Acids, Bases & Salts – Comprehensive Bullet-Point Notes

Syllabus Focus (Valid for Examinations in & after March 2026)

• Study of acids, bases & salts with emphasis on:
– Ions present in mineral acids, alkalis & salts and in their aqueous solutions.
– Simple molecular definitions, characteristic properties, litmus & pH tests.
– Ionisation/dissociation equations for acids, bases & salts.
– Introduction and use of pH scale / Universal indicator.
– Classification & definition of normal, acid, basic salts.
– Decomposition of carbonates, bicarbonates, sulphites & sulphides by dilute acids (lab work mandatory).
– Six preparative routes for normal salts: direct combination, displacement, precipitation, neutralisation of an insoluble base, neutralisation of an alkali (titration), action of dilute acids on (bi)carbonates.

Key Definitions (Molecular View)

• Acid – compound that yields only hydronium ions in water. Example dissociation: $\text{HCl} + 2\,\text{H}2\text{O} \;\longrightarrow\; \text{H}3\text{O}^+ + \text{Cl}^-$
• Base – oxide or hydroxide of a metal (or $\text{NH}3$ derivative) that reacts with hydronium ions to give salt + water only. E.g. $\text{CuO} + 2\,\text{HCl} \rightarrow \text{CuCl}2 + \text{H}2\text{O}$ • Alkali – water-soluble base that furnishes only hydroxyl ions. Example: $\text{NaOH}{(aq)} \rightarrow \text{Na}^+ + \text{OH}^-$ (All alkalis are bases; converse is not true.)
• Salt – product of partial/complete replacement of ionisable $\text{H}^+$ of an acid by a metal or $\text{NH}_4^+$ . Dissociates to give cation ≠ $\text{H}^+$ & anion ≠ $\text{OH}^-$ .

Indicator Basics

• Indicators = weak organic acids/bases whose ionic/molecular forms have different colours → colour governed by $[\text{H}^+]$ .
• Litmus: purple (neutral) → red (acidic) / blue (alkaline).
• Methyl orange: pink (pH < 3.1) → orange (3.1–4.4) → yellow (>4.4).
• Phenolphthalein: colourless (pH < 8.3) → pink (8.3–10.0).
• Universal indicator / pH paper: continuous colour graduation, pH 0 (red) → 7 (green) → 14 (violet).

pH Concept

• Ionic product of water: $Kw = [\text{H}^+][\text{OH}^-] = 10^{-14}\;\text{mol}^2\,\text{dm}^{-6}\;(25^{\circ}\text{C})$ . • $\text{pH} = -\log{10}[\text{H}^+]$ (mol dm$^{-3}$).
• Pure water: $[\text{H}^+] = 10^{-7}$ → pH 7 (neutral).
• pH < 7 acidic, pH > 7 basic; farther from 7 → stronger.

Classification of Acids

By source
– Organic (plant-derived): citric, oxalic, tartaric, acetic (weak).
– Inorganic (mineral): $\text{HCl},\;\text{H}2\text{SO}4,\;\text{HNO}_3$ (generally strong).
By composition
– Hydracids: only H + non-metal (no O) e.g. $\text{HCl},\;\text{HI}$ .
– Oxyacids: H + element + O e.g. $\text{HNO}3,\;\text{H}2\text{SO}_4$ .
By strength (degree of ionisation)
– Strong: nearly 100 % dissociation → solution contains almost only ions. Examples $\text{HCl},\;\text{H}2\text{SO}4,\;\text{HNO}_3$ .
– Weak: partial dissociation, equilibrium mixture of molecules & ions (acetic, citric, carbonic, formic).
By concentration
– Concentrated: high proportion of acid (>1 mol dm$^{-3}$).
– Dilute: <1 mol dm$^{-3}$.
By basicity (no. of ionisable hydrogens)
– Monobasic: $\text{HCl},\;\text{HNO}3,\;\text{CH}3\text{COOH}$ .
– Dibasic: $\text{H}2\text{SO}4,\;\text{H}2\text{CO}3$ (two-step ionisation).
– Tribasic: $\text{H}3\text{PO}4$ (three-step, gives 3 salt types).

