21 – Buffers and neutralization

A system that minimises pH changes when small amounts of an acid or a base are added

A weak acid and its conjugate base
– weak acid (HA) removes added alkali
– conjugate base (A-) removes added acid

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When does a buffer lose its buffering ability?

As soon as one component has reacted, the solution loses its buffering ability towards added acids and alkali.
As the buffer works, the pH does change a small amount – you should not assume that the pH stays completely constant.

Preparing weak acid buffer solutions using a weak acid and its salt.

E.g. Weak acid – CH3COOH, salt – CH3COONa

Ethanoic acid partially dissociates and the amount of ethanoate ions in the solution is very small.

Salt completely dissolves and is a good source of the conjugate base, CH3COO-, which helps to increase the concentration of CH3COO- so that more H+ can be mopped up.

Preparing weak acid buffer solutions by partial neutralisation of the weak acid

Adding an aqueous solution of alkali to an excess of weak acid. The weak acid is partially neutralised by the alkali, forming the conjugate base. Some of the weak acid is left over unreacted. The resulting solution contains a mixture of the salt of the weak acid and the unreacted weak acid.

Conjugate base removes added acid

[H+] increases, which causes more H+ to react with A- . The equilibrium position shifts to the left and H+ is used up.

Weak acid removes added alkali

[OH-] increase, so H+ + OH- –> H2O.
[H+] decreases, so equilibrium position shifts to the right to restore the concentration of H+

Choosing the components of a buffer solution

When [HA] = [A-]:
– the pH of the buffer solution is equal to the pKa of HA
– the operating is typically over about two pH units, centred at the pH of the pKa value.

Calculating the pH of a buffer solution

[H+] = Ka x [HA]/[A-] = Ka x [weak acid]/[salt of weak acid]

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What pH is blood plasma?

pH needs to be maintained between 7.35 and 7.45
Normal healthy blood should have a pH of 7.4
pH of blood is maintained by a mixture of buffers, with carbonic acid- hydrogen carbonate being the most important.

What happens if blood pH slips outside optimal range?

Acidosis – falls below 7.35, causes shortness of breath, fatigue and even shock or death.

Alkalosis – above 7.45, causes nausea, light-headedness and muscle spasms.

* A difference of just O.3pH units is a two-fold difference in H+ concentration

The carbonic acid-hydrogen carbonate ion buffer system

How does the body prevent H2CO3 build up?

By converting the H2CO3 into carbon dioxide gas, which is exhaled by the lungs.

Henderson-Hasselbalch equation (alternative way of calculating pH of a buffer solution)

pH = pKa + log [A-]/[HA]

pH titration curves

a plot of the pH of the solution as a function of the volume of added titrant

Key features of a pH titration curve

equivalence point of a titration

the volume of one solution that exactly reacts with the volume of another solution. The equivalence point is the centre of the vertical section of the pH curve.

4 different types of titration curve

end point of titration

the point at which the indicator changes colour
The end point does not always equal the equivalence point.

How sensitive is the end point?

Different indicators have different Ka values and change colour over different pH ranges.
At the end point:
– [HA] = [A-]
– Ka = H+ , pH = pKa
The pH of the end point is the same as the pKa for HA.
Sensitivity of an indicator depends on the indicator itself and eyesight.

Indicators and their pH ranges

Use an indicator that has a colour range that coincides with the vertical section of the pH titration curve. Ideally the end point and equivalence point would coincide, although this may not be possible as the end point may give a slightly different volume to the equivalence point. Any difference will be very small (1 or 2 drops)