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The rate of a first order reaction is $1.5 \times {10^{ - 2}}mol{L^{ - 1}}{\min ^{ - 1}}$ at $0.5M$ concentration of the reactant.

The half life of the reaction is:

A. $0.383\min $

B. $23.1\min $

C. $8.73\min $

D. $7.53\min $

The half life of the reaction is:

A. $0.383\min $

B. $23.1\min $

C. $8.73\min $

D. $7.53\min $

In a zero order reaction for every $10 ^\circ $ rise in temperature,the rate is doubled.If the temperature is increased from $10 ^\circ C$ to $100 ^\circ C$ ,the rate of reaction will become

a) $256$ times

b) $512$ times

c) $64$ times

d) $128$ times

a) $256$ times

b) $512$ times

c) $64$ times

d) $128$ times

What do you understand by the rate of a reaction? Explain why the rate of a reaction cannot be measured by dividing the total amount of reactant consumed by the time taken.

(a) What do you mean by rate of a reaction?

(b) What will be the effect of temperature on the rate of a reaction?

(b) What will be the effect of temperature on the rate of a reaction?

If the rate respect to ${{\text{O}}_{2\;\;}}$, $NO$, $N{O_2}$ are $ - \dfrac{{\vartriangle \left[ {{{\text{O}}_{2\;\;}}} \right]}}{{\vartriangle {\text{t}}}}\; = \; - \dfrac{1}{2}\dfrac{{\vartriangle \left[ {NO} \right]}}{{\vartriangle {\text{t}}}}\;\; = \; + \dfrac{1}{2}\dfrac{{\vartriangle \left[ {N{O_2}} \right]}}{{\vartriangle {\text{t}}}}$ , then the corresponding chemical equation is $2NO\; + \;{{\text{O}}_2}\; \to \;N{O_2}$.

A.True

B.False

A.True

B.False

1) For the reaction \[R \to P\], the concentration of the reactant changes from 0.03M to 0.02M in 25 minutes. Calculate the average rate of reaction using units of time both in minutes and seconds.

(2) In a reaction, \[2A \to P\] products, the concentration of A decreases from 0.5 to 0.4 mol \[{L^{ - 1}}\]in 10 minutes. Calculate the rate during this interval.

(2) In a reaction, \[2A \to P\] products, the concentration of A decreases from 0.5 to 0.4 mol \[{L^{ - 1}}\]in 10 minutes. Calculate the rate during this interval.

A radioactive material decays by simultaneous emission of two particles with respective half-lives 1620 and 810 years. The time, in years, after which one-fourth of the material remains is:

A. 1080

B. 2430

C. 3240

D. 4860

A. 1080

B. 2430

C. 3240

D. 4860

In the following reaction:

\[{\text{xA}} \to {\text{yB}}\]

\[{\text{log [}}\dfrac{{{\text{ - d[A]}}}}{{{\text{dt}}}}{\text{] = log[}}\dfrac{{{\text{d[B]}}}}{{{\text{dt}}}}{\text{] + log2}}\]

Where –ve sign indicates rate of disappearance of the reactant. Thus x:y is :

\[1)1:2\]

\[2)2:1\]

\[3)3:1\]

\[4)3:10\;\]

\[{\text{xA}} \to {\text{yB}}\]

\[{\text{log [}}\dfrac{{{\text{ - d[A]}}}}{{{\text{dt}}}}{\text{] = log[}}\dfrac{{{\text{d[B]}}}}{{{\text{dt}}}}{\text{] + log2}}\]

Where –ve sign indicates rate of disappearance of the reactant. Thus x:y is :

\[1)1:2\]

\[2)2:1\]

\[3)3:1\]

\[4)3:10\;\]

The decomposition of dimethyl ether leads to the formation of \[C{{H}_{4}}\] , \[{{H}_{2}}\] and \[CO\] and the reaction rate is given by

Rate = \[k{{\left[ C{{H}_{3}}OC{{H}_{3}} \right]}^{\dfrac{3}{2}}}\]

The rate of reaction is followed by increase in pressure in a closed vessel, so the rate can also be expressed in terms of the partial pressure of dimethyl ether, i.e.,

Rate = \[k{{\left( {{p}_{C{{H}_{3}}OC{{H}_{3}}}} \right)}^{\dfrac{3}{2}}}\]

If the pressure is measured in bar and time in minutes, then what are the units of rate and rate constants?

Rate = \[k{{\left[ C{{H}_{3}}OC{{H}_{3}} \right]}^{\dfrac{3}{2}}}\]

The rate of reaction is followed by increase in pressure in a closed vessel, so the rate can also be expressed in terms of the partial pressure of dimethyl ether, i.e.,

Rate = \[k{{\left( {{p}_{C{{H}_{3}}OC{{H}_{3}}}} \right)}^{\dfrac{3}{2}}}\]

If the pressure is measured in bar and time in minutes, then what are the units of rate and rate constants?

What will be the instantaneous rate of reaction in terms of its reactants for the given equation?

\[2KCl{O_3} \to 2KCl + 3{O_2}\]

\[2KCl{O_3} \to 2KCl + 3{O_2}\]

How many types of reaction are there on the basis of speed\[?\]

What is the order of a chemical reaction \[A+2B\text{ }\xrightarrow{k}\text{ }C\], if the rate of formation of C increases by a factor of \[2.82\] on doubling the concentration of A and increases by a factor of \[9\] on tripling the concentration of B?9

A.$7/2.$

B.$7/4.$

C.$5/2.$

D.$5/4.$

A.$7/2.$

B.$7/4.$

C.$5/2.$

D.$5/4.$

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