Question

Why do buffer systems use either a weak acid or a weak base?

Weak acids and weak bases do not react easily.
Weak acids and weak bases are partially dissociated.
Weak acids and weak bases are unaffected by temperature.
Weak acids and weak bases are found in abundance on the earth.

Answer 1

Weak acids and weak bases are partially dissociate

Step-by-step explanation:

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Answer 2

Weak acids and weak bases are only partially dissociated.

Step-by-step explanation:

When a small amount of an acid or a base is added to water (or a solution of some strong acid or base of some certain
`pH \!`,) the
`pH` of that solution tends to change significantly.

In contrast, when the same amount of acid or base is added to a buffer system of some certain
`pH\!`, the
`pH` of this system tends to change by a much smaller amount.

Weak acids and weak bases dissociate only partially. That allows the buffer solutions to contain a large reservoir of both the (undissociated) weak acid (or weak base) and the corresponding conjugate ion, while maintaining the required
`pH`.

This reservoir of the partially-dissociated weak acid (or weak base) and its conjugate ion allows the buffer system to absorb
`\rm H^(+)` and
`\rm OH^(-)` that were added to the solution without much change to the
`pH`.

For example, consider a
`1\; \rm L` buffer solution that initially included
`0.1\; \rm mol` of acetic acid
`\rm CH_3COOH` (a weak acid with
`pK_\text{a} = 4.765`) and
`0.1\; \rm mol\!` of sodium acetate
`\rm CH_3COONa`. The
`pH` of this solution would be
`4.765`.

In contrast, if a solution of
`\rm HCl` (a strong acid) and an equal amount of
`\rm NaCl` also needs to be
`4.765`, only
`10^(-4.765)\; \rm mol \approx 0.0000172\; \rm mol` of each species could be included.

Assume that
`0.001\; \rm mol` of
`{\rm NaOH}\, (s)` is added to the buffer solution of
`\rm CH_3COOH` and
`\rm CH_3COONa`. The quantity of the partially-dissociated weak acid in this solution is much larger than that of the
`\rm OH^(-)` added to the solution.

Only a small portion of the weak base would be consumed.The
`pH` of the solution would likely change by less than
`0.001`.

In contrast, assume that
`0.001\; \rm mol` of
`{\rm NaOH}\, (s)` is added to the
`1\; \rm L` solution with
`0.0000172\; \rm mol` of
`\rm HCl` (fully-dissociated) and
`\rm NaCl`.

All the
`\rm HCl\!` in this solution would be consumed. The
`pH` of this solution would go well above
`7`.