Buffers resist changes in pH because they contain a weak acid and its conjugate base (or a weak base and its conjugate acid). When a small amount of acid is added, the conjugate base neutralizes the extra H+, helping keep the pH stable. When a small amount of base is added, the weak acid donates protons to neutralize it, again helping to stabilize the pH. This balance is often described by pH ≈ pKa + log([A-]/[HA]), meaning the pH stays near the acid’s pKa as long as both species are present. A common example is a solution containing a weak acid and its conjugate base, like acetic acid and acetate, which buffers around the pH near 4.8. In contrast, a solution that cannot neutralize added acids or bases won’t resist pH changes; a strong acid and strong base in equal amounts simply neutralize to water and salt and don’t provide ongoing buffering; and a solution with a high concentration of H3O+ is strongly acidic and lacks buffering behavior.

Under Arrhenius theory, bases are substances that increase the concentration of hydroxide ions in an aqueous solution. When a base dissolves, it releases OH- into the water, as in NaOH ⇌ Na+ + OH-. Even bases like ammonia, which don’t contain OH- themselves, raise OH- levels by reacting with water: NH3 + H2O ⇌ NH4+ + OH-. This increase in hydroxide ions is what makes the solution basic. Hydronium ions (H3O+) are produced by acids, not bases, in water, and free H+ isn’t present as a separate species in solution. So the ion yielded by a base in water is hydroxide, OH-.

When a strong acid and a strong base are mixed in equal amounts, they neutralize to form a salt and water. The H+ from the acid combines with the OH- from the base to make H2O, while the remaining ions—Na+ and Cl− in this case—pair up to form the salt NaCl. Because the volumes and concentrations are the same, the reaction goes to completion with no excess acid or base left. So the resulting solution contains the salt NaCl (and water as the other product). The other substances listed are either consumed in the reaction (HCl and NaOH) or are simply water, which is not a salt.

pH measures how acidic or basic a solution is by looking at the hydrogen ion concentration. Pure water undergoes a tiny self-ionization: H2O ⇌ H+ + OH−. At room temperature, the concentrations of H+ and OH− are equal, each about 1.0 × 10^-7 M. The pH is the negative logarithm of the hydrogen ion concentration, so pH = -log10(1.0 × 10^-7) = 7. This makes pure water neutral on the pH scale. Temperature can shift this neutrality a bit because the water ionization constant Kw changes with temperature. Values like 1 or 4 would mean a lot more H+ (acidic), while 10 would mean more OH− (basic).

The main concept is the common-ion effect for a weak acid. Acetic acid dissociates as HA ⇌ H+ + A−. When you add a salt that provides the conjugate base A−, you increase [A−] in the solution. Le Chatelier’s principle says the system will shift to counter this stress, moving the equilibrium toward the left to form more HA and reduce the amount of H+. So, the dissociation is suppressed and [H+] decreases (pH rises). Since acetic acid is weak, it does not fully ionize, and adding its conjugate base won’t cause complete ionization. The other options don’t fit because increasing dissociation isn’t happening, the equilibrium does change, and complete ionization isn’t achieved by this addition.