Three moles of PCl₃ three moles of PCl₃, and two moles of Cl, are taken in a closed vessel. If at equilibrium the vessel has 1.5 moles of PCl₅ the number of moles of PCl₃, present in it is

• 5

• 3

• 6

• 4.5

For the reaction $2\mathrm{HI} \left(\mathrm{g}\right)\hspace{0.17em}\rightleftharpoons {\mathrm{H}}_2\left(\mathrm{g}\right) + {\mathrm{I}}_2\left(\mathrm{g}\right)\hspace{0.17em}-\mathrm{Q} \mathrm{kJ}$, the equilibrium constant depends upon

• temperature

• pressure

• catalyst

• volume

N₂ + 3H₂ $\rightleftharpoons$2NH₃ + heat.What is the effect of the increase of temperature on the equilibrium of the reaction? 

• Equilibrium is shifted to the left

• Equtlibriumis shifted to the right

• Equilibrium is unaltered

• Reaction rate does not change

654.

30 cc of $\frac{\mathrm{M}}{3}$, 20 cc of $\frac{\mathrm{M}}{2}$ HNO₃ and 40 cc of $\frac{\mathrm{M}}{4}$ NaOH solutions are mixed and the volume was made up to 1 dm³. The pH of the resulting solution is

• 8

• 2

• 1

• 3

10^-6 M NaOH is diluted 100 times. The pH of the diluted base is

• between 7 and 8

• between 5 and 6

• between 6 and 7

• between 10 and 11

2HI (g) $\rightleftharpoons$ H₂ (g) + I₂ (g)

The equilibrium constant of the above reaction is 6.4 at 300 K. If 0.25 mole each of H₂ and I₂ are added to the system, the equilibrium constant will be

• 6.4

• 0.8

• 3.2

• 1.6

Consider the following gaseous equilibria with equilibrium constants K₁ and K₂ respectively.

$$\begin{gathered} {\mathrm{SO}}_2\left(\mathrm{g}\right) + \frac{1}{2}{\mathrm{O}}_2\left(\mathrm{g}\right) \rightleftharpoons {\mathrm{SO}}_3\left(\mathrm{g}\right) \\ 2{\mathrm{SO}}_3\left(\mathrm{g}\right) \rightleftharpoons 2{\mathrm{SO}}_2 \left(\mathrm{g}\right) + {\mathrm{O}}_2\left(\mathrm{g}\right) \\ \mathrm{The} \mathrm{equilibrium} \mathrm{constants} \mathrm{are} \mathrm{related} \mathrm{as} \end{gathered}$$

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

• $2{\mathrm{K}}_1= {\mathrm{K}}_2^2$

• ${\mathrm{K}}_2=\frac{2}{{\mathrm{K}}_1^2}$

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

In the reversible reaction,

A(s) + B(g) $\rightleftharpoons$C (g) + D(g) ΔG^° = -350kJ

Which one of the following statements is true?

• The entropy change is negative

• Equilibrium constant is greater than one

• The reaction should be instantaneous

• The reaction is thermodynamically not feasible

Consider the folowing equilibria with equilibrium constants K₁ and K₂ respectively,

SO₂  (g) + $\frac{1}{2}{\mathrm{O}}_2 \left(\mathrm{g}\right) \rightleftharpoons$SO₃ (g)

2SO₃ (g) $\rightleftharpoons$ 2SO₂ (g) + O₂ (g)

The equilibrium constants are related as

• 2K₁ = K${}_2^2$

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

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

• ${\mathrm{K}}_2 = \frac{2}{{\mathrm{K}}_1^2}$

0.023 g of sodium metal is reacted with 100cm³ of water. The pH of the result solution is

• 10

• 11

• 9

• 12