Classification of Bases / Alkalis

• By acidity (replaceable $\text{OH}^-$ per formula unit)
– Monoacidic: $\text{NaOH},\;\text{KOH},\;\text{NH}4\text{OH}$ . – Diacidic: $\text{Ca(OH)}2,\;\text{Cu(OH)}2$ (sparingly soluble). – Triacidic: $\text{Al(OH)}3,\;\text{Fe(OH)}3$ (insoluble). • Strength parallels acid logic: strong alkalis (LiOH, NaOH, KOH) vs weak (NH$4$OH, Ca(OH)$_2$). Concentration terminology as for acids.

Hydronium & Hydroxyl Ion Formation (Particle Pictures)

• Hydronium: $\text{H}^+$ from acid binds lone pair on O of $\text{H}2\text{O}$ forming $\text{H}3\text{O}^+$ via coordinate covalent bond $\text{O}\rightarrow\text{H}$ .
$\text{HCl} \xrightarrow[]{\text{water}} \text{H}^+ + \text{Cl}^-;\quad \text{H}^+ + \text{H}2\text{O} \rightarrow \text{H}3\text{O}^+$
• Hydroxyl: $\text{H}2\text{O}$ donates $\text{H}^+$ to electron-rich $\text{NH}3$ → $\text{NH}_4^+ + \text{OH}^-$ , bond $\text{N}\rightarrow\text{H}$ .

Preparation of Acids (Representative)

Direct union of non-metal with H$2$ (at elevated T/light): $\text{H}2 + \text{Cl}_2 \rightarrow 2\,\text{HCl}$ .
Dissolving acidic oxides in water: $\text{CO}2 + \text{H}2\text{O} \rightarrow \text{H}2\text{CO}3$ ; $\text{SO}3 + \text{H}2\text{O} \rightarrow \text{H}2\text{SO}4$ .
From salts via displacement with conc. $\text{H}2\text{SO}4$ (< $200^{\circ}\text{C}$ ): $\text{NaCl} + \text{H}2\text{SO}4 \rightarrow \text{NaHSO}_4 + \text{HCl}\uparrow$ .
Oxidation of non-metals: $\text{S} + 6\,\text{HNO}3 \rightarrow \text{H}2\text{SO}4 + 2\,\text{H}2\text{O} + 6\,\text{NO}_2$ .

Preparation of Bases (Representative)

• Metal + O$2$ → basic oxide; soluble oxide + H$2$O → alkali. Example $2\text{K} + 2\text{H}2\text{O} \rightarrow 2\text{KOH} + \text{H}2$ .
• Thermal decomposition of carbonates/nitrates: $2\text{Pb(NO}3)2 \xrightarrow{\Delta} 2\text{PbO} + 4\,\text{NO}2 + \text{O}2$ .
• Precipitating hydroxides: $\text{AlCl}3 + 3\text{NaOH} \rightarrow 3\text{NaCl} + \text{Al(OH)}3\downarrow$ .

Physical Properties Snapshot

Acids: sour, corrosive (strong mineral acids), turn blue litmus red.
Bases/alkalis: bitter, soapy, caustic (NaOH/KOH), turn red litmus blue.

Core Chemical Reactions

• Neutralisation (ionic): $\text{H}^+{(aq)} + \text{OH}^-{(aq)} \rightarrow \text{H}2\text{O}{(l)}$ ; exothermic (heat of neutralisation per Eq. gram-equiv.).
• Metal + dilute acid (except $\text{HNO}3$ for Na/K/Ca) liberates H$2$: $\text{Zn} + 2\text{HCl} \rightarrow \text{ZnCl}2 + \text{H}2$ .
• Less-volatile acid displaces more-volatile acid from salt on heating: $\text{NaNO}3 + \text{H}2\text{SO}4 \rightarrow \text{NaHSO}4 + \text{HNO}3\uparrow$ . • Alkali + ammonium salt → ammonia liberation: $\text{NH}4\text{Cl} + \text{NaOH} \xrightarrow{\Delta} \text{NaCl} + \text{NH}3\uparrow + \text{H}2\text{O}$ .

