The rate of a first order reaction is 1.5 × 10^-2 ~mol ~L^-1 ~min^-1 at 0.5 M concentration of the reactant. The half life of the reaction is:

For the reaction 2SO₂(g) + O₂(g) ⇌ 2SO₃(g), ΔH = -198 kJ. Which condition favours forward reaction?

The unit of rate constant for a first-order reaction is:

$\begin{aligned} &\text { Consider the following compounds: }\\ &\mathrm{\underline{K}O}_2, \mathrm{H}_2 \mathrm{\underline{O}}_2 \text { and } \mathrm{H}_2 \mathrm{\underline{S}O}_4 \text {. } \end{aligned}$ The oxidation states of the underlined elements in them are, respectively,

If the molar conductivity $\left(\Lambda_{\mathrm{m}}\right)$ of a $0.050 \mathrm{~mol} \mathrm{~L}^{-1}$ solution of a monobasic weak acid is $90 \mathrm{~S} \mathrm{~cm}^2 \mathrm{~mol}^{-1}$, its extent (degree) of dissociation will be [Assume $\Lambda_{+}^{\circ}=349.6 \mathrm{~S} \mathrm{~cm}^2 \mathrm{~mol}^{-1}$ and $\left.\Lambda_{-}^{\circ}=50.4 \mathrm{~S} \mathrm{~cm}^2 \mathrm{~mol}^{-1}.\right]$

If the rate constant of a reaction is $0.03 \mathrm{~s}^{-1}$, how much time does it take for $7.2 \mathrm{~mol} \mathrm{~L}^{-1}$ concentration of the reactant to get reduced to $0.9 \mathrm{~mol} \mathrm{~L}^{-1}$ ? (Given $: \log 2=0.301$ )