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Historical Development of Acid/Base Theories

Historical Development of Acid/Base Theories. Year 12 Chemistry. Antoine Lavoisier (1743-1794). He classified all chemicals into three categories – acids, bases and salts He believed that all acids contained oxygen and it was this that gave them their sour taste

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Historical Development of Acid/Base Theories

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  1. Historical Development of Acid/Base Theories Year 12 Chemistry

  2. Antoine Lavoisier (1743-1794) • He classified all chemicals into three categories – acids, bases and salts • He believed that all acids contained oxygen and it was this that gave them their sour taste • Flaw: not all acids contain oxygen and metal oxides form bases

  3. Humphrey Davy (1778-1829) • Showed that all acids do not contain oxygen • Proposed that acids are hydrogen containing materials following the discovery of HCl • Flaw: not all substances that contain hydrogen are acids

  4. Svante Arrhenius (1859-1927) • Acids dissociate in water forming H+ as one product • Bases dissociate in water forming OH- as one product • Neutralisation involves the reaction of H+ and OH- forming a salt in water • Flaws: • theories only apply to aqueous solutions • Some substances such as NH3 are bases and do not contain OH- • Relative strengths not addressed • Amphoteric substances not addressed

  5. Bronsted-Lowry (~1923) • An acid is a proton (H+) donor • A base is a proton acceptor • Examples: • HCl + H2O <===> H3O+ + Cl¯ • NH3 + H2O <===> NH4+ + OH¯ Identify the acids/bases Any other acids/bases here?

  6. B-L Examples HCl + H2O  H3O+ + Cl- NH3 + H2SO4  NH4+ + HSO4- HBr + NH2+ NH3+ + Br- Identify the B-L acids and bases in each of these reactions.

  7. Amphiprotic Substances • Amphiprotic substances are those that can act as bases and acids. They can donate or accept protons. • Water is an obvious example Notice in the previous slide that water reacts with both acids and bases.

  8. Exercise Bisulfate is another amphiprotic substance. Construct chemical equations to show this property.

  9. Gilbert Lewis (~1920) Lewis Theory • An acid is an electron pair receptor • A base is an electron pair donor Note that BF3 would not be an acid under the B-L Theory as there is no H+ to donate.

  10. Other Lewis examples A cobalt metal complex has 6 dative bonds formed by the donation of e- pairs from the ammonia molecules (Lewis bases) Which of these substances are Lewis bases? Lewis acids?

  11. Exercise Water reacts with carbon dioxide to form carbonic acid. Write out the structures and show how the electrons are transferred, thereby identifying the Lewis acid and Lewis base.

  12. Bronsted-Lowry (conjugates) • Acid 1 + Base 2  Base 1 + Acid 2 • What does this mean? • An acid reacts and forms a conjugate base which can also accept a proton • A base reacts and forms a conjugate acid which can donate a proton • Conjugate pairs differ only by one H+ Conjugate pair 1 Conjugate pair 2

  13. Bronsted-Lowry (conjugates) • Example: • HNO3 + H2O <===> H3O+ + NO3¯ Identify the conjugate pairs in this reaction • Here, nitric acid and the nitrate ion are conjugates and water and the hydronium ion are also conjugates

  14. Strong vs. Weak Strong acids completely dissociate in water HA + H2O  H3O+ + A- In general, these reactions are reversible, but for a strong acid the equilibrium is far right. So, HA + H2O  H3O+ + A- Examples: HCl, HNO3, H2SO4 Weak acids Partially dissociate in water HA + H2O  H3O+ + A- There is an equilibrium established with weak acids, which means that there is less H3O+ ions in solution. CH3COOH + H2O  H3O+ + CH3COO- Examples: CH3COOH, H2CO3 What do you think is the difference in electrical conductivity?

  15. Strong vs. Weak Strong bases completely dissociate in water NaOH  Na+ + OH- Again, as with acids, the equilibrium is far right. So, 1 mol NaOH  1 mol OH- Examples: group I hydroxides, Ba(OH)2 Weak bases Partially dissociate in water NH3 + H2O  NH4++ OH- Only about 1% of ammonia dissociates into hydroxide ions. 1 mol NH3  << 1mol OH- Examples: NH3, other amines Again, the electrical conductivity is greater for strong bases

  16. Bronsted-Lowry (equilibrium) • Predicting Equilibrium • The direction of acid-base equilibria is away from the stronger acid base side and towards the weaker acid base side • The stronger the acid, the weaker its conjugate base • The stronger the base, the weaker its conjugate acid

  17. Bronsted-Lowry (equilibrium) • Reactions that proceed to a large extent: • A strong acid will force the equilibrium in the opposite direction (in this case, forward or right) • HCl + H2O <===> H3O+ + Cl¯ • Reactions that proceed to a small extent: • If the weaker of the two acids and the weaker of the two bases are reactants (appear on the left side of the equation), the reaction is said to proceed to only a small extent: • NH3 + H2O <===> NH4+ + OH¯ • Identify the conjugate acid base pairs in each reaction.

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