Acid-base equilibria

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I'm struggling to get my head around buffers. I understand the basic principles of buffers but I don't understand how to work out how the two equilibriums work together.
For example, one exam question is:
 
"Mixtures of ammonia, NH₃ and ammonium chloride, NH₄Cl, can be used as a basis of buffer solutions. Explain how this mixture acts as a buffer solution."
 
The two equilibria are:
 
NH₄⁺ --> NH₃ + H⁺
 
NH₃ + H₂O --> NH₄⁺ + OH^-
 
Please can you explain how adding acid and alkali effects the equilibrium and why.
060608
 

Corrie writes ....
 
A buffer usually consists of a mixture of a weak acid and its conjugate base, or vice versa. They are in equilibrium in solution, just as the weak acid and its base would be if you just dissolved the weak acid in water.
 
In the ammonia buffer system you mention, the NH₄⁺ ion (from NH₄Cl) is the weak acid and NH₃ is the base. The buffer equilibrium (only one!) involved is therefore:
 
NH₄⁺(aq) <=> NH₃(aq) + H⁺(aq)
 
(You should always write down the equilibrium in an answer about buffers, even if it is only AH <=> B + H⁺ for a general case involving acid, AH, and base, B. Then you have something to refer to in your answer.)
 
Now back to your ammonia-based buffer solution, in which the equilibrium above exists, with lots of NH₄⁺ ions present and lots of NH₃ present.
 
1. Adding some strong acid to the buffer solution:
This would increase the concentration of H⁺ ions in the buffer solution, lowering the pH. The disturbed equilibrium will therefore move to the left to lower the H⁺ concentration, until the equilibrium is re-established, with the pH back to (almost) its original value.
 
2. Adding some strong alkali is a little more complicated:
If some strong alkali (i.e. OH^- ions) is added to the buffer solution, the first thing that happens is that the OH^- ions react with the H⁺ ions present to form water:
 
OH^- + H⁺ -> H₂O
 
Because of the reduction in the H⁺ ion concentration that results, the buffer equilbrium moves to the right to produce more H⁺ ions and restore the equilibrium. This continues until all the alkali has reacted and the pH is restored to its original value.
 
Obviously there is a limit to the amount of strong acid or alkali a buffer can 'cope' with, so it is usually assumed that the amount of strong acid or base added is small compared to the amount of buffer substances present, so that the pH of the buffer does not change significantly.

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updated: 18 June 2008