Acid-Base Equilibria

Acid-Base Equilibria

1. • Acids are proton (H+) donors • Bases are proton acceptors

2. Strong acids donate all their protons in aqueous solution, eg nitric acid: - HNO3 (aq) + H2O (l) → H3O+ (aq) + NO3- (aq)

   • Weak acids donate only some of their protons, ie they only partly dissociate into H+ ions and negative ions. The free ions form a dynamic equilibrium with the undissociated acid, eg ethanoic acid: - - + CH3COOH (aq) + H2O (l)  CH3COO (aq) + H3O (aq)

3. Water is accepting a proton, so is classed as a base. There is a second base in the equilibrium, which is CH3COO- (in the reverse direction). This brings about the idea of conjugate acids and bases: CH3COO- is the conjugate base of CH3COOH and H3O+ is the conjugate acid of H2O-

4. Ka = [H+][A-] [HA]

5. Note that it is called Ka, a = acid. • The stronger the acid, the more dissociated it is, giving more ions, so Ka will be a larger value. • A weak acid has less dissociation, so Ka will be smaller.

6. Pure water can conduct small amounts of electricity because it has ions dissolved in it. These ions are there because water dissociates into H+ and OH- ions, producing the following equilibrium: -Kc = [H+][OH-] [H2O], The amount of water that dissociates is tiny, so the equilibrium lies almost completely to the left. This means that the concentration of water can be thought of as constant and we can combine it into the equilibrium constant and rename it Kw, the ionic product of water, written as: -Kw = [H+][OH-].

7. pH for strong bases - When a strong base is dissolved in water, the OH- ions have the same concentration as the original base. If we want to find out the pH, the [H+] ion concentration needs to be known, even in a solution of a base. These will come from the water, so if we use the expression for Kw we can find out [H+] concentration: - Kw = [H+][OH-], so [H+] = Kw - [OH

8. A buffer is a solution whose pH stays almost constant if a small amount of acid or alkali is added to it. It removes any added H+ or OH- ions. It is usually made from a mixture of a weak acid and a salt of the same acid, eg CH3COOH and CH3COONa.

9. There are two dissociations that occur in an acidic buffer: - • The salt dissociates fully: - - + CH3COONa (aq) → CH3COO (aq) + Na (aq) • The acid partly dissociates and sets up an equilibrium: - CH3COOH (aq)  CH3COO- (aq) + H+ (aq), This equilibrium will be pushed far to the left as there is a high concentration of ethanoate ions from the salt, so the acid won’t dissociate very much. The concentration of H+ ions are increased as you add acid to the buffer, so + according to Le Chatelier, equilibrium will shift to the left. This removes the H ions as they react with the ethanoate ions. A base added to the buffer will react with the H+ ions. This shifts the equilibrium to the right and releases more H+ ions from the ethanoic acid..

10. A basic buffer can be made using a mixture of ammonium chloride and ammonia solution. The important equilibrium is: - NH4+  NH3 + H+ • Adding a base removes H+ ions, causing the equilibrium to move to the right, producing more H+. + • Adding an acid shifts the equilibrium to the left, removing the H ions from the acid.

11. Strong acid/strong base - NaOH is added slowly to HCl and you see the pH increasing very slowly until the volume of base added equals the volume of acid, then you see a very sharp increase in pH. This region of the curve is vertical, which rises until it levels off at the end. The middle of the vertical line is called the equivalence point, where amount of alkali added = amount of acid.

12. Weak acid/strong base - NaOH is added to CH3COOH and the pH rises slowly to around 4. When the volume of NaOH reaches half the volume of the acid, the pH levels off. A mixture is formed with unreacted acid and the salt of the acid and this acts as a buffer. pH will not change much then for a while. pH then increases vertically after pH 6.5 to around pH 12.

13. Strong acid/weak base - You’ll see that the pH up to 7 is the same as the first curve. The line then becomes vertical but levels off due to the buffer effect. The pH then increases up to a pH of around 12.

14. Weak acid/weak base - There is no clear vertical line in this graph, so it makes it harder to find an equivalence point. If you don’t have this, then it’s hard to choose an indicator for the reaction (see below). It’s better to use a pH probe to find the equivalence point.

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