Acid-base equilibria
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I am having problems calculating [H+] for solutions of acids and bases. The examples and explanations on pages 336–7 of the Nuffield students' book don't seem very clear to me.
The pages you refer to explain how you can us the equilibrium constant, Kw for the ionisation of water to calculate the value of [OH-] in solutions of strong acids or the value of [H+] in solutions of strong bases.
With a strong acid we assume it is completely ionised. Strong acids are HCl, HNO3, and H2SO4.
Note that sulphuric acid ionises mainly to H+ and HSO4-. HSO4- is a weak acid and does not ionise to any great extent in aqueous solution.
One mole of the strong acid HCl produces one mole of H+ ion when it dissolves in water, so the H+ concentration is assumed to be the same as the concentration of the acid. The same is true for HNO3.
NaOH or KOH are strong bases so we can assume complete ionisation. Each mole of NaOH produces one mole OH– ion when it dissolves, so the OH– concentration is the same as the NaOH concentration.
If we know the OH– concentration, we can find the H+ concentration by using Kw (see page 336). For example. in 0.1M NaOH the value of [OH–] is 0.1 mol dm–3.
Now Kw = [H+] x [OH–] = 1 x 10–14 mol2 dm–6.
So [H+] = Kw divided by [OH–] which is 1 x 10 –13 mol dm–3.
When you are dealing with a weak acid, ionisation is incomplete and often very slight. You need to use the dissociation constant of the acid. In an exam this would be given, but there is a full list in the Nuffield Book of data, table 6.5 on page 122. The method of calculating the H+ concentration for a weak acid is set out on page 339 of the Students Book and I don't think I can improve on it!
Jenkin
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updated: 21 August 2003
