Chemical Equlibrium flash cards

1. 1.What is meant by dynamic equilibrium?

   A state in a reversible process where the forward and backward reactions occur at equal rates, so the concentrations of reactants and products remain constant with time. It is dynamic because both reactions continue at the molecular level.

   Hint: Rates equal, not reactions stopped.

2. 2.State two key characteristics of chemical equilibrium.

   (1) It is dynamic — forward and reverse reactions proceed at equal rates. (2) It can be attained only in a closed system, and equilibrium concentrations remain constant at constant temperature.

   Hint: Closed system, constant T.

3. 3.What is a reversible reaction?

   A reaction in which products can react to re-form the reactants, i.e. it proceeds in both forward and backward directions simultaneously under the same conditions, denoted by $\rightleftharpoons$.

   Hint: Two-way arrow.

4. 4.Give one example each of physical equilibrium (solid-liquid and liquid-vapour).

   Solid-liquid: ice $\rightleftharpoons$ water at $0\,^{\circ}C$ and 1 atm. Liquid-vapour: water $\rightleftharpoons$ water vapour in a closed vessel (saturated vapour pressure).

   Hint: Melting point; saturated vapour.

5. 5.At the melting point, what is true about the rates in the ice–water equilibrium?

   The rate of melting of ice equals the rate of freezing of water, so the amounts of ice and water remain constant at $0\,^{\circ}C$ and 1 atm.

   Hint: Rate melt = rate freeze.

6. 6.For a liquid–vapour equilibrium in a closed container, what remains constant at a given temperature?

   The vapour pressure (saturated vapour pressure) remains constant, because the rate of evaporation equals the rate of condensation.

   Hint: Saturated vapour pressure.

7. 7.State the Law of Mass Action.

   At constant temperature, the rate of a chemical reaction is proportional to the product of the active masses (molar concentrations) of the reactants, each raised to a power equal to its stoichiometric coefficient.

   Hint: Rate $\propto$ product of active masses.

8. 8.For the reaction $aA + bB \rightleftharpoons cC + dD$, write the expression for the equilibrium constant $K_c$.

   $K_c = \dfrac{[C]^c [D]^d}{[A]^a [B]^b}$, where the concentrations are the equilibrium molar concentrations.

   Hint: Products over reactants, powers = coefficients.

9. 9.What does the magnitude of the equilibrium constant $K_c$ tell you about a reaction?

   A large $K_c$ ($>10^3$) means products predominate at equilibrium (reaction nearly complete). A small $K_c$ ($<10^{-3}$) means reactants predominate. $K_c \approx 1$ means appreciable amounts of both.

   Hint: Large = product-favoured.

10. 10.Define active mass as used in the law of mass action.

    Active mass is the molar concentration ($mol\,L^{-1}$) of a species; for gases it may be expressed as partial pressure. Pure solids and pure liquids are taken as unity (constant).

    Hint: Molar concentration; solids/liquids = 1.

11. 11.Why are pure solids and pure liquids omitted from the equilibrium constant expression?

    Their concentration (mole per unit volume, related to density) is essentially constant and is absorbed into the constant; hence their active mass is taken as unity.

    Hint: Constant density, active mass = 1.

12. 12.Write $K_p$ for the reaction $aA(g) + bB(g) \rightleftharpoons cC(g) + dD(g)$.

    $K_p = \dfrac{p_C^{\,c}\, p_D^{\,d}}{p_A^{\,a}\, p_B^{\,b}}$, where $p_i$ are the equilibrium partial pressures.

    Hint: Same form as $K_c$ but with partial pressures.

13. 13.State the relationship between $K_p$ and $K_c$.

    $K_p = K_c (RT)^{\Delta n_g}$, where $\Delta n_g =$ (moles of gaseous products) $-$ (moles of gaseous reactants) and $R$ is the gas constant, $T$ the temperature in kelvin.

    Hint: $(RT)^{\Delta n_g}$.

14. 14.For $N_2(g) + 3H_2(g) \rightleftharpoons 2NH_3(g)$, what is $\Delta n_g$ and the $K_p$–$K_c$ relation?

    $\Delta n_g = 2 - (1+3) = -2$, so $K_p = K_c (RT)^{-2} = \dfrac{K_c}{(RT)^2}$.

    Hint: Fewer gas moles on product side.

15. 15.For a reaction with $\Delta n_g = 0$ (e.g. $H_2 + I_2 \rightleftharpoons 2HI$), how do $K_p$ and $K_c$ compare?

    $K_p = K_c$, because $(RT)^{0} = 1$. Equal moles of gas on both sides.

    Hint: $\Delta n_g = 0 \Rightarrow K_p = K_c$.

16. 16.Are $K_c$ and $K_p$ dimensionless quantities?

    Strictly, the thermodynamic equilibrium constant is dimensionless (activities used). In NCERT numericals, $K_c$/$K_p$ may carry units depending on $\Delta n$; often units are dropped by convention.

    Hint: Thermodynamic $K$ is unitless.

17. 17.How does $K_c$ change if a reaction is reversed?

    The equilibrium constant of the reversed reaction is the reciprocal: $K_c' = \dfrac{1}{K_c}$.

    Hint: Reverse $\Rightarrow$ reciprocal.

18. 18.If all coefficients of a reaction are multiplied by $n$, how does $K_c$ change?

    The new equilibrium constant becomes $(K_c)^n$.

    Hint: Multiply coefficients $\Rightarrow$ power $n$.

19. 19.If two reactions are added, how are their equilibrium constants combined?

    The equilibrium constant of the overall (added) reaction is the product of the individual constants: $K = K_1 \times K_2$.

    Hint: Add reactions $\Rightarrow$ multiply K's.

20. 20.Define the reaction quotient $Q_c$.

    $Q_c$ has the same algebraic form as $K_c$ but uses the concentrations at any arbitrary instant (not necessarily equilibrium): $Q_c = \dfrac{[C]^c[D]^d}{[A]^a[B]^b}$.

    Hint: Same form as K, any instant.

21. 21.How is the direction of a reaction predicted by comparing $Q_c$ with $K_c$?

    If $Q_c < K_c$: reaction proceeds forward (more products form). If $Q_c > K_c$: reaction proceeds backward. If $Q_c = K_c$: system is at equilibrium.

    Hint: Q<K forward; Q>K backward.

22. 22.QUESTION: For a reaction $Q_c = 0.5$ and $K_c = 2.0$. In which direction does the reaction proceed?

    Since $Q_c (0.5) < K_c (2.0)$, the reaction proceeds in the forward direction to form more products until $Q_c = K_c$.

    Hint: Compare Q and K.

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