Acids, Bases and Salts

🟢 Lite — Quick Review (1h–1d)

Rapid summary of acids, bases, and salts for NABTEB chemistry.

Acids, Bases, and Salts are fundamental classes of compounds with distinct properties.

Arrhenius Definitions:

• Acid: A substance that produces $H^+$ ions in aqueous solution (e.g., $HCl \rightarrow H^+ + Cl^-$)
• Base: A substance that produces $OH^-$ ions in aqueous solution (e.g., $NaOH \rightarrow Na^+ + OH^-$)
• Salt: Produced when an acid reacts with a base: $\text{Acid} + \text{Base} \rightarrow \text{Salt} + \text{Water}$

Bronsted-Lowry Definitions (more general):

• Acid: A proton ($H^+$) donor
• Base: A proton ($H^+$) acceptor
• Conjugate acid-base pairs differ by one proton

Properties of Acids:

• Taste sour (e.g., lemon juice — citric acid)
• Turn litmus blue → red
• React with metals to produce hydrogen gas: $2HCl + Zn \rightarrow ZnCl_2 + H_2$
• React with carbonates to produce $CO_2$: $2HCl + CaCO_3 \rightarrow CaCl_2 + CO_2 + H_2O$
• React with bases (neutralisation)

Properties of Bases:

• Taste bitter
• Feel soapy/slippery
• Turn litmus red → blue
• React with acids (neutralisation)
• Some bases dissolve in water (alkalis): e.g., $NaOH$, $KOH$

pH Scale:

• pH 1–6: Acidic (lower pH = stronger acid)
• pH 7: Neutral (e.g., pure water)
• pH 8–14: Basic/alkaline (higher pH = stronger base)

⚡ NABTEB Exam Tip: The pH scale is logarithmic — pH 4 is 10 times more acidic than pH 5, and 100 times more acidic than pH 6. Always state pH values to 1 decimal place.

---

🟡 Standard — Regular Study (2d–2mo)

For NABTEB students who want thorough understanding of acids, bases, and salts.

Strong vs Weak Acids:

Property	Strong Acids	Weak Acids
Definition	Completely dissociates in water	Partially dissociates in water
Examples	HCl, $H_2SO_4$, $HNO_3$	$CH_3COOH$, $H_2CO_3$, $H_3PO_4$
pH (same concentration)	Lower pH	Higher pH
Electrical conductivity	Higher	Lower
Reaction rate (with metals)	Faster	Slower

Strong vs Weak Bases:

Property	Strong Bases	Weak Bases
Definition	Completely dissociates in water	Partially dissociates in water
Examples	$NaOH$, $KOH$, $Ca(OH)_2$	$NH_3$, $Mg(OH)_2$
pH (same concentration)	Higher pH	Lower pH

Neutralisation Reactions:

Acid + Base → Salt + Water

Examples: $$HCl + NaOH \rightarrow NaCl + H_2O$$ $$H_2SO_4 + 2KOH \rightarrow K_2SO_4 + 2H_2O$$ $$HNO_3 + NH_3 \rightarrow NH_4NO_3$$

Salt Preparation:

1. Direct combination: Metal + Non-metal → Salt $$2Na + Cl_2 \rightarrow 2NaCl$$

2. Acid + Metal: (for soluble salts — metals above hydrogen in reactivity series) $$Zn + H_2SO_4 \rightarrow ZnSO_4 + H_2$$

3. Acid + Base (Neutralisation): $$HCl + NaOH \rightarrow NaCl + H_2O$$

4. Acid + Carbonate: $$2HCl + Na_2CO_3 \rightarrow 2NaCl + CO_2 + H_2O$$

5. Precipitation (double decomposition): $$AgNO_3(aq) + NaCl(aq) \rightarrow AgCl(s) + NaNO_3(aq)$$

Types of Salts:

• Soluble salts: Nitrates, ammonium salts, chlorides (except AgCl, PbCl₂), sulphates (except BaSO₄, PbSO₄)
• Insoluble salts: Carbonates (except Na₂CO₃, K₂CO₃), hydroxides (except NaOH, KOH, Ca(OH)₂)

pH Calculations:

$$pH = -\log[H^+]$$

For pure water at 25°C: $[H^+] = 10^{-7}$ mol/L, so pH = 7.

For a strong acid with concentration $c$ mol/L: $$[H^+] = c \times \text{basicity of acid}$$

For a strong base with concentration $c$ mol/L: $$[OH^-] = c \times \text{acidity of base}$$

⚡ NABTEB Exam Tip: During a titration, the indicator must change colour at the equivalence point (when acid and base have reacted completely). Phenolphthalein is suitable for strong acid-strong base titrations (endpoint pH ~9). Methyl orange is suitable for strong acid-weak base titrations (endpoint pH ~4).

