QCM : Fundamentals of Acid-Base Chemistry — 20 questions

Explication

A Brønsted acid is defined as a species that donates a proton. A proton acceptor is a Brønsted base, not an acid.

Explication

Conjugate acid-base pairs differ by exactly one proton, with one member formed when the other donates or accepts H+. The other options do not capture this proton-transfer relationship.

Explication

For two conjugate pairs, K greater than 1 means the first acid is stronger than the second, and the second base is stronger than the first. This follows the competition for proton transfer.

Explication

In aqueous solution, any acid stronger than hydronium is converted completely to H3O+, so their apparent strengths are leveled. This is the leveling effect of water.

Explication

pKa is defined as -log10(Ka). A smaller pKa therefore corresponds to a stronger acid.

Explication

Because pKa decreases as acid strength increases, a smaller pKa means the acid is stronger and its conjugate base is weaker. This is the Golden Rule of strength.

Explication

Neutrality means [H3O+] equals [OH-], as set by water autoionization. The pH is not always 7 unless the temperature and equilibrium constant correspond to that value.

Explication

A basic solution has more OH- than H3O+, so its pH is above the neutral benchmark. The opposite describes an acidic solution.

Explication

Neglecting water is valid when the acid-generated hydronium is much larger than that from water, expressed as [H3O+]eq ≥ 10Ke. That keeps the pH error acceptably small.

Explication

For a weak acid, the approximation is acceptable only if the degree of dissociation alpha stays below 10%. This ensures the initial acid concentration is barely changed.

Explication

A strong acid is treated as completely dissociated in water, so hydronium is produced in large amount. The weak-acid behavior described in other options does not apply.

Explication

Water levels strong acids so any acid stronger than hydronium is converted into H3O+, making them look equally strong. The same leveling idea applies to strong bases and hydroxide.

Explication

Weak acids dissociate only partially, so an equilibrium is established in solution. Complete dissociation is a feature of strong acids, not weak acids.

Explication

Because pKa = -log10(Ka), a higher pKa means a smaller Ka and therefore a weaker acid. By the Golden Rule, that corresponds to a stronger conjugate base.

Explication

Strong acids are treated as completely dissociated, so their hydronium contributions add directly. Henderson-Hasselbalch applies to weak acid/conjugate base mixtures, not strong acids.

Explication

In a strong acid plus weak acid mixture, the strong acid controls the acidity because it contributes the dominant hydronium concentration. The weak acid is not the main determinant of pH.

Explication

For a 1:1 strong acid–strong base titration, equivalence corresponds to a neutral solution at 298 K, so pH is 7.00. The weak-acid buffer ideas do not apply here.

Explication

At half-equivalence, half the weak acid has been converted to conjugate base, so [AH] = [A−] and the logarithmic term in Henderson-Hasselbalch is zero. That makes pH = pKa.

Explication

A buffer resists major pH changes when small amounts of acid or base are added, which is the defining property. Equal hydronium and hydroxide is the definition of neutrality, not buffering.

Explication

Buffer efficiency peaks when [Base]/[Acid] = 1, which corresponds to pH = pKa. The useful buffer window is roughly [Base]/[Acid] between 0.1 and 10.