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Equilibria involving acids and bases

Equilibria involving acids and bases. Chapter 17. Acidity of Solutions. Lowry-Brønsted theory Acids are proton donors Bases are proton acceptors. An acid base reaction involves the transfer of a proton from an acid to a base.

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Equilibria involving acids and bases

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  1. Equilibria involving acids and bases Chapter 17

  2. Acidity of Solutions • Lowry-Brønsted theory • Acids are proton donors • Bases are proton acceptors. • An acid base reaction involves the transfer of a proton from an acid to a base. • Some substances can both donate and accept protons, these are amphiprotic solutions.

  3. The ionisation constant of water • Water will react with itself in a process called self-ionisation: H2O(l) + H2O(l) H3O+(aq) + OH-(aq) • Write the equilibrium constant for this reaction. • The concentration of water is usually fairly constant so we can write the equilibrium law the ionisation of water as: • Kw = [H3O+][OH-] • Where Kw is the ionisation of water

  4. The ionisation constant of water • This law applies to both pure water and all aqueous solutions. • In pure water at 25°C, chemists have found that the concentration of both H3O+ and OH- ions is 10-7M. • The value of Kw at 25°C can be calculated: • Kw = [H3O+][OH-] = 10-7 x 10-7 = 1.0 x 10-14 M2

  5. Acidic and Basic Solutions • In acidic solutions the concentration of H3O+ ions will be greater than 10-7 M at 25°C. • The opposite is true for basic solutions. • In summary at 25°C: • In pure water and neutral solutions: [H3O+] = [OH-] = 10-7 M • In acidic solutions: [H3O+] > 10-7 M and [OH-] < 10-7 M • In basic solutions: [H3O+] < 10-7 M and [OH-] > 10-7 M • The higher the concentration of hydronium ions in a solution the more acid the solutions is.

  6. pH • pH = -log10[H3O+] • [H3O+] = 10-pH • Remember for basic solutions we are given the OH- concentration • Can anyone remember how we can convert this to a H3O+ concentration? • Lets look at the worked examples on page 287

  7. How is pH affected by temperature? Remember: Kw = [H3O+][OH-] = 10-7 x 10-7 = 1.0 x 10-14 M2 at 25°C • But what happens if the temperature is not 25°C • Ionisation of water is endothermic • So: • As the temperature increases the equilibrium constant increases • This causes Kw to rise which results in a decrease in pH • If the temperature decreases it favours a reverse reaction which is exothermic • Kw will decrease and the pH will increase

  8. Your Turn • Page 288 • Questions 1 - 3

  9. Acidity Constants • For the reaction • HCl(aq) + H2O(l) H3O+(aq) + Cl-(aq) • The expression for the equilibrium constant can be written as: • Water is the solvent in aqueous solutions and its concentration is virtually constant. • So we can remove water from the equilibrium constant and call it Ka instead [H3O+][Cl-] K = [HCl][H2O]

  10. Acidity Constant [H3O+][Cl-] • Ka is the acidity constant. • The value of Ka for HCl is 107 M at 25°C. • This means that in HCl solutions, most of the HCl has been converted to product. • This is why HCl is classified as a strong acid. Ka = [HCl]

  11. Acidity Constants • These ideas can be generalised to solutions of any weak acid represented by HA: • HA(aq) + H2O(l) A-(aq) + H3O+(aq) • [H3O+] = [A-] and [HA] does not change during the ionisation

  12. Calculations Involving Acidity Constants • Calculate the pH and percentage hydrolysis of a 0.50 M ethanoic acid solution, given that the Ka for ethanoic acid is 1.75 x 10-5 M

  13. Your Turn • Page 292 • Question 4 and 5

  14. Buffers: Using equilibrium to resist change • Buffers are solutions that can absorb the addition of acids or bases with little change of pH. • They can be made most easily by mixing a weak acid and a salt of its conjugate base. • Buffers are especially important in environmental and living systems where they maintain a delicate chemical balances essential to life. • Consider the equation for ethanoic acid in water as a buffer. • What happens to the equilibrium when you add a strong acid or a strong base???

  15. pH in the body • A number of reactions that occur in the body involve acid-base reactions. • Without a means of controlling acidity the pH of body fluids could fluctuate from extremely basic to extremely acid. • Our bodily systems can only operate in a small range. • Turn to page 292 for an example.

  16. Your Turn • Page 292 • Question 6 • End of Chapter questions are due

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