Polyprotic Acids, Polybasic Bases and Buffers.

.

.

.

.

.

Polyprotic Acids, Polybasic Bases and Buffers.

Have more than one hydrogen which can be removed in sequence on titration with a base and form more than one type of conjugate base as a result.

H SO  is a biprotic acid with two hydrogens and forms the two conjugate bases, Bisulphate HSO     and sulphate SO

H PO (phosphoric acid) is a triprotic acid. Each hydrogen reacts in turn giving three different conjugate bases as in the following curve:

Three different equilibria form at different pH's

The formation of H PO  is complete at about pH 4 so methyl orange could be used to detect this end-point.

The formation of HPO  is complete at about pH 9 so phenolphthalein could be used to detect this end point.

The third end point is very difficult to determine as                                  which is close to the value of K

This means that PO   is competing with OH- from water for H+ ions and the reversible equilibria means that there is never a pH value where only PO    is present in the solution.

Able to acccept more than 1 H+ from an acid, at different stages of the reaction.

In the carbonate-bicarbonate equilibrium 25cm  of 0.01mol dm  sodium carbonate solution titrated with 0.01 mol dm  hydrochloric acid solution would give the following graph:

There are clearly two possible end points at pH 9  and pH 4.

The carbonate ion, CO   must accept H+ first to form bicarbonate ion, HCO

The bicarbonate then reacts to give carbonic acid, H CO  as the endpoint of the titration.

The following equilibria are important in the transfer of CO  from cells in the body to the lungs via the blood.

Buffer solutions are useful in many chemical reactions as they allow the addition of a small amount of acid/ base without changing the pH.

They normally consist of either a weak acid and its salt or a weak base and its salt.

Ammonium acetate buffer is a mixture of ammonium acetate and acetic acid.

.

.

Adding H+ or OH- will shift the (top) ethanoic acid equilibrium to the left or right respectively (le chatelier's principle)

When acid is added to the buffer it reacts with the weak base present (CH COO )

When base is added to the buffer it reacts with the weak acid present (CH COOH)

The flat portion before the equivalence point is called the buffer region. In this part of the pH scale the acid and conjugate base are both present in significant concentrations and the solution resists changes in pH.

As base is added to a solution in this buffer region, the acid reacts with it to form ethanoate ion, without a large change in pH. If additional acid was added to the solution in the buffer region it would react with the conjugate base, ethanoate ion, and again the pH would not change significantly.

This point lies in the middle of the buffer region. At this point the concentrations of ethanoic acid and ethanoate ion are equal as the quantity of base added is half that required to reach the equivalence point.

The equilibrium expression at the half equivalence point is:

.

.

At the half equivalence point,                           so,

Therefore it follows that at this point,

This equation allows the pH of a buffer solution to be calculated when the concentrations of acid and conjugate base in it are not the same:

.

.

.

e.g. for the ammonium acetate buffer

- [acid] is the the concentration of ethanoic acid which is only partially dissociated.

- [base] is the concentration of acetate (ethanoate) ion obtained from the ammonium acetate salt.  The salt is completely dissociated so this value should be the same as the starting concentration of the salt.