PPT - M6 Focus1 Acid Base Properties

HSC Chemistry Module 6: Equilibrium and Acid Reactions

Focus 1: Properties of Acids and Bases

  • Inquiry Question: What is an acid and what is a base?

6.1.1: IUPAC Nomenclature and Properties of Acids and Bases

  • IUPAC: International Union of Pure and Applied Chemistry; responsible for naming compounds.

  • Nomenclature: Rules for naming compounds in science, differing for:

    • Inorganic compounds (not containing carbon)

    • Organic compounds (containing carbon)

    • Polymers

  • Types of Acids and Bases:

    • Inorganic Bases: oxides and hydroxides of metals, ammonia.

    • Organic Bases: typically contain nitrogen (e.g., DNA base pairs).

    • Inorganic Acids: derived from inorganic compounds (e.g., hydrochloric acid, sulfuric acid).

    • Organic Acids: derived from organic compounds, often with carboxylic acid groups (-COOH).

6.1.2: Indicators for Acids and Bases

  • Indicators: Substances (vegetable and synthetic dyes) that change colour based on solution acidity.

  • Original Method: Based on the effect of substances on vegetable dyes (e.g., litmus).

  • Investigation Task: Demonstrate characteristics of acids and bases using indicators (reference: Chemistry in Focus, page 122).

6.1.3: Acid Reactions

Types of Reactions:

  1. Neutralisation (Acid & Base) Reactions:

    • General Equation: ACID + BASE → SALT + WATER

    • Example: Hydrochloric Acid + Sodium Hydroxide → Sodium Chloride + Water

      • HCl + NaOH → NaCl + H2O

  2. Reaction with Carbonates:

    • General Equation: ACID + CARBONATE → SALT + CARBON DIOXIDE + WATER

    • Example: Sulfuric Acid + Magnesium Carbonate → Magnesium Sulfate + Carbon Dioxide + Water

  3. Dilute Acid & Metal Reactions:

    • General Equation: DILUTE ACID + METAL → SALT + HYDROGEN GAS

    • Example: Nitric Acid + Sodium → Sodium Nitrate + Hydrogen Gas

      • HNO3 + Na → NaNO3 + H2

6.1.4: Applications of Neutralisation

  • Skill Development:

    • Provide examples from everyday life and industrial processes.

    • Example Requirements:

      • Identify chemicals involved.

      • Write a balanced chemical equation.

      • Describe purpose and context of use.

      • Justify why this application is beneficial.

6.1.5: Enthalpy of Neutralisation

  • Definition:

    • Measure of total energy in a substance or group when reactants bond-breaking occurs.

  • Change in Enthalpy (ΔH):

    • Formula: ΔH = enthalpy of products - enthalpy of reactants

    • Measured in kJ mol-1

    • Positive ΔH: energy absorption (endothermic)

    • Negative ΔH: energy release (exothermic)

  • Standard Conditions:

    • 25°C (298K), 100kPa

    • Enthalpy changes during neutralisation: Generally negative, indicating exothermic nature.

6.1.6: Historical Definitions of Acids and Bases

Antoine Lavoisier (1776):

  • Proposed that acids contained oxygen.

  • Limitations:

    • Did not explain lack of acidity in metal oxides or common acids like HCl.

Humphry Davy (1815):

  • Suggested acids contained hydrogen.

  • Limitations:

    • Classified NH3 as acidic, which is not correct.

Svante Arrhenius (1884):

  • Defined acids as H+ producers, bases as OH- producers in solutions.

  • Limitations:

    • Did not account for solvent effects on strength; failed to explain ammonia's behavior.

Bronsted-Lowry Theory (1923):

  • Acids as proton donors, bases as proton acceptors.

  • Limitations:

    • Relied on solvent presence; could not explain acid-base reactions without a solvent.

Gilbert Lewis (1923):

  • Defined acids as electron pair acceptors and bases as electron pair donors.

  • Advancement:

    • Broader definition not limited to protons or solvents (e.g., BF3 reacting with NH3).

  • Limitations:

    • Did not explain strength differences or provide clarity on reactions of common acids like HCl.