Acids, Bases, Salts and pH

🟢 Lite — Quick Review (1h–1d)

Rapid summary for last-minute revision before your NECO exam.

Arrhenius Definitions:

• Acid: A substance that produces H⁺ ions in aqueous solution. Example: $\text{HCl} \rightarrow \text{H}^+ + \text{Cl}^-$
• Base: A substance that produces OH⁻ ions in aqueous solution. Example: $\text{NaOH} \rightarrow \text{Na}^+ + \text{OH}^-$

Bronsted–Lowry Definitions (preferred at NECO level):

• Acid: Proton (H⁺) donor
• Base: Proton (H⁺) acceptor
• Conjugate acid–base pairs differ by one proton

pH Scale: $0$ to $14$

• pH $< 7$: acidic
• pH $= 7$: neutral (at 25°C)
• pH $> 7$: basic/alkaline

$$\text{pH} = -\log_{10}[\text{H}^+]$$ $$[\text{H}^+] = 10^{-\text{pH}}$$

Common Acids and Bases:

Strong Acids	Strong Bases
HCl, HBr, HI, HNO₃, H₂SO₄, HClO₄	NaOH, KOH, Ca(OH)₂, Ba(OH)₂

Salt: Ionic compound formed from the reaction of an acid and a base (neutralisation).

⚡ NECO Tip: Strong acids completely dissociate — their pH is very low. A $0.1$ M HCl solution has $[H^+] = 0.1$ M, so pH $= 1$. Weak acids only partially dissociate — a $0.1$ M CH₃COOH solution has $[H^+] \ll 0.1$ M, so pH $> 1$.

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🟡 Standard — Regular Study (2d–2mo)

Standard content for NECO Chemistry students with a few days to months.

pH Calculations

For strong monoprotic acids: $[\text{H}^+] = C$ (concentration), so $\text{pH} = -\log C$.

Example: Calculate pH of $0.001$ M HCl. $$[\text{H}^+] = 0.001 = 10^{-3} \text{ M} \Rightarrow \text{pH} = -\log(10^{-3}) = 3$$

pOH and the Relationship: $$\text{pH} + \text{pOH} = 14 \text{ (at } 25°C\text{)}$$ $$\text{pOH} = -\log[\text{OH}^-]$$

Example: Find pH of $0.01$ M NaOH. $$[\text{OH}^-] = 10^{-2} \text{ M} \Rightarrow \text{pOH} = 2 \Rightarrow \text{pH} = 14 - 2 = 12$$

Weak Acids and $K_a$:

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

$$[\text{H}^+] = \sqrt{K_a \times C} \quad \text{(for dilute solutions)}$$

Example: Find $[H^+]$ in $0.1$ M acetic acid ($K_a = 1.8 \times 10^{-5}$). $$[\text{H}^+] = \sqrt{1.8 \times 10^{-5} \times 0.1} = \sqrt{1.8 \times 10^{-6}} = 1.34 \times 10^{-3} \text{ M}$$ $$\text{pH} = -\log(1.34 \times 10^{-3}) = 2.87$$

Buffer Solutions:

A buffer resists pH change. Acidic buffer: weak acid + its salt (e.g., CH₃COOH + CH₃COONa).

Types of Salts:

• Acidic salts: Formed from strong acid + weak base (e.g., NH₄Cl — produces acidic solution)
• Basic salts: Formed from weak acid + strong base (e.g., Na₂CO₃ — produces basic solution)
• Neutral salts: Formed from strong acid + strong base (e.g., NaCl, K₂SO₄ — neutral solution)

Indicators and Titration:

Indicator	pH Range (approx.)	Colour Change
Methyl orange	3.1 – 4.4	Red → Yellow
Phenolphthalein	8.2 – 10.0	Colourless → Pink
Litmus	5.5 – 8.2	Red → Blue

⚡ NECO Common Mistakes:

• Confusing strong and weak acids — same concentration but very different pH values
• Forgetting that pH is a logarithmic scale — pH 3 has 10× more H⁺ than pH 4
• Not knowing which indicator to use for which titration
• In buffer calculations, using the wrong formula for $[H^+]$

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🔴 Extended — Deep Study (3mo+)

Comprehensive coverage for NECO and JAMB Chemistry preparation.

Strong Acid vs Weak Acid — Degree of Dissociation:

For strong acids, $\alpha \approx 1$ (complete dissociation). For weak acids: $$\alpha = \sqrt{\frac{K_a}{C}}$$

pH of Strong Acid–Strong Base Titration:

At equivalence point: pH = 7 (for strong acid + strong base).

Before equivalence: excess $H^+$ present → pH determined by excess acid. After equivalence: excess $OH^-$ present → determine pOH then convert to pH.

pH of Weak Acid–Strong Base Titration:

At equivalence point: pH > 7 (weak acid’s conjugate base makes solution slightly basic). The salt formed (e.g., CH₃COONa) hydrolyses: $$\text{CH}_3\text{COO}^- + \text{H}_2\text{O} \rightleftharpoons \text{CH}_3\text{COOH} + \text{OH}^-$$

Salt Hydrolysis:

For salt of weak acid and strong base: solution is basic ($pH > 7$). For salt of strong acid and weak base: solution is acidic ($pH < 7$). For salt of weak acid and weak base: pH depends on $K_a$ and $K_b$: $$\text{pH} = \frac{1}{2}\left(\text{p}K_a + \text{p}K_b\right) + \frac{1}{2}\log C$$

Solubility Product ($K_{sp}$):

For $\text{AB} \rightleftharpoons \text{A}^+ + \text{B}^-$: $K_{sp} = [\text{A}^+][\text{B}^-]$

Example: If $K_{sp}$ of AgCl = $1.8 \times 10^{-10}$, the maximum $[Ag^+]$ in solution is $\sqrt{1.8 \times 10^{-10}} = 1.34 \times 10^{-5}$ M.

Ion Product and Precipitation:

If ion product $> K_{sp}$: precipitation occurs. If ion product $< K_{sp}$: no precipitation (unsaturated solution).

Neutralisation Enthalpy:

The enthalpy of neutralisation for strong acid + strong base is approximately $-57.3$ kJ/mol (independent of which acid or base, since all produce $\text{H}_2\text{O}$ from $\text{H}^+ + \text{OH}^-$).

NECO/JAMB Patterns:

• NECO frequently asks: calculate pH of dilute acids; identify conjugate acid-base pairs; complete neutralisation titration calculations
• Hydrolysis of salts questions appear regularly in Section B
• Be able to sketch and label a titration curve (pH vs volume added)

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