# Confusion regarding the mechanism of acidic/basic buffer solutions

While learning about acidic and basic buffer solutions we were told that they can resist pH change if we add small amounts of acid or base by neutralizing the $$\ce{H+}$$/$$\ce{OH-}$$ ions from the added acid or base. But I am confused about the fact that if we add a small amount of base to an acidic buffer then after neutralization of the excess $$\ce{OH-}$$ ions from the base by the $$\ce{H+}$$ ions from the buffer, shouldn't the concentration of $$\ce{H+}$$ ions in the solution decrease, thereby making the solution acidic and reducing the $$\mathrm{pH}$$. (vice-versa for the addition of acid to a basic buffer)

• Buffers resisting pH change does not mean pH is constant. It means pH changes a lot less than without them being present. See the Henderson-Hasselbalch ( not -bach) equation on Libretexts or on Wikipedia and also buffer solution. It was question of 1-2 minutes of search and copy/paste. You may want to pay more attention to searching before asking. Feb 17 '21 at 7:21

While learning about acidic and basic buffer solutions we were told that they can resist pH change if we add small amounts of acid or base by neutralizing the $$\ce{H+/OH−}$$ ions from the added acid or base.

Yes, the pH changes less when you add small amounts of acid or base to a buffer, and more if you add it to pure water or an unbuffered aqueous solution.

But I am confused about the fact that if we add a small amount of base to an acidic buffer then after neutralization of the excess $$\ce{OH−}$$ ions from the base by the H+ ions from the buffer, shouldn't the concentration of $$\ce{H+}$$ ions in the solution decrease,

Yes, that is what happens. As a consequence, the weak acid/base reaction of the buffer substances is no longer at equilibrium, and there will be a net reaction to replenish some of the $$\ce{H+}$$ that was lost.

thereby making the solution acidic and reducing the pH. (vice-versa for the addition of acid to a basic buffer)

This is probably a typo. When you add a small amount of base, the pH would go up, buffered or not. In the presence of buffer, it goes up a little, and in the absence of buffer, a lot.

(vice-versa for the addition of acid to a basic buffer)

Buffers stabilize the pH near the $$\mathrm{p}K_\mathrm{a}$$ of the conjugate weak acid in "both directions", so even if the pH of a solution is slightly acidic, it is able to buffer against addition of acid or base.

This is linked to the fact that in a weak acid-base pair, the respective concentrations are linked to each other through their ratio. We have

$$K_\mathrm{a} = \frac{[\ce{A-}][\ce{H3O+}]}{[\ce{AH}]},$$

therefore

$$\mathrm{pH} = \mathrm{p}K_\mathrm{a} + \log\frac{[\ce{A-}]}{[\ce{HA}]}.$$

In water, the reaction is written

$$\ce{AH + H2O -> A- + H3O+}$$

If we add a little strong acid, then we shift the equilibrium a little in the direction of formation of the acid but as the concentrations are related to each other by the $$K_\mathrm{a},$$ the consequence on the $$\mathrm{pH}$$ is attenuated.

For a strong acid, the addition of another strong acid contributes totally to the increase in $$\mathrm{pH}$$ because there is no balance that counterbalances the addition of the acid (as is the case for a weak acid).