NESA Chemistry Calculating the Equilibrium Constant

This is incorrect because the stoichiometric coefficient of the reactant (2) must be used as an exponent, not a multiplier. The correct term is $[\text{HI}]^2$, not $2[\text{HI}]$.

This is incorrect because it uses atomic species (H, I) instead of the molecular species ($\text{H}_2$, $\text{I}_2$) actually present in the reaction, and it treats coefficients as multipliers rather than exponents.

This is incorrect because it substitutes atomic concentrations for molecular products and uses coefficients as multipliers in the numerator.

This value is $(K_c)^2$, which would be the equilibrium constant if the first reaction were multiplied by 2 but not reversed.

This value is $2 \times K_c$. When a reaction is multiplied by a coefficient, the equilibrium constant must be raised to that power, not multiplied by it.

This value results from incorrect mathematical manipulation of the equilibrium constant. The correct operation is to take the inverse square of the original $K_c$.

Statement (iii) is incorrect. While it is true that $[H^+] = [OH^-]$ in pure water, water only ionises to a very small extent, not completely.

Statements (iii) and (iv) are both incorrect. Water does not ionise completely, and its self-ionisation produces equal amounts of hydrogen and hydroxide ions, not twice as many.

This option is incorrect because statements (iii) and (iv) are false. Water only partially ionises and produces a 1:1 ratio of hydrogen to hydroxide ions.

Equal concentrations of reactants and products do not generally occur at equilibrium, especially when the equilibrium constant $K_c$ is significantly different from 1.

A greater concentration of reactants than products would only be expected if the equilibrium constant $K_c$ was much less than 1, indicating a reactant-favored system.

The coefficients in the balanced chemical equation dictate the stoichiometric ratio in which substances react and form, not their absolute concentrations at equilibrium.

This expression incorrectly places the reactants in the numerator and the products in the denominator. Additionally, pure liquids like $\text{H}_2\text{O}$ are excluded from equilibrium expressions.

This option incorrectly places reactants in the numerator and includes liquid water. Equilibrium expressions follow the ratio $\frac{[\text{products}]}{[\text{reactants}]}$.

This expression incorrectly includes water in the denominator. Pure liquids are omitted from equilibrium constants, and the reactant $\text{NH}_3$ is missing.

Incorrect. When a reaction equation is multiplied by a factor of 2, the equilibrium constant must be squared, not multiplied by 2.

Incorrect. Although the units are correctly squared, the numerical value was incorrectly multiplied by 2 instead of being squared.

Incorrect. The numerical value was correctly squared, but the units must also be squared from $\text{M}^{-2}$ to $\text{M}^{-4}$.

The reaction is endothermic ($\Delta\text{H} > 0$). Decreasing the temperature shifts the equilibrium to the left, which decreases the equilibrium constant $K$.

Changing the partial pressure of a reactant or product shifts the equilibrium position, but it does not change the value of the equilibrium constant $K$, which is only temperature-dependent.

Changing the total pressure of the system will shift the equilibrium position (in this case, to the left where there are fewer moles of gas), but it will not change the equilibrium constant $K$.

This expression is inverted. The equilibrium constant $K$ is defined as the ratio of product concentrations to reactant concentrations, not the other way around.

This option is both inverted (reactants over products) and incorrectly includes the solid $Cu(OH)_2$. Pure solids are excluded from equilibrium expressions because their concentration remains constant.

This option incorrectly includes the concentration of the solid reactant, $Cu(OH)_2$. In heterogeneous equilibria, pure solids and liquids are omitted from the $K$ expression.

This option incorrectly places reactants in the numerator and products in the denominator, and it fails to raise the concentrations to the power of their stoichiometric coefficients.

While this option correctly places products over reactants, it fails to raise the concentrations to the power of their stoichiometric coefficients found in the balanced equation.

This expression is the inverse of the correct equilibrium constant; it incorrectly places the reactants in the numerator and the products in the denominator.

This incorrectly assumes the concentration of $\text{O}_2(g)$ is proportional to the amount of solid $\text{Li}_2\text{O}_2$. The amounts of pure solids do not affect the equilibrium position.

This ratio has no basis in the equilibrium expression. The amounts of solid reactants and products do not change the equilibrium concentration of the gas.

This is incorrect because the equilibrium concentration of $\text{O}_2(g)$ depends only on the temperature-dependent equilibrium constant, not on the initial amounts of solids.

This option represents a 1:1 ratio of A to B. Since the initial mixture already contained more A than B and the equilibrium shifts further toward A upon cooling, the final mixture must have a ratio of A to B significantly greater than 1.

This option shows a higher amount of product B than reactant A. Decreasing the temperature of an endothermic reaction favors the reverse reaction, which would decrease the amount of B and increase the amount of A.

This option implies a shift toward the product B. According to Le Chatelier's principle, removing heat (lowering temperature) from an endothermic reaction drives the system toward the reactants, so A should be the predominant species.