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Chapter 6: Acid, Base and Salt

Form 4 Chemistry Bab 6: Acid, Base and Salt

6.1 The Role of Water in Showing Acidic and Alkaline Properties

An acid is a chemical substance that ionises in water to produce hydrogen ions ($\text{H}^+$) or hydroxonium ions ($\text{H}_3\text{O}^+$).

A base is a substance that reacts with acids to form salt and water. An alkali is a base that is soluble in water and ionises to produce hydroxide ions ($\text{OH}^-$).

Importance of Water

  • Acids and alkalis show their acidic or alkaline properties only in the presence of water because water allows them to ionise and release free-moving $\text{H}^+$ or $\text{OH}^-$ ions.
  • In the absence of water (e.g., dry gas or dissolved in organic solvents like glacial ethanoic acid in propanone), acids exist as neutral molecules and do not turn dry litmus paper red or conduct electricity.

Basicity of Acids

Basicity refers to the number of hydrogen ions ($\text{H}^+$) produced by one molecule of acid when it ionises in water:

  • Monoprotic acid: Produces $1$ $\text{H}^+$ ion per molecule (e.g., $\text{HCl}$, $\text{HNO}_3$).
  • Diprotic acid: Produces $2$ $\text{H}^+$ ions per molecule (e.g., $\text{H}_2\text{SO}_4$).
  • Triprotic acid: Produces $3$ $\text{H}^+$ ions per molecule (e.g., $\text{H}_3\text{PO}_4$).

6.2 pH Value

pH is a logarithmic measure of the concentration of hydrogen ions ($\text{H}^+$) in a solution:

$$\text{pH} = -\log[\text{H}^+]$$ $$\text{pOH} = -\log[\text{OH}^-]$$ $$\text{pH} + \text{pOH} = 14$$
  • Acidic solution: $\text{pH} < 7$ (Higher $[\text{H}^+]$ concentration = Lower pH value).
  • Neutral solution: $\text{pH} = 7$.
  • Alkaline solution: $\text{pH} > 7$ (Higher $[\text{OH}^-]$ concentration = Higher pH value).

6.3 Strength of Acids and Alkalis

  • Strong Acid: An acid that completely ionises (100%) in water to produce a high concentration of $\text{H}^+$ ions (e.g., $\text{HCl}$, $\text{HNO}_3$, $\text{H}_2\text{SO}_4$). $$\text{HCl}(aq) \rightarrow \text{H}^+(aq) + \text{Cl}^-(aq)$$
  • Weak Acid: An acid that partially ionises in water to produce a low concentration of $\text{H}^+$ ions (e.g., $\text{CH}_3\text{COOH}$, $\text{H}_2\text{CO}_3$). $$\text{CH}_3\text{COOH}(aq) \rightleftharpoons \text{CH}_3\text{COO}^-(aq) + \text{H}^+(aq)$$
  • Strong Alkali: Completely ionises in water to produce high concentration of $\text{OH}^-$ ions (e.g., $\text{NaOH}$, $\text{KOH}$).
  • Weak Alkali: Partially ionises in water to produce low concentration of $\text{OH}^-$ ions (e.g., $\text{NH}_3(aq)$).

6.4 Chemical Properties of Acids and Alkalis

Reactions of Acids

  • $\text{Acid} + \text{Reactive Metal} \rightarrow \text{Salt} + \text{Hydrogen gas } (\text{H}_2)$
  • $\text{Acid} + \text{Metal Carbonate} \rightarrow \text{Salt} + \text{Water} + \text{Carbon dioxide gas } (\text{CO}_2)$
  • $\text{Acid} + \text{Base/Alkali} \rightarrow \text{Salt} + \text{Water}$

Reactions of Alkalis

  • $\text{Alkali} + \text{Acid} \rightarrow \text{Salt} + \text{Water}$
  • $\text{Alkali} + \text{Ammonium Salt} \rightarrow \text{Salt} + \text{Water} + \text{Ammonia gas } (\text{NH}_3)$
  • $\text{Alkali} + \text{Metal Cation Solution} \rightarrow \text{Insoluble Metal Hydroxide Precipitate} + \text{Salt}$

6.5 Concentration of Aqueous Solution and Volumetric Analysis

  • Concentration ($g\text{ dm}^{-3}$): $\text{Mass of solute (g)} / \text{Volume of solution (dm}^3)$
  • Molarity ($mol\text{ dm}^{-3}$): $\text{Moles of solute (mol)} / \text{Volume of solution (dm}^3)$
  • Interconversion: $\text{Concentration } (g\text{ dm}^{-3}) = \text{Molarity } (mol\text{ dm}^{-3}) \times \text{Molar Mass } (g\text{ mol}^{-1})$
  • Dilution Formula: $M_1V_1 = M_2V_2$
  • Neutralisation Titration Formula: $$\frac{M_a V_a}{M_b V_b} = \frac{a}{b}$$ where $a$ and $b$ are stoichiometric coefficients from the balanced equation $a\text{Acid} + b\text{Base} \rightarrow \text{Products}$.

