk (pH 6.5)
7	Neutral	10⁻⁷	Pure water
8–9	Weakly basic	10⁻⁸ to 10⁻⁹	Sea water (pH 8), Baking soda solution (pH 8.3)
10–11	Strong base	10⁻¹⁰ to 10⁻¹¹	Milk of magnesia (pH 10), Ammonia solution (pH 11)
12–14	Very strong base	10⁻¹² to 10⁻¹⁴	Soda lye (pH 13), NaOH solution (pH 14)

[Memory Hook] pH 7 = neutral; each pH unit = 10-fold change in H⁺ concentration

[Exam Trap] pH is logarithmic, not linear. A pH change from 5 to 3 means [H⁺] increases by 100-fold, not 2-fold!

pOH and Relationship to pH

pOH = -log₁₀[OH⁻]

Relationship: pH + pOH = 14 (at 25°C)

Example: If pH = 3, then pOH = 14 - 3 = 11

Part 3: Strong vs. Weak Acids and Bases

Strong Acids (Completely Ionize)

Definition: Acids that dissociate completely in water.

The Big 6 strong acids (must memorize):

1. HCl (Hydrochloric acid)
2. HBr (Hydrobromic acid)
3. HI (Hydroiodic acid)
4. HNO₃ (Nitric acid)
5. H₂SO₄ (Sulfuric acid)
6. HClO₄ (Perchloric acid)

Ionization: HCl → H⁺ + Cl⁻ (100% dissociation)

If you have 0.1 M HCl, [H⁺] = 0.1 M (complete ionization)

Strong Bases (Completely Ionize)

Definition: Bases that dissociate completely in water.

Common strong bases:

• Group 1 hydroxides: LiOH, NaOH, KOH (soluble)
• Group 2 hydroxides: Ca(OH)₂, Ba(OH)₂ (some soluble)

Ionization: NaOH → Na⁺ + OH⁻ (100% dissociation)

If you have 0.1 M NaOH, [OH⁻] = 0.1 M

Weak Acids (Partially Ionize)

Definition: Acids that ionize incompletely; most molecules remain un-ionized.

Common weak acids:

• CH₃COOH (Acetic acid, vinegar)
• HF (Hydrofluoric acid)
• H₂CO₃ (Carbonic acid)
• H₃PO₄ (Phosphoric acid)

Ionization: CH₃COOH ⇌ H⁺ + CH₃COO⁻

Only ~1% ionizes; 99% remains as CH₃COOH molecules

Ka (acid dissociation constant): Quantifies strength of weak acid

• Large Ka = stronger acid (more ionization)
• Small Ka = weaker acid (less ionization)

Weak Bases (Partially Ionize)

Definition: Bases that ionize incompletely.

Common weak bases:

• NH₃ (Ammonia): NH₃ + H₂O ⇌ NH₄⁺ + OH⁻
• Amines: RNH₂
• Some metal hydroxides: Al(OH)₃, Zn(OH)₂

[Memory Hook] Strong acids/bases: complete ionization; weak acids/bases: partial ionization (equilibrium)

[Exam Trap] Students confuse "concentration" with "strength." A weak acid can be concentrated (high molarity) or dilute (low molarity). Its weakness refers to incomplete ionization, not concentration.

Part 4: Indicators

Indicators are organic compounds that change color based on pH. They work because their molecular structure changes when H⁺ ions are added or removed.

Common Indicators

Indicator	Color in Acid	Color in Base	pH Range
Litmus	Red	Blue	5–8
Methyl Orange	Red	Yellow	3–4
Methyl Red	Red	Yellow	5–6
Phenolphthalein	Colorless	Pink/Magenta	8–10
Methylene Blue	Colorless	Blue	10–11

Applications:

• Acid-base titrations: Use an indicator to detect endpoint (where color just changes)
• pH testing strips: Contain multiple indicators for broad pH ranges
• Qualitative tests: Distinguish acidic from basic solutions quickly

Example: To test if a solution is acidic or basic:

• Add litmus paper: turns red if acidic, blue if basic
• Add phenolphthalein: turns pink if basic, stays colorless if acidic

[Memory Hook] Phenolphthalein: colorless in acid, pink in base (most common in titrations)

Part 5: Neutralization and Salt Formation

Neutralization Reaction

Definition: Reaction between an acid and a base producing a salt and water.

General form: Acid + Base → Salt + Water

H⁺ + OH⁻ → H₂O (ionic equation)

Example: HCl + NaOH → NaCl + H₂O

• HCl (acid): provides H⁺
• NaOH (base): provides OH⁻
• Combine: H⁺ + OH⁻ → H₂O
• Remaining ions (Na⁺, Cl⁻) form salt NaCl

Stoichiometry of neutralization: For a monoprotic acid and monobasic base: moles of acid = moles of base at equivalence point

Salts and Hydrolysis

A salt is the ionic compound formed when the H⁺ from an acid combines with the OH⁻ from a base.

