Consider the following equilibrium process at 686 degree C: CO2 (g) + H2O (g) >=> CO (g) + H2O (g) The equilibrium concentrations of the reacting species are [CO] = 0.050 M, [H2] = 0.045 M, [CO2] = 0.086 M, and [H2O] = 0.040 M. (a) Calculate Kc for the reaction at 686°C. (b) If we add CO2 to increase its concentration to 0.50 mol/L, what will the concentrations of all the gases be when equilibrium is reestablished?

The dissociation of molecular iodine into iodine atoms is represented as I2 (g) <=> 2 I (g) At 1000 K, the equilibrium constant Kc for the reaction is 3.80*10^-5. Suppose you start with 0.0456 mole of I2 in a 2.30-L flask at 1000 K. What are the concentrations of the gases at equilibrium?

At equilibrium, the pressure of the reacting mixture CaCO3(s) <=> CaO(s) + CO2(g) is 0.105 atm at 350 degree C. Calculate Kp and Kc for this reaction.

A 2.50-mole quantity of NOCl was initially in a 1.50-L reaction chamber at 400 degree C. After equilibrium was established, it was found that 28.0 percent of the NOCl had dissociated: 2 NOCl (g) <=> 2 NO (g) + Cl2 (g) Calculate the equilibrium constant Kc for the reaction.

Ammonium carbamate, NH4CO2NH2, decomposes as follows: NH4CO2NH2(s) <=> 2 NH3 (g) + CO2 (g) Starting with only the solid, it is found that at 40°C the total gas pressure (NH3 and CO2) is 0.363 atm. Calculate the equilibrium constant Kp.

The equilibrium constant Kc for the reaction H2 (g) + Br2 (g) <=> 2 HBr (g)
is 2.18 $\times$ 10^6 at 730°C. Starting with 3.20 moles of HBr in a 12.0-L reaction vessel, calculate the concentrations of H2, Br2, and HBr at equilibrium.

At 700 K, the equilibrium constant for the reaction $\mathrm { CCl } _ { 4 } ( g ) \rightleftharpoons \mathrm { C } ( s ) + 2 \mathrm { Cl } _ { 2 } ( g )$ is $K_p= 0.76.$ A flask is charged with 2.00 atm of $CCl_4,$ which then reaches equilibrium at 700 K. What fraction of the $CCl_4$ is converted into C and $Cl_2?$

At 1000 K, a sample of pure NO2 gas decomposes: 2 NO2 (g) <=> 2 NO (g) + O2(g) The equilibrium constant Kp is 158. Analysis shows that the partial pressure of O2 is 0.25 atm at equilibrium. Calculate the pressure of NO and NO2 in the mixture.

Consider the following equilibrium system involving SO2, Cl2, and SO2Cl2 (sulfuryl dichloride): SO2 (g) + Cl2 (g) <=> SO2Cl2 (g) Predict how the equilibrium position would change if (a) Cl2 gas were added to the system; (b) SO2Cl2 were removed from the system; (c) SO2 were removed from the system. The temperature remains constant

Consider the following equilibrium process: PCl5 (g) <=> PCl3(g) + Cl2(g) ; $\Delta H^\circ$ = 92.5 kJ/mol Predict the direction of the shift in equilibrium when (a) the temperature is raised; (b) more chlorine gas is added to the reaction mixture; (c) some PCl3 is removed from the mixture; (d) the pressure on the gases is increased; (e) a catalyst is added to the reaction mixture.

Heating solid sodium bicarbonate in a closed vessel establishes the following equilibrium: 2 NaHCO3 (s) <=> Na2CO3 (s) + H2O(g) + CO2 (g) What would happen to the equilibrium position if (a) some of the CO2 were removed from the system; (b) some solid Na2CO3 were added to the system; (c) some of the solid NaHCO3 were removed from the system? The temperature remains constant.

The equilibrium constant Kc for the following reaction is 1.2 at 375 degree C. N2 (g) + H2 (g) <=> 2 NH3 (g) (a) What is the value of KP for this reaction? (b) What is the value of the equilibrium constant Kc for 2 NH3 (g) <=> N2 (g) + 3 H2 (g) (c) What is the value of Kc for 1/2 N2 (g) + 3/2 H2 (g) <=> NH3 (g) ? (d) What are the values of Kp for the reactions described in (b) and (c)?

The equilibrium constant (Kc) for the reaction 2HCl(g) <=> H2(g) + Cl2(g) is 4.17 $\times$ 10^-34 at 25 degree C. What is the equilibrium constant for the reaction H2(g) + Cl2(g) <=> 2HCl(g) at the same temperature?

List four factors that can shift the position of an equilibrium. Only one of these factors can alter the value of the equilibrium constant. Which one is it?