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HSC Chemistry

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  1. Module 1: Properties and Structure of Matter
    1.1 Properties of Matter
  2. 1.2 Atomic Structure and Atomic Mass
  3. 1.3 Periodicity
  4. 1.4 Bonding
  5. Module 2: Introduction to Quantitative Chemistry
    2.1 Chemical Reactions and Stoichiometry
  6. 2.2 Mole Concept
  7. 2.3 Concentration and Molarity
  8. 2.4 Gas Laws
  9. Module 3: Reactive Chemistry
    3.1 Chemical Reactions
  10. 3.2 Predicting Reactions of Metals
  11. 3.3 Rates of Reactions
  12. Module 4: Drivers of Reactions
    4.1 Energy Changes in Chemical Reactions
  13. 4.2 Enthalpy and Hess's Law
  14. 4.3 Entropy and Gibbs Free Energy
  15. Module 5: Equilibrium and Acid Reactions
    5.1 Static and Dynamic Equilibrium
    5 Topics
  16. 5.2 Factors that Affect Equilibrium
    2 Topics
  17. 5.3 Calculating the Equilibrium Constant
    2 Topics
  18. 5.4 Solution Equilibria
  19. Module 6: Acid/Base Reactions
    6.1 Properties of Acids and Bases
    7 Topics
  20. 6.2 Using Brønsted–Lowry Theory
    2 Topics
  21. 6.3 Quantitative Analysis
    1 Topic
  22. Module 7: Organic Chemistry
    7.1 Nomenclature
    2 Topics
  23. 7.2 Hydrocarbons
    2 Topics
  24. 7.3 Products of Reactions Involving Hydrocarbons
  25. 7.4 Alcohols
    1 Topic
  26. 7.5 Reactions of Organic Acids and Bases
  27. 7.6 Polymers
    2 Topics
  28. Module 8: Applying Chemical Ideas
    8.1 Analysis of Inorganic Substances
    3 Topics
  29. 8.2 Analysis of Organic Substances
  30. 8.3 Chemical Synthesis and Design
  31. Working Scientifically
    Working Scientifically Overview
    1 Topic
Lesson 19, Topic 3
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Acid-Base Models and Theories

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Antoine Lavoisier (1776) – the oxygen theory

  • Lavoisier demonstrated that non-metallic oxides (e.g. sulfur trioxide) dissolved in water made the water acidic.
  • Lavoisier proposed that oxygen was the cause of acidity. Lavoisier’s theory could not explain why metal oxides were not acidic.

Humphry Davy (1810) – the hydrogen theory

  • Davy observed that when HCl gas and H2S gas dissolved in water they produced acidic solutions. These molecules contained no oxygen atoms. Davy’s theory attributed acidity to the presence of replaceable hydrogen. These results were not consistent with Lavoisier’s oxygen theory.
  • Davy’s theory could not explain why not all hydrogen compounds were acidic.

Svante Arrhenius (1884) – the hydrogen ion theory

  • Arrhenius observed that acidic solutions conducted electricity just like salt solutions. He proposed that acids released hydrogen ions when they dissolved in water.
  • Arrhenius proposed that alkalinity was due to the presence of hydrogen ions in solution.
  • Arrhenius explained how acidic solutions neutralised alkaline solutions in terms of the interaction of the aqueous hydrogen ions and hydroxide ions:
\text{OH}^- (aq) + \text H^+ (aq) \rightarrow \text H_2 \text O (l)
  • Arrhenius’ theory also explained the difference between strong acids such as hydrochloric acid (HCl) and weak acids, such as hydrofluoric acid (HF) as being due to their different degrees of dissociation in water.

Johannes Brønsted and Thomas Lowry (1923)—the proton theory

  • Hydrogen ions are protons. These protons associate with water molecules to form hydronium ions (H3O+). B/L acids donate protons and B/L bases accept protons.
  • Bases include substances that neutralise acids but which are not classified as alkalis as no OH ions are present. In the following equation the hydronium ion donates a proton to the carbonate ion:
\text H_3\text O^+ (aq) + \text{CO}_3^{2-} (aq) \rightarrow \text H_2\text O(l) + \text{HCO}_3^- (aq)

Gilbert Lewis (1923) – the electron pair theory

  • A Lewis acid accepts an electron pair and a Lewis base donates an electron pair.
  • When hydrogen ions and hydroxide ions undergo neutralisation, the H+ ion accepts a pair of non-bonding electrons from the oxygen atom of the OH ion to form a covalent bond in the resulting water molecule.