This question tests understanding of weak acid and base equilibria. Methylamine ($\text{CH}_3\text{NH}_2$) is a weak base with $K_b = 4.4 \times 10^{-4}$, establishing equilibrium $\text{CH}_3\text{NH}_2 + \text{H}_2\text{O} \rightleftharpoons \text{CH}_3\text{NH}_3^+ + \text{OH}^-$. Since $K_b$ is small (much less than 1), the equilibrium lies far to the left, meaning most methylamine molecules remain unreacted. Therefore, at equilibrium $[\text{CH}_3\text{NH}_2] > [\text{CH}_3\text{NH}_3^+]$. Choice B incorrectly assumes complete reaction to produce $[\text{OH}^-] = 0.10,\text{M}$, which would only occur for a strong base like NaOH. For weak bases, use the $K_b$ expression to recognize that small $K_b$ means little reaction, so the molecular form predominates.

This question tests understanding of weak acid and base equilibria. Hydrofluoric acid (HF) is a weak acid with $K_a = 6.8 \times 10^{-4}$, meaning it only partially dissociates in water according to $ \text{HF} \rightleftharpoons \text{H}^{+} + \text{F}^{-} $. Since $K_a$ is relatively small (much less than 1), the equilibrium lies far to the left, meaning most HF molecules remain undissociated. At equilibrium, $[\text{HF}] > [\text{F}^{-}]$ because only a small fraction of the original 0.10 M HF dissociates. Choice B ($[\text{H}^{+}] = 0.10,\text{M}$) is incorrect because it assumes complete dissociation, which only occurs for strong acids. For weak acids, always set up the equilibrium expression $K_a = \frac{[\text{H}^{+}][\text{F}^{-}]}{[\text{HF}]}$ and recognize that most of the acid remains in molecular form.

This question tests understanding of weak acid and base equilibria. For weak acids at the same initial concentration, the acid with larger Ka dissociates more and produces higher [H+]. Since HA has Ka = 1.0 × 10^-3 (larger) and HB has Ka = 1.0 × 10^-6 (smaller), HA dissociates more extensively. This means the HA solution has higher [H+] and therefore lower pH than the HB solution. Choice A incorrectly states that higher Ka leads to higher pH, confusing the inverse relationship between [H+] and pH. For weak acid comparisons, remember: larger Ka → more dissociation → higher [H+] → lower pH.

This question tests understanding of weak acid and base equilibria. Both acetic acid and HCN are weak acids that establish equilibria with their conjugate bases and H+ ions. The strength of a weak acid is determined by its Ka value - larger Ka means stronger acid and more dissociation. Since acetic acid has Ka = 1.8 × 10^-5 while HCN has Ka = 6.2 × 10^-10, acetic acid is the stronger acid and will produce more H+ ions at equilibrium. Choice D incorrectly assumes both acids fully dissociate, which only happens with strong acids like HCl or HNO3. For comparing weak acids at the same concentration, the one with larger Ka always produces more H+ and thus has lower pH.

This question assesses understanding of weak acid and base equilibria. Weak acids such as HA undergo incomplete dissociation, producing H+ and A- in equal amounts from the ionization reaction. The equilibrium Ka = [H+][A-]/[HA] governs the concentrations, but stoichiometry ensures [H+] = [A-] neglecting water's contribution. This equality holds because each dissociated HA molecule yields one H+ and one A-. Choice C is a tempting distractor, claiming $[A-]{eq}$ = $[HA]{eq}$, but this is wrong as weak acids have $[HA]{eq}$ >> $[A-]{eq}$. For weak species, write the equilibrium first before thinking about pH.

This question assesses understanding of weak acid and base equilibria. Weak bases like B partially react with water, not fully, leading to an equilibrium where OH- is produced. Kb = [BH+][OH-]/[B] describes this equilibrium, with a value like $4.0×10^{-4}$ indicating moderate dissociation and increased [OH-] compared to pure water. This governs concentrations, raising [OH-] and thus making the solution basic. Choice A is a common misconception, stating [H+] increases because BH+ is an acid, but actually [H+] decreases due to higher [OH-]. For weak species, write the equilibrium first before thinking about pH.

This question assesses understanding of weak acid and base equilibria. Weak bases like NH3 do not fully accept protons from water, resulting in partial reaction and an equilibrium state. The equilibrium is quantified by Kb = [NH4+][OH-]/[NH3], which dictates the concentrations of ions produced relative to the undissociated base. Water's concentration is omitted from Kb as it is constant, focusing the expression on the base's interaction. A common distractor is choice E, which includes [H2O] in the numerator, but this is erroneous since water is the solvent. For weak species, write the equilibrium first before thinking about pH.

This question assesses understanding of weak acid and base equilibria. Weak acids like acetic acid partially dissociate, forming H+ and CH3COO- in equilibrium. Ka = $1.8×10^{-5}$ indicates limited ionization, but the reaction stoichiometry ensures [H+] = [CH3COO-] from each dissociation event. This equality governs the concentrations, neglecting water's autoionization. Choice A is a common misconception, suggesting $[CH3COO-]{eq}$ > $[CH3COOH]{eq}$ because ions are favored, but small Ka means the opposite. For weak species, write the equilibrium first before thinking about pH.

This question assesses understanding of weak acid and base equilibria. Weak acids like HNO2 partially dissociate in water, establishing an equilibrium rather than complete ionization. The equilibrium constant Ka is defined as [H+][NO2-]/[HNO2], reflecting the extent of dissociation and controlling the concentrations of products relative to reactants. Since water is the solvent, its concentration is constant and not included in the Ka expression, ensuring the focus on the acid's ionization. A common misconception is choice E, which includes [H2O] in the numerator, but this is wrong because water's concentration is incorporated into Ka. For weak species, write the equilibrium first before thinking about pH.

This question assesses understanding of weak acid and base equilibria. Weak acids like HNO2 undergo incomplete dissociation in aqueous solution, establishing an equilibrium where HNO2 partially breaks into H+ and NO2- ions. The acid dissociation constant Ka is defined as the ratio of products to reactants, specifically Ka = [H+][NO2-]/[HNO2], reflecting the extent of dissociation and controlling the equilibrium concentrations. This expression ensures that the concentrations satisfy the equilibrium condition, with water not included as it's the solvent. A common misconception is that Ka = [HNO2]/[H+][NO2-], but this inverts the actual definition and would imply a constant for association rather than dissociation. For weak species, write the equilibrium first before thinking about pH.