Assume that the decomposition of HNO₃ can be represented by the following equation

$4{\mathrm{HNO}}_3\left(\mathrm{g}\right)\rightleftharpoons 4{\mathrm{NO}}_2\left(\mathrm{g}\right)+2{\mathrm{H}}_2\mathrm{O}\left(\mathrm{g}\right)+{\mathrm{O}}_2\left(\mathrm{g}\right)$

and the reaction approaches equilibrium at 400K temperature and 30 atm pressure. At equilibrium, partial pressure of HNO₃ is 2 atm. Calculate K_c in (mol/L)³ at 400 K:

(Use : R=0.08 atm-L/mol-K)

1. 4

2. 8

3. 16

4. 32

A 250 mL flask contains NO(g) at a pressure of 0.46 atm, and a 100 mL flask contains O₂(g) at 0.86 atm, both at 350 K. The two flasks are connected by a stopcock. When the stopcock is opened, the gases mix, and the following reactions occur:

2 NO(g) + O₂(g) → 2 NO₂(g) (this reaction goes to completion)

followed by the equilibrium:

2 NO₂(g) ⇌ N₂O₄(g)

After both processes, the total pressure of the mixture in the combined system (total volume = 350 mL) is 0.37 atm. Calculate Kp for the equilibrium 2 NO₂(g) ⇌ N₂O₄(g) at 350 K. (Kp has units of atm⁻¹.)

1. 0.645 atm^-1 

2. 1.290 atm^-1 

3. 20 atm^-1 

4. 5.46 atm^-1 

The equilibrium constant for the reactions

${\mathrm{N}}_2+3{\mathrm{H}}_2\rightleftharpoons 2{\mathrm{NH}}_3<mspace\; width="1em"></mspace>\mathrm{and}<mspace\; width="1em"></mspace>\frac{1}{2}{\mathrm{N}}_2+\frac{3}{2}{\mathrm{H}}_2\rightleftharpoons {\mathrm{NH}}_3$

are K₁ and K₂ respectively. The correct relationship between K₁ and K₂ is

1. ${\mathrm{K}}_1=\frac{{\mathrm{K}}_2}{2}$

2. ${\mathrm{K}}_2=\sqrt{{\mathrm{K}}_1}$

3. ${\mathrm{K}}_2={{\mathrm{K}}_1}^2$

4. ${\mathrm{K}}_1={\mathrm{K}}_2$

Find the value of Kc for the equilibrium:

NH₃(g) ⇌ ½N₂(g) + 3/2 H₂(g)

Given: 1 mol of N₂ and 2 mol of H₂ are taken in a 4 L vessel, and 50% of N₂ is converted into NH₃.

1. 256

2. 16

3. $\frac{1}{16}$

4. None of the above

If $\mathrm{pKa}=-\log <mspace\; width="1em"></mspace>{\mathrm{K}}_{\mathrm{a}}=4<mspace\; width="1em"></mspace>\mathrm{and}<mspace\; width="1em"></mspace>{\mathrm{K}}_{\mathrm{a}}={\mathrm{Cx}}^2$ then Van't Hoff factor of weak monobasic acid when C=0.01 M is

1. 1.02

2. 1.01

3. 1.20

4. 1.10

The equilibrium constant (K_p) for the decomposition of water is given by:
\(H_2O(g) \rightarrow H_2(g) + {1 \over 2} O_2 (g)\)

Which one of the following is the correct option?
[α = degree of dissociation & p = pressure of the reaction mixture at equilibrium]

Which of these can not be a Bronsted acid?

1. ${\mathrm{HCO}}_3^{-}$

2. ${\mathrm{HCOO}}^{-}$

3. ${\mathrm{H}}_2{\mathrm{PO}}_3^{-}$

4. ${\mathrm{H}}_3{\mathrm{O}}^{+}$

Find the correct order of proton-accepting tendency (Brønsted basic strength) among HS⁻, I⁻, RNH₂, and NH₃.

Which one of the following orders of acid strength is correct?

1. $\mathrm{RCOOH}>\mathrm{HC}\equiv \mathrm{CH}>\mathrm{HOH}>\mathrm{ROH}$

2. $\mathrm{RCOOH}>\mathrm{ROH}>\mathrm{HOH}>\mathrm{HC}\equiv \mathrm{CH}$

3. $\mathrm{RCOOH}>\mathrm{HOH}>\mathrm{ROH}>\mathrm{HC}\equiv \mathrm{CH}$

4. $\mathrm{RCOOH}>\mathrm{HOH}>\mathrm{HC}\equiv \mathrm{CH}>\mathrm{ROH}$

A centinormal solution of a monobasic acid is 100% ionized. Its pH is:

1. 2

2. 4

3. 3

4. 1