Uses (Selected)

Acids: boric acid (eye-wash), citric (preserves), oxalic (ink remover), carbonic (soft drinks), tartaric (baking powder), acetic (vinegar), HCl (pickling metals).
Bases: NaOH (soap), Ca(OH)$2$ (bleaching powder), Mg(OH)$2$ / Al(OH)$3$ (antacids), Ca(OH)$2$ (water softening, degreasing).

Acid Rain Reference

• pH < 5.6 precipitation containing $\text{H}2\text{SO}4,\;\text{H}2\text{SO}3,\;\text{HNO}3,\;\text{HNO}2$ formed from anthropogenic $\text{SO}2,\;\text{NO}x$ .
• Reaction chain: $\text{SO}2 + \text{H}2\text{O} \rightarrow \text{H}2\text{SO}3$ ; $2\text{SO}2 + \text{O}2 \rightarrow 2\text{SO}_3$ etc.
• Effects: soil nutrient leaching, aquatic pH drop, infrastructural corrosion.

Salts – Detailed Classification

Normal salts: complete H-replacement; no ionisable H. Eg $\text{Na}2\text{SO}4$ .
Acid salts: partial replacement, retain replaceable H; show acidic properties. Eg $\text{NaHSO}4,\;\text{NaHCO}3$ .
Basic salts: contain metallic cation + $\text{OH}^-$ + acid anion (from incomplete neutralisation of polybasic base). Eg $\text{Cu(OH)Cl},\;\text{Cu(OH)NO}_3$ .
Double salts: crystalline combination of two simple salts that dissociate to give both ions in solution (e.g. potash alum $\text{K}2\text{SO}4\,\cdot\,\text{Al}2(\text{SO}4)3\,\cdot\,24\text{H}2\text{O}$ , Mohr’s salt $\text{(NH}4)2\text{SO}4\,\cdot\,\text{FeSO}4\,\cdot\,6\text{H}_2\text{O}$ ).
Mixed salts: two different positive or negative radicals inside one formula, e.g. $\text{NaKCO}_3,\;\text{Ca(OCl)Cl}$ .
Complex salts: yield complex ion in solution; e.g. $\text{K}2[\text{HgI}4],\;\text{Na[Ag(CN)}2],\;\text{Na}2\text{ZnO}2,\;[\text{Cu(NH}3)4]\text{SO}4$ .

Solubility Rules (Cold Water)

• Always soluble: $\text{Na}^+,\;\text{K}^+,\;\text{NH}4^+$ salts; all nitrates/nitrites; all bicarbonates; most chlorides & sulphates. • Sparingly/insoluble exceptions: – $\text{PbSO}4,\;\text{AgSO}4,\;\text{CaSO}4,\;\text{BaSO}4$ . – $\text{PbCl}2$ (soluble in hot water), $\text{AgCl},\;\text{HgCl}$ .
– All sulphites, sulphides, carbonates, oxides, hydroxides, phosphates (except Na/K/NH$_4$ derivatives).

General Synthetic Routes for Salts

Direct combination (synthesis): $2\text{Fe} + 3\text{Cl}2 \rightarrow 2\text{FeCl}3$ .
Displacement (active metal + dil. acid): $\text{Zn} + \text{H}2\text{SO}4 \rightarrow \text{ZnSO}4 + \text{H}2$ .
Precipitation (double decomposition): $\text{Pb(NO}3)2 + 2\text{NaCl} \rightarrow 2\text{NaNO}3 + \text{PbCl}2\downarrow$ .
Neutralisation of insoluble base: $\text{CuO} + \text{H}2\text{SO}4 \rightarrow \text{CuSO}4 + \text{H}2\text{O}$ .
Neutralisation of an alkali (titration): $2\text{NaOH} + \text{H}2\text{SO}4 \rightarrow \text{Na}2\text{SO}4 + 2\text{H}_2\text{O}$ .
Action of dilute acids on (bi)carbonates: $\text{Na}2\text{CO}3 + 2\text{HCl} \rightarrow 2\text{NaCl} + \text{H}2\text{O} + \text{CO}2$ .

