Understanding Chemical Equilibrium

For a reaction at equilibrium, if the rate of the forward reaction increases, what change must occur in the reverse reaction?

Consider the reaction $aA + bB \rightleftharpoons cC + dD$. How is the equilibrium constant expression, $K_c$, formulated?

What condition is necessary to calculate the value of $K_c$ from known equilibrium amounts?

Consider the reversible reaction $N_2(g) + 3H_2(g) \rightleftharpoons 2NH_3(g)$. If, at equilibrium, the partial pressures are $P_{N_2} = 2 atm$, $P_{H_2} = 3 atm$, and $P_{NH_3} = 4 atm$, what is the value of $K_p$?

At 700 K, the reaction $2SO_2(g) + O_2(g) \rightleftharpoons 2SO_3(g)$ has an equilibrium constant $K_c = 4.2 \times 10^5$. If the reaction mixture contains [SO₂] = 0.20 M, [O₂] = 0.25 M, and [SO₃] = 85 M, which direction will the reaction proceed to reach equilibrium?

In the reaction $CO(g) + 2H_2(g) \rightleftharpoons CH_3OH(g)$, what change will occur if the volume of the container is decreased?

Given the endothermic reaction $N_2O_4(g) \rightleftharpoons 2NO_2(g)$, what effect will increasing the temperature have on the value of $K_c$?

Consider the equilibrium: $2SO_2(g) + O_2(g) \rightleftharpoons 2SO_3(g)$ If at equilibrium, some $O_2$ is added to the system, what will happen to the amount of $SO_3$?

Given the reaction $H_2(g) + I_2(g) \rightleftharpoons 2HI(g)$, with $K_c = 50$ at a certain temperature. If a reaction vessel initially contains [H₂] = 0.2 M, [I₂] = 0.2 M, and [HI] = 0.0 M, what will happen initially in order for said reaction to obtain equilibrium?

For the reaction $A(g) + B(g) \rightleftharpoons C(g)$, the equilibrium constant $K_c$ is 4. If the initial concentrations are [A] = 2 M, [B] = 2 M, and [C] = 0 M, what is the equilibrium concentration of C?

How does a catalyst affect a reaction at equilibrium?

What is the molar composition of the equilibrium mixture? $PCl_5(g) \rightleftharpoons PCl_3(g) + Cl_2(g)$. Initially 1.00 mol $PCl_5$ in a 1.00-L container, and at equilibrium 0.135 mol $PCl_3$.

Methanol is commercially made by $2H_2(g) + CO(g) \rightleftharpoons CH_3OH(g)$. What is the $K_c$ expression for this reaction?

Carbon dioxide decomposes at elevated temperature by $2CO_2(g) \rightleftharpoons 2CO(g) + O_2(g)$. At 3000 K, 2.00 mol $CO_2$ placed in 1.00-L container, at equilibrium, 0.90 mol $CO_2$ remains. What is the value of $K_c$ at this temperature?

Write the $K_c$ expression for the following heterogeneous reaction: $H_2O(g) + C(s) \rightleftharpoons CO(g) + H_2(g)$

Given: $aA + bB \rightleftharpoons cC + dD; K_1$. What is the $K_2$ if the reaction is reversed?

Given: $aA + bB \rightleftharpoons cC + dD; K_1$. What is the $K_2$ if the reaction is doubled?

How are $K_p$ and $K_c$ related? $aA(g) + bB(g) \rightleftharpoons cC(g) + dD(g)$

If the value of $K_c$ at $227^\degree C$ is 0.0952 for the following reaction: $CH_3OH(g) \rightleftharpoons CO(g) + 2H_2(g)$, what is $K_p$ at this temperature?

For $K_c$ = 0.82 for a reaction, which statement best describes the composition of the equilibrium mixture?

Nickel(II) oxide reduced to the metal by $CO(g) + NiO(s) \rightleftharpoons CO_2(g) + Ni(s)$. If the partial pressure of CO is 100. mmHg and the total pressure of CO and $CO_2$ does not exceed 1.0 atm, will this reaction occur at 1500 K at equilibrium? ($K_p$ = 700. at 1500 K)

Nitrogen and oxygen form nitric oxide by $N_2(g) + O_2(g) \rightleftharpoons 2NO(g)$. If an equilibrium mixture at $25^\circ C$ contains 0.040 M $N_2$ and 0.010 M $O_2$, what is the concentration of NO in this mixture? The equilibrium constant at $25^\circ C$ is $1.0 \times 10^{-30}$

Hydrogen iodide decomposes by $2HI(g) \rightleftharpoons H_2(g) + I_2(g)$. At 800 K, the equilibrium constant, $K_c$, for this reaction is 0.016. IF 0.50 mol HI placed in a 5.0-L flask, what is the composition of the equilibrium mixture in molarities?

$N_2O_4$ decomposes by $N_2O_4(g) \rightleftharpoons 2NO_2(g)$. At $100^\circ C$, $K_c$ = 0.36. If a 1.00-L flask initially contains 0.100 M $N_2O_4$, what will be the equilibrium concentration of $NO_2$?

Initial conditions: $H_2(g) + F_2(g) \rightleftharpoons 2HF(g)$; $K_c = 1.15 \times 10^2$. 3.000 mol of each species is put in 1.500-L vessel. What is the equilibrium concentration of each species?

The following reaction is at equilibrium: $COCl_2(g) \rightleftharpoons CO(g) + Cl_2(g)$. If chlorine gas is added to the reaction mixture what direction will the reaction shift?

The following reaction is at equilibrium: $COCl_2(g) \rightleftharpoons CO(g) + Cl_2(g)$. If carbon monoxide gas is removed from the mixture what direction will the direction shift?

In which direction will each reaction shift when the volume of the reaction container is increased?

If $\Delta H^\degree$ = 484 kJ for the reaction $2H_2O(g) \rightleftharpoons 2H_2(g) + O_2(g)$, would you expect this reaction to be favorable at high or low temperatures?

The Fischer-Tropsch process involves $8CO(g) + 17H_2(g) \rightleftharpoons C_8H_{18}(g) + 8H_2O(g)$. Suppose the reaction is at equilibrium at 200°C, then cooled to condense the octane, and then the remaining gases are reheated to 200°C. In which direction will the equilibrium shift?

A typical reaction that occurs in the Fischer–Tropsch process is $8CO(g) + 17H_2(g) \rightleftharpoons C_8H_{18}(g) + 8H_2O(g)$. In which direction will the equilibrium shift when the pressure is increased?

Which statement accurately describes dynamic equilibrium in a reversible chemical reaction?

For the synthesis of ammonia, $N_2(g) + 3H_2(g) \rightleftharpoons 2NH_3(g)$, indicate how $K_p$ and $K_c$ are related.

For a reaction where solids or pure liquids are involved, what is their contribution to the $K_c$ expression?