In a reaction,the concentration of reactant is increased two times and three times,then the increases in the rate of reaction were four times and nine times respectively. The order of reaction is:

Consider the reaction,$2A + B \rightarrow$ products. When concentration of $B$ alone was doubled,the half-life did not change. When the concentration of $A$ alone was doubled,the rate increased by two times. The unit of rate constant for this reaction is

When the concentration of a reactant,$A$,in a reaction $A \to \text{Products}$ is doubled,the rate of reaction increases seven times. The order of the reaction is between:

The hypothetical reaction $A_2 + B_2 \rightarrow 2AB$ follows the mechanism given below. What is the overall order of the reaction?
$A_2 \rightleftharpoons A + A$ ...... (fast);
$A + B_2 \rightarrow AB + B$ ...... (slow);
$A + B \rightarrow AB$ ...... (fast)

For the reaction $2A + B \rightarrow C + D$,select the correct rate law based on the following data:
$1$. $[A] = 0.1, [B] = 0.1, \text{Initial Rate} = 7.5 \times 10^{-3}$
$2$. $[A] = 0.3, [B] = 0.2, \text{Initial Rate} = 9.0 \times 10^{-2}$
$3$. $[A] = 0.3, [B] = 0.4, \text{Initial Rate} = 3.6 \times 10^{-1}$
$4$. $[A] = 0.4, [B] = 0.1, \text{Initial Rate} = 3.0 \times 10^{-2}$

Which of the following statements regarding the order of a reaction is incorrect?