The rate law for the reaction $A + B \rightarrow \text{product}$ is $\text{rate} = k[A][B]$. When will the rate of reaction increase by a factor of $2$?

For the reaction $2A + B \rightarrow C$,the rate law is given by: $\text{Rate} = k[A][B]$. Which of the following statements is correct regarding this reaction?

The rate constant of a reaction is $1.388 \times 10^{-3} \ mol^{-2} \ L^{2} \ s^{-1}$. The order of the reaction is:

For a reaction,$AB_5 \to AB + 4B$,the rate can be expressed in the following ways:
$-\frac{d[AB_5]}{dt} = K[AB_5]$ ; $\frac{d[B]}{dt} = K_1[AB_5]$
What is the correct relation between $K$ and $K_1$?

Consider the following reactions:
$A$ $\rightarrow P_1; B$ $\rightarrow P_2; C$ $\rightarrow P_3; D$ $\rightarrow P_4$
The orders of the above reactions are $a, b, c,$ and $d,$ respectively. The following graph is obtained when $\log[\text{rate}]$ vs. $\log[\text{conc.}]$ are plotted:
Among the following,the correct sequence for the order of the reactions is:

Assertion : The kinetics of the reaction $mA + nB + pC \to m'X + n'Y + p'Z$ obey the rate expression as $\frac{dX}{dt} = k[A]^m[B]^n$.
Reason : The rate of the reaction does not depend upon the concentration of $C$.