6.6 Neutralisation

Neutralisation is the reaction between an acid and a base/alkali to produce salt and water only. Ionic equation for neutralisation:

$$\text{H}^+(aq) + \text{OH}^-(aq) \rightarrow \text{H}_2\text{O}(l)$$

Applications in daily life: Toothpaste (alkaline) neutralises mouth acid; antacids ($\text{Mg(OH)}_2$) neutralise excess stomach acid; slaked lime ($\text{Ca(OH)}_2$) treats acidic soil.

6.7 Salts, Solubility, and Preparation Methods

A salt is an ionic compound formed when the hydrogen ion ($\text{H}^+$) from an acid is replaced by a metal ion or an ammonium ion ($\text{NH}_4^+$).

Solubility Rules Summary

  • NO3- (Nitrates): ALL soluble.
  • Na+, K+, NH4+ (SPAN salts): ALL soluble.
  • Cl- (Chlorides): ALL soluble EXCEPT PBA ($\text{PbCl}_2$, $\text{BaCl}_2$ is soluble, $\text{AgCl}$). *(Note: $\text{PbCl}_2$ is soluble in hot water)*.
  • SO42- (Sulfates): ALL soluble EXCEPT PBC ($\text{PbSO}_4$, $\text{BaSO}_4$, $\text{CaSO}_4$).
  • CO32- (Carbonates): ALL INSOLUBLE EXCEPT Sodium, Potassium, Ammonium carbonates.

Salt Preparation Methods

  1. Soluble Non-SPAN Salts: Acid + Excess Insoluble Metal / Metal Oxide / Metal Carbonate (Filter, evaporate to saturated state, cool, filter crystals, dry with filter paper).
  2. Soluble SPAN Salts: Acid + Alkali via Titration (using indicator to find endpoint, repeat without indicator, evaporate, cool, filter, dry).
  3. Insoluble Salts: Double Decomposition Reaction (Precipitation) by mixing two aqueous soluble salt solutions.

6.8 Qualitative Analysis of Salts (Qualitative Testing)

Action of Heat on Salts

  • $\text{CO}_3^{2-}$ decomposes to release $\text{CO}_2$ gas (turns lime water cloudy).
  • $\text{NO}_3^-$ decomposes to release $\text{O}_2$ and brown $\text{NO}_2$ gas.
  • Oxide residue colors:
    • $\text{ZnO}$: Yellow when hot, white when cold.
    • $\text{PbO}$: Brown when hot, yellow when cold.
    • $\text{CuO}$: Black residue.

Confirmatory Tests for Anions

  • $\text{CO}_3^{2-}$: Add dilute acid $\rightarrow$ effervescence of gas that turns lime water cloudy.
  • $\text{Cl}^-$: Add dilute $\text{HNO}_3$, then add $\text{AgNO}_3$ solution $\rightarrow$ white precipitate ($\text{AgCl}$).
  • $\text{SO}_4^{2-}$: Add dilute $\text{HCl}$, then add $\text{BaCl}_2$ solution $\rightarrow$ white precipitate ($\text{BaSO}_4$).
  • $\text{NO}_3^-$ (Brown Ring Test): Add dilute $\text{H}_2\text{SO}_4$, add $\text{FeSO}_4$ solution, carefully slant test tube and trickle concentrated $\text{H}_2\text{SO}_4$ down tube $\rightarrow$ brown ring forms at junction.

Cation Testing with Sodium Hydroxide ($\text{NaOH}$) and Ammonia ($\text{NH}_3$)

  • $\text{Cu}^{2+}$: Blue precipitate, insoluble in excess $\text{NaOH}$, soluble in excess $\text{NH}_3$ to form dark blue solution.
  • $\text{Fe}^{2+}$: Green precipitate, insoluble in excess.
  • $\text{Fe}^{3+}$: Brown precipitate, insoluble in excess.
  • $\text{Al}^{3+}, \text{Pb}^{2+}, \text{Zn}^{2+}$: White precipitate, soluble in excess $\text{NaOH}$. To differentiate:
    • In excess $\text{NH}_3$: Only $\text{Zn}^{2+}$ precipitate dissolves.
    • Add $\text{KI}$ solution: $\text{Pb}^{2+}$ forms yellow precipitate ($\text{PbI}_2$), whereas $\text{Al}^{3+}$ forms no precipitate.
  • $\text{NH}_4^+$: No precipitate with $\text{NaOH}$; on heating, releases pungent $\text{NH}_3$ gas that turns moist red litmus paper blue.
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