Chemistry: Buffers and Henderson-Hasselbalch

Understanding Chemistry: Buffers and Henderson-Hasselbalch

The resistance comes from a fast chemical adjustment, not from stopping added substances from entering the solution. A weak acid exists in balance with its dissolved hydrogen ions and its conjugate base. When extra hydrogen ions arrive, the conjugate base joins with many of them to form more weak acid.

When hydroxide ions arrive, the weak acid supplies hydrogen ions that turn hydroxide into water. The balance shifts in the needed direction.

This is an example of Le Chatelier's principle. The pH still changes a little because the chemical balance is never unlimited or perfectly fixed.

The Henderson-Hasselbalch equation connects pH to the acid strength and to the relative amounts of the two buffer forms. It comes from rearranging the expression for the acid dissociation constant. The equation says pH equals pKa plus the logarithm of the concentration of conjugate base divided by the concentration of weak acid.

The logarithm matters because pH is a logarithmic scale. A tenfold change in the ratio changes the predicted pH by one unit. A larger amount of conjugate base gives a higher pH.

A larger amount of weak acid gives a lower pH. In typical school calculations, concentration is used as a close estimate of chemical activity. This estimate becomes less accurate in very concentrated solutions.

Buffer capacity describes how much strong acid or strong base a buffer can handle before its pH changes greatly. Capacity depends on the total amount of buffer chemicals present. Two solutions can have the same pH and the same conjugate base to weak acid ratio, yet one can be far more useful as a buffer because it contains larger amounts of both forms.

A dilute buffer has the right ratio but runs out quickly. Dilution often causes little immediate pH change because the ratio stays nearly the same.

It does reduce capacity. Temperature can matter too, since acid dissociation constants can change with temperature.

When solving buffer problems, deal with the added strong acid or strong base first. This reaction is usually treated as complete. Use amounts in moles to find how much weak acid and conjugate base remain after that reaction.

Only then use their new ratio in the Henderson-Hasselbalch equation. Students often use the original concentrations by mistake. It is also important to choose a buffer whose pKa is near the required pH.

In blood, the carbonic acid and bicarbonate system helps keep pH near 7.4, while breathing and kidney function help control the components. In laboratories, buffers keep enzyme tests, DNA work, and indicator measurements within conditions where the results remain reliable.