The plot of pH-metric titration of weak base $$NH_4OH$$ vs strong acid $$HCl$$ looks like

In a pH-metric titration of a weak base $$\mathrm{(NH_4OH)}$$ against a strong acid $$\mathrm{(HCl)}$$, the titration curve begins at a basic pH.

Since $$\mathrm{NH_4OH}$$ is a weak base, the initial pH is moderately basic and not extremely high.

$$\mathrm{Initial\ pH \approx 11}$$

As $$\mathrm{HCl}$$ is gradually added, the weak base gets neutralised:

$$\mathrm{NH_4OH + HCl \longrightarrow NH_4Cl + H_2O}$$

A buffer solution containing:

$$\mathrm{NH_4OH \ and\ NH_4^+}$$

is formed.

Because buffer solutions resist sudden pH changes, the curve shows a gradual downward slope in this region.

At the equivalence point:

$$\mathrm{NH_4Cl}$$

is present in solution.

Since this salt is formed from a strong acid and a weak base, it undergoes cationic hydrolysis:

$$\mathrm{NH_4^+ + H_2O \rightleftharpoons NH_4OH + H^+}$$

This produces excess $$\mathrm{H^+}$$ ions, making the solution acidic.

Therefore, the equivalence point lies below $$\mathrm{pH = 7}$$.

$$\mathrm{Equivalence\ pH \approx 5\ to\ 5.5}$$

Beyond the equivalence point, excess $$\mathrm{HCl}$$ dominates the solution.

Hence, the pH decreases sharply and finally becomes strongly acidic.

$$\mathrm{Final\ pH \approx 1\ to\ 2}$$

This corresponds to option A. Thus, A is the right option.