Kp Kc Factors Affecting Them

For the reaction, 
CH₄(g) + 2O₂(g) CO₂(g) + 2H₂O(l) 
Δ_rH = -170.8 kJ mol^-1 
Which of the following statements is not true?

The following equilibrium constants are given:
N₂ + 3H₂ ⇌ 2NH₃; K₁
N₂ + O₂ ⇌ 2NO; K₂
H₂ + 1/2O₂ ⇌ H₂O; K₃

The equilibrium constant for the oxidation of NH₃ by oxygen to give NO is:

1. $K_2{K}_3^3/K_1$

2. $K_2{K}_3^2/K_1$

3. $K_2^2{K}_3/K_1$

4. $K_1{K}_2/K_3$

The dissociation equilibrium of a gas AB₂ can be represented as
$2A{B}_2\left(g\right)$ $\leftrightharpoons$ $2AB\left(g\right)$ $+$ $B_2\left(g\right)$ $$
The degree of dissociation is ‘x’ and is small compared to 1. The expression relating the degree of dissociation (x) with equilibrium constant K_P and total pressure p is:
1. (2KP/p)
2. (2Kp /p)^1/3
3. (2KP/p)^1/2
4. (KP/P)

The value of K_P1 and K_p2 for the reactions
$X$ $\leftrightharpoons$ $Y$ $+Z$ $......\left(i\right)$
$and$ $A$ $\leftrightharpoons 2B$ $......\left(ii\right)\hspace{0.17em}$
are in the ratio of 9 : 1. If the degree of dissociation of X and A is equal, then the total pressure at equilibrium(i) and (ii) are in the ratio:

The value of the equilibrium constant of the reaction
is 8.0.
The equilibrium constant of the reaction
will be-

The dissociation constants for acetic acid and HCN at 25 °C are 1.5 x 10^-5 and 4.5 x 10^-10, respectively. The equilibrium constant for the equilibrium,
CN^- + CH₃COOH ⇌ HCN + CH₃COO^-
would be:
1. $3.0\backslash times{10}^5$
2. $3.0\backslash times{10}^{-5}$
3. $3.0\backslash times{10}^{-4}$
4. $3.0\backslash times{10}^4$

The equilibrium reaction that doesn't have equal values for K_c and K_p is:

1.
2.
3.
4.

For the reaction ${\mathrm{N}}_2\left(\mathrm{g}\right)$ $+$ ${\mathrm{O}}_2\left(\mathrm{g}\right)\leftrightharpoons 2\mathrm{NO}\left(\mathrm{g}\right)$ the equilibrium constant is K₁. The equilibrium constant is K₂ for the reaction $2\mathrm{NO}\left(\mathrm{g}\right)$ $+$ ${\mathrm{O}}_2\left(\mathrm{g}\right)$ $\leftrightharpoons 2{\mathrm{NO}}_2\left(\mathrm{g}\right)$ $$
The value of K for the reaction given below will be:
${\mathrm{NO}}_2\left(\mathrm{g}\right)\leftrightharpoons \frac{1}{2}{\mathrm{N}}_2\left(\mathrm{g}\right)$ $+{\mathrm{O}}_2\left(\mathrm{g}\right)$ $$

1. $\frac{1}{4}$ $\left(4\right)$ ${\mathrm{K}}_1$ ${\mathrm{K}}_2$

2. ${\left[\frac{1}{{\mathrm{K}}_1{\mathrm{K}}_2}\right]}^{1/2}$

3. $\frac{1}{\left({\mathrm{K}}_1{\mathrm{K}}_2\right)}$

4. $\frac{1}{\left(2{\mathrm{K}}_1{\mathrm{K}}_2\right)}$

The equilibrium constant for the reaction:

What will be the equilibrium constant for the reaction: N₂(g) + O₂(g) ⇄ NO(g)?

1. $K^{\frac{1}{2}}$

2. $\frac{1}{2}K$

3. K

4. $K^2$

Consider the given equilibrium constants:

The equilibrium constant (K) of the following reaction will be:

A 20-litre container at 400 K contains CO₂ (g) at pressure 0.4 atm and an excess of SrO (neglecting the volume of solid SrO). The volume of the container is now decreased by moving the movable piston fitted in the container.
The maximum volume of the container, when the pressure of CO₂ attains its maximum value will be:
(Given that: SrCO₃(s) ⇋ SrO(s) + CO₂(g) , Kp = 1.6 atm)