Step-wise Preparations (Laboratory Protocol Highlights)

• Iron(III) chloride (volatile, deliquescent): heat Fe wire in dry Cl$2$ stream; sublimate $\text{FeCl}3$ vapours into chilled receiver; store with fused $\text{CaCl}2$ . • FeSO$4$·7H$2$O / ZnSO$4$·7H$2$O: treat Fe or Zn with dil. $\text{H}2\text{SO}4$ → filter, concentrate, crystallise green (Fe) or colourless (Zn) heptahydrate. • PbCl$2$: convert insoluble $\text{PbO}$ or $\text{PbCO}3$ to soluble $\text{Pb(NO}3)2$ via dil. $\text{HNO}3$ , then precipitate with $\text{NaCl}$ ; dissolve ppt in hot water & cool → needle crystals.
• CaCO$3$: mix aq. $\text{CaCl}2$ + $\text{Na}2\text{CO}3$ → white ppt, wash & dry to amorphous powder.
• CuSO$4$·5H$2$O: react black $\text{CuO}$ (or blue $\text{Cu(OH)}2$ ) with warm dil. $\text{H}2\text{SO}4$ ; filter, concentrate & crystallise blue pentahydrate. • Na$2$SO$4$·10H$2$O (titration): standardise neutralisation of NaOH with dil. $\text{H}2\text{SO}4$ using phenolphthalein → evaporate neutral solution & crystallise Glauber’s salt.

Hydrolysis Behaviour of Salts (Solution pH)

Salt type	Example	Products of hydrolysis	Nature of solution
Weak base + strong acid	$\text{NH}_4\text{Cl}$	$\text{NH}_4\text{OH} + \text{HCl}$	Slightly acidic
Strong base + weak acid	$\text{NaHCO}<em>3,\;\text{Na}</em>2\text{CO}_3$	$\text{NaOH} + \text{H}<em>2\text{CO}</em>3$	Alkaline
Strong base + strong acid	$\text{NaCl}$	$\text{NaOH} + \text{HCl}$ (no net change)	Neutral

Water of Crystallisation, Deliquescence & Efflorescence

• Water of crystallisation fixed (loose) combination: gypsum ( $\text{CaSO}4\cdot2\text{H}2\text{O}$ ), blue vitriol ( $\text{CuSO}4\cdot5\text{H}2\text{O}$ ), washing soda ( $\text{Na}2\text{CO}3\cdot10\text{H}2\text{O}$ ). • Deliquescent salts absorb moisture until they liquefy: $\text{FeCl}3,\;\text{CaCl}2$ (anh.), $\text{MgCl}2$ .
• Efflorescent salts lose water, crumble to powder: $\text{Na}2\text{SO}4\cdot10\text{H}2\text{O},\;\text{MgSO}4\cdot7\text{H}_2\text{O}$ .

Field / Industrial Significance of pH & Indicators

• Agriculture: match soil pH to crop (citrus: slightly alkaline, rice: acidic, sugarcane: neutral).
• Dairy: souring milk detected when pH drops below 6.6.
• Medicine: diagnostic pH testing of blood (≈7.4) & urine; antacid formulation.
• Technology: biochemical & organic syntheses carried out under controlled pH windows.

Representative pH Values (25 °C)

• $0.1\,\text{N}\;\text{HCl}$ ≈ pH 1.0; $0.1\,\text{N}\;\text{H}2\text{SO}4$ ≈ 1.2.
• Vinegar (acetic) ≈ 2.9; grape juice (tartaric) ≈ 3.5.
• Milk (lactic) ≈ 6.6; human blood ≈ 7.3.
• Sea water ≈ 8.5; $0.1\,\text{N}\;\text{NH}_4\text{OH}$ ≈ 11.1; $0.1\,\text{N}\;\text{NaOH}$ ≈ 13.0.

These bullet-point notes consolidate every major & minor concept, definition, equation, example, laboratory method and real-world linkage mentioned across the transcript pages, ready for direct use as a full replacement study guide.