Given the following two reactions:
     A + B ⇌ C + D, with rate constant K₁
     E + F ⇌ G + H, with rate constant K₂

If C + D + E + F produces a product, what is the rate constant for this reaction?

The equilibrium constant Kp for the following reaction is:
\(\mathrm{MgCO}_{3(\mathrm{~s})} \rightleftharpoons \mathrm{MgO}_{(\mathrm{s})}+\mathrm{CO}_{2(\mathrm{~g})}\)
1. \(\mathrm{K_P} =\mathrm{P}_{\mathrm{CO}_2}\)
2. \(\mathrm{K_P} =\mathrm{P}_{\mathrm{CO}_2} \times \frac{\mathrm{P}_{\mathrm{CO}_2 \times \mathrm{P}_{\mathrm{MgO}^{\mathrm{}}}}}{\mathrm{P}_{\mathrm{MgCO}_3}}\)
3. \(\mathrm{K_P} =\frac{\mathrm{P}_{\mathrm{CO}_2}+\mathrm{P}_{\mathrm{MgO}^{\mathrm{}}}}{\mathrm{P}_{\mathrm{MgCO}_3}}\)
4. \(\mathrm{KP} =\frac{\mathrm{P}_{\mathrm{MgCO}_3}}{\mathrm{P}_{\mathrm{CO}_2} \times \mathrm{P}_{\mathrm{MgO}}}\)

What happens to the equilibrium constant of a reversible reaction if the concentration of reactants is increased?
1. It depends on the amount of concentration
2. It remains unchanged
3. It decrease
4. It increase

The ionization constant of \(CH_{3} COOH\) is 1.7 × \(\left(10\right)^{- 5}\) and the concentration of \(H^{+ }\)ions is 3.4 × \(\left(10\right)^{- 4}\). The initial concentration of \(CH_{3} COOH\) will be:
1.  3.4 × \(\left(10\right)^{- 4}\)
2.  3.4× \(\left(10\right)^{- 3}\)
3.  6.8 × \(\left(10\right)^{- 4}\)
4.  6.8 × \(\left(10\right)^{- 3}\)

Find the factor on which the equilibrium pressure of O₂ depends for the reaction:

2BaO₂(s) ⇌ 2BaO(s) + O₂(g)  ΔH > 0

The reaction quotient (Q) for the reaction:

${\mathrm{N}}_2\left(\mathrm{g}\right)+3{\mathrm{H}}_2\left(\mathrm{g}\right)\rightleftharpoons 2{\mathrm{NH}}_3\left(\mathrm{g}\right)$

is given by $\mathrm{Q}=\frac{{\left[{\mathrm{NH}}_3\right]}^2}{\left[{\mathrm{N}}_2\right]{\left[{\mathrm{H}}_2\right]}^3}$.
Under what conditions will the reaction proceed from right to left?

1. Q = K_C
2. Q < K_C
3. Q > K_C
4. Q = 0

(where K_C is the equilibrium constant)