The following diagrams represent the equilibrium state for three different reactions of the type $\mathrm{A}+\mathrm{X} \rightleftharpoons \mathrm{AX}(\mathrm{X}=\mathrm{B}, \mathrm{C}$, or $\mathrm{D})$ :
Which reaction has the largest equilibrium constant?

A $0.321 \mathrm{~kg}$ mass is attached to a spring with a force constant of $13.3 \mathrm{~N} / \mathrm{m}$. If the mass is displaced $0.256 \mathrm{~m}$ from equilibrium and released, what is its speed when it is $0.128 \mathrm{~m}$ from equilibrium?

Consider the following equilibrium process at $686^{\circ} \mathrm{C}$.

$$\mathrm{CO}_2(\mathrm{~g})+\mathrm{H}_2(\mathrm{~g}) \rightleftharpoons \mathrm{CO}(\mathrm{g})+\mathrm{H}_2 \mathrm{O}(\mathrm{g})$$

The equilibrium concentrations of the reacting species are $[\mathrm{CO}]=0.050 \mathrm{M},\left[\mathrm{H}_2\right]=0.045 \mathrm{M}$, $\left[\mathrm{CO}_2\right]=0.086 \mathrm{M}$, and $\left[\mathrm{H}_2 \mathrm{O}\right]=0.040 \mathrm{M}$.
If we add $\mathrm{CO}_2$ to increase its concentration to $0.50 \mathrm{~mol} \mathrm{~L}^{-1}$, what will the concentrations of all the gases be when equilibrium is reestablished?

Consider the following equilibrium process at $686^{\circ} \mathrm{C}$.

$$\mathrm{CO}_2(\mathrm{~g})+\mathrm{H}_2(\mathrm{~g}) \rightleftharpoons \mathrm{CO}(\mathrm{g})+\mathrm{H}_2 \mathrm{O}(\mathrm{g})$$

The equilibrium concentrations of the reacting species are $[\mathrm{CO}]=0.050 \mathrm{M},\left[\mathrm{H}_2\right]=0.045 \mathrm{M}$, $\left[\mathrm{CO}_2\right]=0.086 \mathrm{M}$, and $\left[\mathrm{H}_2 \mathrm{O}\right]=0.040 \mathrm{M}$.
Calculate $K_{\mathrm{c}}$ for the reaction at $686^{\circ} \mathrm{C}$.

The equilibrium constant $K_{\mathrm{c}}$ for the decomposition of phosgene $\left(\mathrm{COCl}_2\right)$ is $4.69 \times 10^{-3}$ at $527^{\circ} \mathrm{C}$ :

$$\mathrm{COCl}_2(\mathrm{~g}) \rightleftharpoons \mathrm{CO}(\mathrm{g})+\mathrm{Cl}_2(\mathrm{~g})$$

Calculate the equilibrium partial pressure of all the components, starting with pure phosgene at $0.785$ bar.

Consider the following equilibrium process at $686^{\circ} \mathrm{C}$

$$\mathrm{CO}_2(g)+\mathrm{H}_2(g) \rightleftarrows \mathrm{CO}(g)+\mathrm{H}_2 \mathrm{O}(g)$$

The equilibrium concentrations of the reacting species are $[\mathrm{CO}]=0.050 \mathrm{M},\left[\mathrm{H}_2\right]=0.045 \mathrm{M}$, $\left[\mathrm{CO}_2\right]=0.086 M$, and $\left[\mathrm{H}_2 \mathrm{O}\right]=0.040 M$.

(a) Calculate $K_{\mathrm{c}}$ for the reaction at $686^{\circ} \mathrm{C}$.

(b) If we add $\mathrm{CO}_2$ to increase its concentration to $0.50 \mathrm{~mol} / \mathrm{L}$, what will the concentrations of all the gases be when equilibrium is reestablished?

HA has a $p K_{a}$ of 15, while HB has a $\mathrm{pK}_{\mathrm{a}}$ of 5. Draw the equilibrium that would result upon mixing HB with NaA. Does the equilibrium favor formation of HA or formation of HB?

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

$$\mathrm{I}_2(\mathrm{~g}) \rightleftharpoons 2 \mathrm{I}(\mathrm{g})$$

At $1000 \mathrm{~K}$, the equilibrium constant $K_{\mathrm{c}}$ for the reaction is $3.80 \times 10^{-5}$. Suppose you start with $0.0456$ mole of $\mathrm{I}_2$ in a $2.30$-L flask at $1000 \mathrm{~K}$. What are the concentrations of the gases at equilibrium?

The equilibrium constant $K_{\mathrm{c}}$ for the reaction

$$\mathrm{H}_2(\mathrm{~g})+\mathrm{Br}_2(\mathrm{~g}) \rightleftarrows 2 \mathrm{HBr}(\mathrm{g})$$

is $2.18 \times 10^6$ at $730^{\circ} \mathrm{C}$. Starting with $3.20$ moles of $\mathrm{HBr}$ in a 12.0-L reaction vessel, calculate the concentrations of $\mathrm{H}_2, \mathrm{Br}_2$, and $\mathrm{HBr}$ at equilibrium.

What is the half-life of a compound if 75 percent of a given sample of the compound decomposes in 60 min? Assume first-order kinetics.

At $1000 \mathrm{~K}$, a sample of pure $\mathrm{NO}_2$ gas decomposes:

$$2 \mathrm{NO}_2(\mathrm{~g}) \rightleftharpoons 2 \mathrm{NO}(\mathrm{g})+\mathrm{O}_2(\mathrm{~g})$$

The equilibrium constant $K_p$ is 160. Analysis shows that the partial pressure of $\mathrm{O}_2$ is $0.25$ bar at equilibrium. Calculate the pressure of $\mathrm{NO}$ and $\mathrm{NO}_2$ in the mixture.

Arrange the following in order of increasing dipole moment. (a) (E)-1-Chloropropene (b) 1,1-Dichloro-2-methylpropene (c) (E)-2,3-Dichloro-2-butene

For the reaction

$$2 \mathrm{C}_2 \mathrm{H}_6(g)+7 \mathrm{O}_2(g) \longrightarrow 4 \mathrm{CO}_2(g)+6 \mathrm{H}_2 \mathrm{O}(g)$$

predict the enthalpy of reaction from the average bond enthalpies in above Table.