Q.iii. How are basic buffer solutions prepared?

A buffer solution resists drastic pH changes when a small amount of strong acid, strong base, or water is added to it. An acidic buffer pairs a weak acid with its salt of a strong base (e.g. CH3COOH + CH3COONa); its pH is given by the Henderson-Hasselbalch equation $pH=pK_a+\log_{10}\dfrac{[salt]}{[acid]}$, where $pK_a=-\log_{10}K_a$. A basic buffer pairs a weak base with its salt of a strong acid (e.g. NH4OH + NH4Cl); its pOH is $pOH=pK_b+\log_{10}\dfrac{[salt]}{[base]}$, where $pK_b=-\log_{10}K_b$. The mechanism (buffer action) relies on having LARGE reservoirs of both the weak acid/base and its conjugate salt simultaneously present: any added strong acid's H+ is consumed by the abundant conjugate base, and any added strong base's OH- is neutralized by the abundant weak acid/base itself, so neither addition meaningfully shifts [H+] or [OH-]. Because dilution scales both reservoir concentrations by the same factor, the [salt]/[acid] ratio -- and hence the pH -- is also unaffected by dilution. These properties (resisting small acid/base additions, dilution, and standing over time) are what make buffers essential in practice: blood pH is held at 7.36-7.42 by an (HCO3-+H2CO3) buffer, soils are naturally buffered by carbonate/phosphate/organic-acid salts, and buffers stabilize products in the paper, dye, paint and pharmaceutical industries (e.g. sodium citrate stabilizing penicillin) as well as in qualitative analytical chemistry (an NH4OH+NH4Cl buffer maintains the pH 8-10 needed to selectively precipitate group IIIA cations).