# Are all solutions of weak acid/bases buffers?

I am having a difficult time understanding what makes a buffer a buffer.

Buffers in my textbook are defined as a solution of a weak acid or base and their conjugate acid/base. So if I were to just dissolve acetic acid (a weak acid) in water why would this not be defined as a buffer. Acetic acid incompletely dissociates into an acetate ion (the conjugate base) and a Hydrogen ion. Why is it that for a base to be made the acetic acid would have to be mixed with something like Sodium acetate if acetic acid incompletely dissociates into a conjugate base in the first place? Could someone please help me understand why this is the case?

That initial steep increase in the pH slows down quickly, and a pH plateau appears in the range of 4.0-5.8. As in this range $[\ce{CH3COOH}]$ is comparable to $[\ce{CH3COO-}]$ (i.e. their ratio is within a tenfold range), we will have a two-component system containing both a weak acid and a weak base—i.e. a buffer. The centre of the plateau is at 50 % titration where $[\ce{CH3COOH}]=[\ce{CH3COO-}]$. Therefore, according to the equation $\mathrm{pH}=\mathrm{p}K_\mathrm{a}+\log{\frac{[\ce{CH3COO-}]}{[\ce{CH3COOH}]}}$, the $\mathrm{pH}$ will just equal the $\mathrm{p}K_\mathrm{a}$. Following the plateau (upon the exhaustion of the buffer capacity), the pH will again increase steeply, and the top of the titration step will be reached, i.e. the point of neutralization at which the degree of titration of the acetic acid will be 100 %
From the figure above, the region marked inside the rectangle is a buffer. Because it can provide the most efficient pH stabilization against both acidic and basic shifts. In this region, the ratio of the two components (acid and conjugated base) is one or close to one. Accordingly, a buffer is effective in the $\mathrm{pH}$ range of its $\mathrm{p}K_\mathrm{a} \pm 1$ $\mathrm{pH}$ unit