Type 1: Salt of Strong Acid + Strong Base (Neutral)

Example: NaCl (from HCl + NaOH)

• Neither ion hydrolyzes
• Solution remains neutral (pH ≈ 7)
• Other examples: KNO₃, Na₂SO₄, Ca(NO₃)₂

Type 2: Salt of Weak Acid + Strong Base (Basic)

Example: NaCH₃COO (sodium acetate, from CH₃COOH + NaOH)

• CH₃COO⁻ ion hydrolyzes: CH₃COO⁻ + H₂O ⇌ CH₃COOH + OH⁻
• Produces OH⁻ ions; solution becomes basic (pH > 7)
• Other examples: Na₂CO₃ (carbonate), NaHCO₃ (bicarbonate)

Type 3: Salt of Strong Acid + Weak Base (Acidic)

Example: NH₄Cl (ammonium chloride, from HCl + NH₃)

• NH₄⁺ ion hydrolyzes: NH₄⁺ + H₂O ⇌ NH₃ + H₃O⁺
• Produces H₃O⁺ (acidic) ions; solution becomes acidic (pH < 7)
• Other examples: FeCl₃, AlCl₃

Type 4: Salt of Weak Acid + Weak Base

Example: NH₄CH₃COO (ammonium acetate)

• Both ions hydrolyze; pH depends on which hydrolysis is stronger
• Usually nearly neutral if both weak acid and weak base have similar strengths

[Memory Hook] Salt hydrolysis: weak acid's conjugate base is basic; weak base's conjugate acid is acidic

Part 6: Buffer Solutions

What is a Buffer?

A buffer is a solution that resists changes in pH when small amounts of acid or base are added.

Composition: Weak acid + its conjugate base salt, OR weak base + its conjugate acid salt

Examples:

• Acetate buffer: CH₃COOH + NaCH₃COO
• Phosphate buffer: H₂PO₄⁻ + HPO₄²⁻
• Ammonia buffer: NH₃ + NH₄Cl

How Buffers Work

Adding H⁺ (acid): The buffer's conjugate base absorbs H⁺: CH₃COO⁻ + H⁺ → CH₃COOH

Adding OH⁻ (base): The buffer's weak acid donates H⁺: CH₃COOH + OH⁻ → CH₃COO⁻ + H₂O

Henderson-Hasselbalch Equation

Formula: pH = pKa + log([A⁻] / [HA])

Where:

• pKa = -log(Ka)
• [A⁻] = concentration of conjugate base
• [HA] = concentration of weak acid

Application: When [A⁻] = [HA], the pH = pKa (the buffer's optimal pH for resistance)

[Memory Hook] Buffer = weak acid + conjugate base (or weak base + conjugate acid)

Biological Buffers

Blood buffer system: H₂CO₃ / HCO₃⁻

• Maintains pH at 7.35–7.45
• Carbonic acid (H₂CO₃) donates H⁺ if pH rises
• Bicarbonate (HCO₃⁻) accepts H⁺ if pH falls

Why it matters: Small changes in blood pH can be fatal (< 6.8 or > 7.8 = death)

Part 7: India-Specific Applications

Water Hardness and Treatment

Hardness is caused by dissolved Ca²⁺ and Mg²⁺ ions.

Sources:

• Limestone regions (CaCO₃): Rajasthan, Andhra Pradesh
• Groundwater in mineral-rich areas

Problems:

• Soap won't lather (forms insoluble Ca²⁺ and Mg²⁺ soaps)
• Kettles accumulate mineral deposits
• Pipes corrode

Treatment:

• Boiling: Removes temporary hardness (CaCO₃ precipitates) CaCO₃(aq) + heat → CaCO₃(s) + CO₂↑ + H₂O

• Sodium carbonate (washing soda): Removes permanent hardness Ca²⁺ + CO₃²⁻ → CaCO₃↓

• Ion exchange resins: Replace Ca²⁺/Mg²⁺ with Na⁺

Soil Chemistry (Acid Soils in India)

Problem: Acidic soils (pH < 6.5) in:

• Western Ghats (tea plantations)
• Nilgiris (coffee plantations)
• Some regions of Jharkhand

Effects:

• Reduced crop yield
• Nutrient absorption problems
• Aluminum toxicity

Remedy:

• Add lime (CaO or Ca(OH)₂): neutralizes acidity
• Add limestone (CaCO₃): slower but effective
• CaCO₃ + 2HCl(soil) → CaCl₂ + H₂O + CO₂

Indian Antacids

Antacids neutralize excess stomach acid (HCl).

Common Indian antacids:

• Calcium carbonate (CaCO₃): Quick relief CaCO₃ + 2HCl → CaCl₂ + H₂O + CO₂↑ (CO₂ causes belching)

• Magnesium hydroxide (Mg(OH)₂): Milk of magnesia Mg(OH)₂ + 2HCl → MgCl₂ + 2H₂O

• Sodium bicarbonate (NaHCO₃): Baking soda NaHCO₃ + HCl → NaCl + H₂O + CO₂↑ (Fast but short-lasting)

Conclusion

Acids and bases are fundamental to chemistry and life. Understanding pH, neutralization, buffers, and salt hydrolysis explains phenomena from industrial chemistry to your body's internal processes. The pH scale connects the atomic world (H⁺ concentration) to observable reality (color changes, physiological effects).

