Aqueous Solutions and Reactions

Reactions in Aqueous Solutions

• Chemical reactions often occur in aqueous solutions.

States in Equations

• (s): solid, (l): liquid, (g): gas
• (aq): aqueous (dissolved in water)
• (↓): precipitate or solid
• (↑): gas produced

Formation of Aqueous Solutions

• Dissociation: Ionic compounds separate from crystal lattice.
• Ionization: Covalent compounds (polar molecules) form ions in water.

Dissociation Example

• $NaCℓ(s) \rightarrow Na^+(aq) + Cℓ^-(aq)$
• Electrolyte: Ions that conduct electricity in water.

Ionization Example

• $HCℓ(g) + H2O(ℓ) \rightarrow H3O^+(aq) + Cℓ^-(aq)$
• Electrolyte

Dissolving Molecules (e.g., Sugar)

• $C12H22O11(s) + H2O(ℓ) \rightarrow C12H22O11(aq)$
• No ions formed; does not conduct electricity.

Dissolution Process

• Dissolving/Dissolution: Ionic solid separates into ions, or covalent compounds form ions in water.
• Hydration: Ions surrounded by water molecules.

Solution Terminology

• Solution: Homogenous mixture.
• Solubility: Max solute that dissolves in a solvent.
• Solute: Dissolved substance (lesser amount).
• Solvent: Dissolving substance (greater amount).

Electrolytes and Conductivity

• Electrolyte: Liquid with free ions that conducts current.
• Free ions are essential for electrical conductivity.

Factors Affecting Conductivity

• Concentration: Higher concentration, greater conductivity.
• Solubility: More ions dissolved, greater conductivity.
• Substance Type: Some dissolve but don't conduct.
• Temperature: Affects solubility.

Precipitation Reactions

• Precipitate: Insoluble substance formed during a reaction.

Solubility Rules (Anions)

• Nitrate (NO3-): All cations soluble.
• Chloride (Cℓ-), Bromide (Br-), Iodide (I-): Mostly soluble, except with $Ag^+, Hg^+, Pb^{2+}$.
• Sulphate (SO42-): Mostly soluble, except with $Ca^{2+}, Sr^+, Ba^{2+}, Ag^+, Pb^{2+}$.
• Carbonate (CO32-), Hydroxide (OH-), Sulphide (S2-): Insoluble except with specific cations like $Li^+, Na^+, K^+, NH4^+$.

Types of Equations for Precipitation Reactions

• Molecular Equation: Reactants/products as molecules.
• Ionic Equation: Soluble substances as ions, precipitate as a molecule.
• Net Ionic Equation: Only participating ions; spectator ions omitted.

Tests for Anions

• Halides (Bromide, Chloride, Iodide):
  • Add $AgNO3$, precipitate indicates halide presence (white for $Cℓ^-$, cream for $Br^-$, pale-yellow for $I^-$.
  • Add dilute $HNO3$ to remove $CO3^{2-}$, confirming halide.
• Sulphates:
  • Add $BaCℓ2$ or $Ba(OH)2$, white precipitate ($BaSO4$) indicates sulphate.
  • Add concentrated nitric acid to confirm (precipitate remains if sulphate).
• Carbonates:
  • React with dilute acid to produce $CO2$ gas.
  • $CO2$ turns limewater milky.

Ion Exchange Reactions

• Reactions where ions are exchanged (e.g., precipitation reactions).

Gas Formation Reactions

• Acid + Metal → Salt + Hydrogen
• Metal Oxide + Heat → Metal + Oxygen
• Acid + Carbonate → Salt + Water + Carbon Dioxide

Tests for Gases

• Oxygen: Glowing splint bursts into flames.
• Hydrogen: Burning splint pops.
• Carbon Dioxide: Turns limewater milky.

Acid-Base Reactions

• Acids neutralize bases.

Properties of Acids

• Taste acidic.
• Influence indicators.
• React with water: $H^+ + H2O \rightarrow H3O^+$
• Neutralize bases.

Properties of Bases

• Taste bitter, feel soapy.
• Influence indicators.
• Conduct current in aqueous solution (except $NH3$).
• Form hydroxide ions ($OH^-$) in water.
• Neutralize acids.

Neutralization Reaction

• Acid + Base → Salt + Water
• $H^+(aq) + OH^-(aq) \rightarrow H2O(ℓ)$

Redox Reactions

• Oxidation: Loss of electrons.
• Reduction: Gain of electrons.
• OIL RIG (Oxidation Is Loss, Reduction Is Gain).