Ionic Compounds

Many compounds are formed by the combination of positively charged ions (cations, usually from metals) and negatively charged ions (anions, usually from non-metals or polyatomic groups). These are held together by electrostatic forces and are called ionic compounds. Examples include common salt (NaCl), magnesium chloride (MgCl₂), etc.

Dissociation

When ionic compounds dissolve in water, the polar water molecules surround the ions and pull them apart. This process, where the compound separates into its constituent ions in solution, is called dissociation.

Example: NaCl(s) --(water)--> Na⁺(aq) + Cl⁻(aq)

The ability of substances to dissociate into ions in water is crucial for understanding acids, bases, and electrolytes.

Svante Arrhenius proposed a theory to define acids and bases based on their behaviour in aqueous solutions:

• Acid: An acid is a substance which, on dissolving in water, gives rise to H⁺ (hydrogen ions) as the only cation. Example: HCl(aq) → H⁺(aq) + Cl⁻(aq). Note: H⁺ ions readily combine with water to form hydronium ions (H₃O⁺).
• Base: A base is a substance which, on dissolving in water, gives rise to OH⁻ (hydroxide ions) as the only anion. Example: NaOH(aq) → Na⁺(aq) + OH⁻(aq).

Strength (Based on Dissociation)

• Strong Acids/Bases: Dissociate almost completely in water, producing a high concentration of H⁺ or OH⁻ ions. Examples: HCl, H₂SO₄, HNO₃ (acids); NaOH, KOH (bases).
• Weak Acids/Bases: Dissociate only partially in water, establishing an equilibrium and producing a low concentration of H⁺ or OH⁻ ions. Examples: CH₃COOH (acetic acid), H₂CO₃ (carbonic acid); NH₄OH (ammonium hydroxide).

Alkalis

Bases that are highly soluble in water are specifically called alkalis. Examples: NaOH, KOH, Ca(OH)₂ (sparingly soluble but considered an alkali).

Basicity and Acidity

• Basicity of an Acid: The number of H⁺ ions obtainable from one molecule of an acid (e.g., HCl - monobasic, H₂SO₄ - dibasic, H₃PO₄ - tribasic).
• Acidity of a Base: The number of OH⁻ ions obtainable from one molecule of a base (e.g., NaOH - monoacidic, Ca(OH)₂ - diacidic, Al(OH)₃ - triacidic).

The concentration of a solution tells us the amount of solute dissolved in a given amount of solvent or solution. Common units include:

• Grams per litre (g/L): Mass of solute (g) / Volume of solution (L).
• Molarity (M): Moles of solute / Volume of solution (L).

Indicators

Indicators are substances that change colour depending on whether they are in an acidic, basic, or neutral solution. They help us identify the nature of a solution.

Universal Indicator: A mixture of indicators showing a range of colours corresponding to different pH values.

The pH Scale

Developed by Soren Sorensen, the pH scale measures the hydrogen ion concentration ([H⁺]) in a solution. It is a logarithmic scale from 0 to 14.

• pH < 7: Acidic solution
• pH = 7: Neutral solution (e.g., pure water)
• pH > 7: Basic (alkaline) solution

The formula is: pH = -log₁₀[H⁺]. A lower pH means higher acidity (higher [H⁺]).

When an acid reacts with a base, they neutralize each other's properties, forming salt and water. This is a neutralization reaction.

General Equation: Acid + Base → Salt + Water

Example: HCl(aq) + NaOH(aq) → NaCl(aq) + H₂O(l)

Ionic Equation: H⁺(aq) + OH⁻(aq) → H₂O(l)

Definition and Formation

A salt is an ionic compound formed during neutralization, consisting of a cation from the base and an anion from the acid.

Types of Salts (Based on Hydrolysis)

• Neutral Salts: Formed from strong acid + strong base (e.g., NaCl, KNO₃). Solution pH ≈ 7.
• Acidic Salts: Formed from strong acid + weak base (e.g., NH₄Cl, CuSO₄). Solution pH < 7.
• Basic Salts: Formed from weak acid + strong base (e.g., CH₃COONa, Na₂CO₃). Solution pH > 7.

Water of Crystallization

Some salts incorporate a fixed number of water molecules into their crystal structure during crystallization. This is called water of crystallization, and the salts are called hydrated salts (e.g., CuSO₄·5H₂O - blue vitriol).

Heating hydrated salts removes this water, often causing a change in colour and structure (e.g., blue CuSO₄·5H₂O becomes white anhydrous CuSO₄).

Properties of Salts

Generally crystalline solids, high melting/boiling points, variable solubility, conduct electricity when molten or dissolved, can be neutral/acidic/basic in solution.

Electrolytes

Substances that conduct electricity when dissolved in water or in the molten state because they produce mobile ions.

• Strong Electrolytes: Dissociate completely (strong acids, strong bases, most salts). High conductivity.
• Weak Electrolytes: Dissociate partially (weak acids, weak bases). Low conductivity.
• Non-electrolytes: Do not dissociate or produce ions (e.g., sugar, alcohol). Do not conduct electricity.

Electrolysis

The process of using direct electric current to decompose an electrolyte. Chemical reactions (oxidation at anode, reduction at cathode) occur.

Electrolysis of Water: Pure water is a poor conductor, so a small amount of acid (like H₂SO₄) is added. Passing current decomposes water:

2H₂O(l) --(Electrolysis)--> 2H₂(g) [at Cathode] + O₂(g) [at Anode]

The volume of H₂ produced is twice the volume of O₂.