The relationship between $\mathrm{pK_a}$ and $\mathrm{pH}$ of a weak acid, $\ce {HA}$, is mathematically represented by Henderson-Hasselbach equation where $\ce {A-}$ is its conjugate base (there is a equation for a weak base as well).
$$\mathrm{pH = pK_a + log\frac{[\ce {A-}]}{[\ce {HA}]}}$$
Thus, the $\mathrm{pH}$ of a buffer of weak acid can be calculated using above equation, which gives close approximate value. As @MollyCooL pointed out, the highest buffer capacity is in effect when $\mathrm{pH}$ of the buffer is equal to $\mathrm{pK_a}$ of the acid, meaning,
$\mathrm{log\frac{[\ce {A-}]}{[\ce {HA}]}= pH - pK_a =0}$. Hence, $\mathrm{\frac{[\ce {A-}]}{[\ce {HA}]}= 1}$ or $\mathrm{[\ce {A-}]=[\ce {HA}]}$. At this $\mathrm{pH}$, the effective buffer capacity is $\mathrm{pK_a \pm 1}$. Because of this reason, it is common practice with biochemists that they made bufer solutions at or about $\mathrm{pK_a}$ of the relavent acid or base. For example, if they need buffer of $\mathrm{pH}\space 4.5$, they use acetic acid/sodium acetate (acetic acid: $\mathrm{pK_a} = 4.76$), but use $\ce {KH2PO4/K2HPO4}$ system to make $\mathrm{pH}\space 7.0$ buffer (phosphoric acid: $\mathrm{pK_{a2}} = 7.2$). To understand this concept, see Plot 1 below:
This plot illustrates various protonation states of acetic acid as the pH changes. When $\mathrm{pH}$ is below $\mathrm{pK_a}$ of the acid ($\mathrm{pH \lt 3.5}$), most of the acid has been protonated. When $\mathrm{pH}$ is significantly above $\mathrm{pK_a}$ ($\mathrm{pH \gt 6.0}$), most of the acid is deprotonated. If the $\mathrm{pH}$ close of equal to the $\mathrm{pK_a}$, the acid is about $\mathrm 50\%$ protonated and $\mathrm 50\%$ deprotonated. That means, acetic acid is not suitable for the situations where sought $\mathrm{pH}$ is below $\mathrm 3.5$ or above $\mathrm 6.0$.
Now, let's do little quick and dirty calculations for your buffer solution:
$$\pu {Final acetic acid concentration = 9.100 \space M\times\frac{0.025 \space L}{0.075 \space L} = 3.033 \space M}$$
$$\pu {Final acetate concentration = 0.100 \space M\times\frac{0.050 \space L}{0.075 \space L} = 0.067 \space M}$$
Now, apply these two values to Henderson-Hasselbach equation:
$$\mathrm{pH = 4.76 + log\frac{[\ce {0.067}]}{[\ce {3.033}]}= 4.76-1.66=3.1}$$
Thus, the buffer system selected is not good choice for $\mathrm {pH} \space 3.1$ buffer. ascorbic acid (ascorbic acid: $\mathrm{pK_{a1}} = 4.10$) or phosphoric acid (phosphoric acid: $\mathrm{pK_{a1}} = 2.15$) would have been a better choice. Nonetheless, citric acid (citric acid: $\mathrm{pK_{a1}} = 3.15$) is the ideal one for this $\mathrm {pH}$.
The plot is from: https://bio.libretexts.org/LibreTexts/University_of_California_Davis/BIS_2A%3A_Introductory_Biology_(Easlon)/Readings/06.1%3A_pKa