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23 MCQ Questions

Q1: According to Arrhenius definition, an acid is a substance that produces:

• A) OH⁻ ions in aqueous solution
• B) H⁺ ions in aqueous solution
• C) Neutralized ions
• D) Negative ions

Q2: The Brønsted-Lowry definition describes an acid as:

• A) A source of H⁺ ions
• B) A proton donor
• C) A substance that increases H⁺ concentration
• D) A substance that tastes sour

Q3: In the reaction CH₃COOH ⇌ H⁺ + CH₃COO⁻, the conjugate base is:

• A) CH₃COOH
• B) H⁺
• C) CH₃COO⁻
• D) All of the above

Q4: At 25°C, if the pH of a solution is 3, what is the pOH?

• A) 3
• B) 7
• C) 11
• D) 14

Q5: A solution with pH = 5 has [H⁺] concentration of:

• A) 10⁻⁵ M
• B) 10⁻⁹ M
• C) 5 M
• D) 0.5 M

Q6: Which of the following is NOT a strong acid?

• A) HCl (Hydrochloric acid)
• B) H₂SO₄ (Sulfuric acid)
• C) CH₃COOH (Acetic acid)
• D) HNO₃ (Nitric acid)

Q7: Weak acids are characterized by:

• A) Complete ionization in water
• B) Partial ionization; equilibrium between ionized and un-ionized forms
• C) Not producing H⁺ ions
• D) Having very low concentration

Q8: If litmus paper turns blue when added to a solution, the solution is:

• A) Acidic
• B) Neutral
• C) Basic
• D) Cannot be determined

Q9: Phenolphthalein is colorless in:

• A) Acidic solution
• B) Neutral solution
• C) Basic solution
• D) All pH ranges

Q10: In a neutralization reaction, HCl + NaOH → NaCl + H₂O, the salt formed is:

• A) Acidic
• B) Basic
• C) Neutral
• D) Depends on concentration

Q11: Sodium acetate (NaCH₃COO) dissolved in water forms a solution that is:

• A) Acidic (due to acetate hydrolysis)
• B) Basic (due to acetate hydrolysis)
• C) Neutral
• D) Depends on temperature

Q12: Ammonium chloride (NH₄Cl) dissolved in water forms a solution that is:

• A) Acidic
• B) Basic
• C) Neutral
• D) Cannot be determined

Q13: A buffer solution consists of:

• A) A strong acid and strong base
• B) A weak acid and its conjugate base (or weak base and its conjugate acid)
• C) Only distilled water
• D) Only salt solution

Q14: In the Henderson-Hasselbalch equation, when [HA] = [A⁻], the pH equals:

• A) 7
• B) pKa
• C) pKb
• D) 14

Q15: The buffer capacity of a buffer solution is greatest when:

• A) [HA] << [A⁻]
• B) [HA] >> [A⁻]
• C) [HA] = [A⁻]
• D) [HA] + [A⁻] is very small

Q16: Water hardness is primarily caused by:

• A) Dissolved oxygen
• B) Dissolved Ca²⁺ and Mg²⁺ ions
• C) Dissolved salts from rock salt
• D) Presence of iron

Q17: To remove temporary hardness from water, the most effective method is:

• A) Adding soap
• B) Boiling the water (CaCO₃ precipitates)
• C) Adding soda ash (Na₂CO₃)
• D) Using ion-exchange resins

Q18: Acidic soils (pH < 6.5) can be neutralized by adding:

• A) Acid
• B) Lime (CaO) or limestone (CaCO₃)
• C) Salt
• D) Fertilizer

Q19: Milk of magnesia (Mg(OH)₂) is used as an antacid because:

• A) It is a strong base and neutralizes stomach acid
• B) It reacts with HCl: Mg(OH)₂ + 2HCl → MgCl₂ + 2H₂O
• C) It increases pH in the stomach
• D) Both B and C

Q20: Which indicator best shows the transition from acidic to basic at pH 8–10?

• A) Methyl orange
• B) Methyl red
• C) Phenolphthalein
• D) Litmus

Q21: In the Kw expression at 25°C, if [H⁺] = 10⁻⁴ M, what is [OH⁻]?

• A) 10⁻⁴ M
• B) 10⁻¹⁰ M
• C) 10⁻¹⁸ M
• D) 4 × 10⁻¹¹ M

Q22: Blood pH is maintained at 7.35–7.45 primarily by the:

• A) Hemoglobin-oxyhemoglobin buffer
• B) H₂CO₃ / HCO₃⁻ buffer system
• C) Phosphate buffer system
• D) Protein buffers only

Q23: A solution with very high buffering capacity against added base should contain:

• A) Excess weak acid and very little conjugate base
• B) Excess conjugate base and very little weak acid
• C) Equal amounts of weak acid and conjugate base
• D) Mostly salt and very little weak acid

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Answer Key: 1-B, 2-B, 3-C, 4-C, 5-A, 6-C, 7-B, 8-C, 9-A, 10-C, 11-B, 12-A, 13-B, 14-B, 15-C, 16-B, 17-B, 18-B, 19-D, 20-C, 21-B, 22-B, 23-A