Equilibrium practice · Mixed difficulty

Option C is correct because the assertion is true (solubility of sparingly soluble salts generally increases with temperature), but the reason is false (enthalpy of solution can be positive or negative).

Option C (CH3COONa, 1.25 M) is correct because increasing the concentration of CH3COONa to 1.25 M will decrease the [H+] enough to start the precipitation of MnS.

The correct answer is D ($3.1 \times 10^{-4}$) because it is calculated using the equilibrium constant expression in terms of partial pressures, considering the mole fractions and total pressure.

Option D ($7.0$) is correct because ammonium acetate is a salt of a weak acid and a weak base, and since pKa = pKb, the pH is 7.

Option B is correct. By manipulating the given equilibrium reactions and their constants, the equilibrium constant for D $\rightleftharpoons$ 4A is found to be Z/(X²Y).

Option A (9.9) is correct because the pH of the buffer solution is calculated using the Henderson-Hasselbalch equation.

Option A ($< {K}_{{SP}}$) is correct because the calculated ionic product is less than the given solubility product, indicating that a precipitate will not form.

Option C (5) is correct because the pH must be greater than or equal to 4.253 to dissolve the precipitate, making pH 5 the appropriate choice.

Option A (5.05) is correct because it accurately calculates the pH of the buffer solution using the Henderson-Hasselbalch equation.

Option B (10) is correct because it is calculated by first finding the molarity of the NaOH solution, then calculating the pOH, and finally determining the pH.

Option A (1.74) is correct because the equilibrium constant Kc is calculated based on the degree of association and the initial concentration of benzoic acid.

Option C is correct. HCl is a strong acid and dissociates completely, resulting in a lower pH than acetic acid, which is a weak acid. The reason is incorrect because HCl has more titrable protons.

Option D (All of the above) is correct because all given mixtures can act as buffer solutions, either by containing a weak acid/base and its conjugate or by forming one upon reaction.

The correct answer is A (A low value for the equivalent weight) because methyl orange changes color at a lower pH than the equivalence point of a weak acid titration, leading to premature endpoint detection and an underestimation of the volume of titrant required, thus a lower calculated equivalent weight.

Option A is correct because the pH range of a buffer solution is typically $\mathrm{pK}_{\mathrm{a}} \pm 1$, so for a weak acid with $\mathrm{pK}_{\mathrm{a}}=6$, the pH range is 5 to 7.

Option B ($5 \times 10^{-10} \mathrm{M}$) is correct because the solubility product expression is used to calculate the solubility of AgCl in the presence of a common ion.

Option C is correct because a buffer solution resists changes in pH upon the addition of small amounts of acid or base.

Option A (2.0) is correct because a lower $\mathrm{p}{K_a}$ value indicates a stronger acid.

Option B (0.0625) is correct because the equilibrium constant Kc is calculated by dividing the product concentration by the reactant concentrations, each raised to the power of their stoichiometric coefficients.

Option B (Pressure is decreased) is correct because decreasing pressure shifts the equilibrium to the side with more moles of gas, which is the reactants' side, thus decreasing ammonia yield.

Option C is correct. The order of relative base strength is determined by the strength of their conjugate acids, with weaker acids corresponding to stronger bases. OH- is the strongest base, followed by CH3COO-, and then Cl-.

Option A ($\frac{2}{3}$) is correct because the equilibrium constant is calculated as the ratio of the concentration of products to reactants at equilibrium, which in this case is $\frac{2}{3}$.

Option C ($\sqrt{\frac{8 y^3}{x}}$) is correct because it correctly applies the solubility product principle and the equilibrium relationship between the concentrations of fluoride and carbonate ions.

Option B (${\mathrm{H}_3}{\mathrm{O}^ + }$) is correct because it can accept a pair of electrons due to the positive charge, making it a Lewis acid.

Option B ($10^{-3}$ to $10^3$) is correct because when the equilibrium constant $K_c$ is within this range, the reaction mixture contains appreciable amounts of both reactants and products.

Option B ($S\, = \,{\left( {\frac{{{{K_{sp}}}}}{{{{6912}}}}} \right)^{1/7}}$) is correct because it represents the correct relationship between molar solubility (S) and the solubility product (Ksp) for Zr3(PO4)4.

Option D is correct because both the assertion and the reason are false. The solubility product of $\mathrm{Ag}_2 \mathrm{CrO}_4$ is $4S^3$, and the expression for $K_{sp}$ is $\mathrm{K_{sp}} = [\mathrm{Ag}^+]^2[\mathrm{CrO}_4^{2-}]$.

Option D ($10^{-13} {M}$) is correct because the hydrogen ion concentration is calculated using the ion product of water ($K_w = 10^{-14}$).

Option D (8) is correct because it is calculated using the Henderson-Hasselbalch equation for a buffer solution, considering the pKb of ammonia and the concentrations of ammonia and ammonium chloride.

Option A ($80 \%$ ) is correct because it's the degree of association calculated using the van't Hoff factor and the freezing point depression.