Hydrogen ion concentration in mol/L in a solution of pH = 5.4 will be

• 3.98 × 10⁸

• 3.88 × 10⁶

• 3.68 × 10^-6

• 3.98 × 10^-6

For the following equilibrium, N₂O₄ $\rightleftharpoons$2NO₂ in the gaseous phase, NO₂ is 50% of the total volume when equilibrium is set up. Hence, per cent of dissociation of N₂O₄ is

• 50%

• 25%

• 66.66%

• 33.33%

Of the given anions, the strongest Bronsted base is

• ClO^-

• ClO${}_3^{-}$

• ClO${}_2^{-}$

• ClO${}_4^{-}$

At equilibrium, if K_P = 1, then

• $∆\mathrm{G}°$ > 1

• $∆\mathrm{G}°$ < 1

• $∆\mathrm{G}°$ = 0

• $∆\mathrm{G}°$ = 1

375.

Which of the following is most acidic?

• H₂O

• H₂S

• H₂Se

• H₂Te

1 mL of 0.01 N HCl is added to 999 mL solution of 0.1 N Na₂SO₄. The pH of the resulting solution will

• 2

• 7

• 5

• 1

The equilibrium constants for the reaction,

Br₂ $\rightleftharpoons$ 2Br → 1

at 500 K and 700 K are 1 × 10^-10 and 1 × 10^-5 respectively. The reaction is

• endothermic

• exothermic

• fast

• slow

A chemist wishes to prepare a buffer solution of pH = 2.90 that efficiently resists a change in pH yet contains only small concentration of buffering agents which one of the following weak acid along with its salt would be best to use 

• m-chlorobenzoic acid (pK_a = 3.98)

• Acetoacetic acid (pK_a = 3.58)

• 2 5-dihydrobenzoic acid (pK_a = 2.97)

• p-chlorocmanic acid (pK_a = 4.41)

NaOH is a strong base. What will the be pH of 5.0 × 10^-2 M NaOH solution? (log 2 = 0.3)

• 14.00

• 13.70

• 13.00

• 12.70

The equilibrium constants of the reactions

${\mathrm{SO}}_2\left(\mathrm{g}\right) + \frac{1}{2}{\mathrm{O}}_2\left(\mathrm{g}\right) \rightleftharpoons {\mathrm{SO}}_3\left(\mathrm{g}\right)$ and $2{\mathrm{SO}}_2\left(\mathrm{g}\right) + {\mathrm{O}}_2\left(\mathrm{g}\right) \rightleftharpoons 2{\mathrm{SO}}_3\left(\mathrm{g}\right)$ are K₁ respectively. The relationship between K₁ and K₂ will be

• ${\mathrm{K}}_1^2=$K₂

• K₂ = √K₁

• K₁= K₂

• ${\mathrm{K}}_2^3 = {\mathrm{K}}_1$