---

🔴 Extended — Deep Study (3mo+)

Comprehensive coverage for thorough NABTEB preparation.

Conjugate Acid-Base Pairs:

In Bronsted-Lowry theory, every acid has a conjugate base, and every base has a conjugate acid:

$$HA + H_2O \rightleftharpoons A^- + H_3O^+$$

• $HA$ is the acid (donates $H^+$)
• $A^-$ is the conjugate base of $HA$ (what remains after $H^+$ is lost)
• $H_2O$ is the base (accepts $H^+$)
• $H_3O^+$ is the conjugate acid of $H_2O$

Examples of Conjugate Pairs:

Acid	Conjugate Base
HCl	Cl⁻
$H_2SO_4$	$HSO_4^-$
$H_3O^+$	$H_2O$
$CH_3COOH$	$CH_3COO^-$
$NH_4^+$	$NH_3$
$H_2O$	$OH^-$

Acid Strength and $K_a$:

For weak acids: $HA \rightleftharpoons H^+ + A^-$ $$K_a = \frac{[H^+][A^-]}{[HA]}$$

Larger $K_a$ = stronger acid.

pKa = $-\log K_a$. Smaller pKa = stronger acid.

Common $K_a$ values at 25°C:

Acid	$K_a$	pKa
HCl	~10⁷	-7
$H_2SO_4$ (first proton)	~10³	-3
$CH_3COOH$	1.8 × 10⁻⁵	4.74
$H_2CO_3$	4.3 × 10⁻⁷	6.37
$H_3PO_4$	7.5 × 10⁻³	2.12

Base Strength and $K_b$:

For weak bases: $B + H_2O \rightleftharpoons BH^+ + OH^-$ $$K_b = \frac{[BH^+][OH^-]}{[B]}$$

$$K_a \times K_b = K_w = 10^{-14}$$

This means: strong acids have weak conjugate bases, and vice versa.

Buffer Solutions:

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

Acidic buffer: Weak acid + its conjugate base (e.g., $CH_3COOH/CH_3COONa$) Basic buffer: Weak base + its conjugate acid (e.g., $NH_3/NH_4Cl$)

Buffer action: When $H^+$ is added, it reacts with the conjugate base; when $OH^-$ is added, it reacts with the weak acid.

Henderson-Hasselbalch Equation: $$pH = pK_a + \log\frac{[\text{conjugate base}]}{[\text{acid}]}$$

Indicators:

Indicator	pH Range	Colour Change
Methyl orange	3.1 – 4.4	Red → Yellow
Methyl red	4.4 – 6.2	Red → Yellow
Bromothymol blue	6.0 – 7.6	Yellow → Blue
Phenolphthalein	8.3 – 10.0	Colourless → Pink
Litmus	5.5 – 8.2	Red → Blue

Titration Curves:

Strong acid vs Strong base:

• pH starts low, rises gradually
• Sharp change near equivalence point (pH 4–10)
• Equivalence point at pH = 7

Weak acid vs Strong base:

• pH starts low, rises more steeply initially
• Weak acid buffer region visible
• Equivalence point at pH > 7 (because conjugate base is basic)

Strong acid vs Weak base:

• Equivalence point at pH < 7 (because conjugate acid is acidic)

Hydrolysis of Salts:

Salts of strong acid + strong base: Neutral (pH ≈ 7) Salts of weak acid + strong base: Basic (pH > 7) — e.g., $Na_2CO_3$ Salts of strong acid + weak base: Acidic (pH < 7) — e.g., $NH_4Cl$ Salts of weak acid + weak base: Depends on relative strengths

Ostwald’s Dilution Law:

Degree of dissociation ($\alpha$) for weak electrolytes: $$\alpha = \sqrt{\frac{K}{c}}$$

As concentration ($c$) decreases, degree of dissociation increases.

⚡ NABTEB Quick Reference:

• $pH = -\log[H^+]$
• $K_a \times K_b = K_w = 10^{-14}$
• $pH + pOH = 14$
• Strong acids: HCl, $H_2SO_4$, $HNO_3$ (completely ionised)
• Weak acids: $CH_3COOH$, $H_2CO_3$, $H_3PO_4$ (partially ionised)
• Neutralisation: Acid + Base → Salt + Water
• Buffer: resists pH change
• Indicator colour change range must bracket the equivalence point pH