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HSC Chemistry
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Module 1: Properties and Structure of Matter1.1 Properties of Matter
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1.2 Atomic Structure and Atomic Mass
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1.3 Periodicity
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1.4 Bonding
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Module 2: Introduction to Quantitative Chemistry2.1 Chemical Reactions and Stoichiometry
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2.2 Mole Concept
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2.3 Concentration and Molarity
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2.4 Gas Laws
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Module 3: Reactive Chemistry3.1 Chemical Reactions
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3.2 Predicting Reactions of Metals
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3.3 Rates of Reactions
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Module 4: Drivers of Reactions4.1 Energy Changes in Chemical Reactions
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4.2 Enthalpy and Hess's Law
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4.3 Entropy and Gibbs Free Energy
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Module 5: Equilibrium and Acid Reactions5.1 Static and Dynamic Equilibrium5 Topics
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5.2 Factors that Affect Equilibrium2 Topics
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5.3 Calculating the Equilibrium Constant2 Topics
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5.4 Solution Equilibria
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Module 6: Acid/Base Reactions6.1 Properties of Acids and Bases7 Topics
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6.2 Using Brønsted–Lowry Theory2 Topics
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6.3 Quantitative Analysis1 Topic
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Module 7: Organic Chemistry7.1 Nomenclature2 Topics
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7.2 Hydrocarbons2 Topics
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7.3 Products of Reactions Involving Hydrocarbons
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7.4 Alcohols1 Topic
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7.5 Reactions of Organic Acids and Bases
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7.6 Polymers2 Topics
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Module 8: Applying Chemical Ideas8.1 Analysis of Inorganic Substances3 Topics
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8.2 Analysis of Organic Substances
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8.3 Chemical Synthesis and Design
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Working ScientificallyWorking Scientifically Overview1 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.