Question

The table below shows recorded concentration data for the following chemical reaction:
$CaC{l}_2\left(aq\right)+2AgN{O}_3\left(aq\right)\to Ca\left(N{O}_3{)}_2\right(aq)+2AgCl(s)$
Use the data to determine the exponents $x$ and $y$ in the rate law:
$rate=k\left[CaC{l}_2{]}^x\right[AgN{O}_3{]}^y$

$\left[CaC{l}_2\right]M$	$\left[AgN{O}_3\right]M$	Rate M/s
$0.050$	$0.050$	$3.33\times {10}^{-3}$
$0.100$	$0.050$	$6.66\times {10}^{-3}$
$0.100$	$0.100$	$13.32\times {10}^{-3}$

• A. $x=2,y=3$
• B. $x=1,y=2$
• C. $x=1,y=1$
• D. $x=2,y=2$

Answer 1

The correct option is C $x=1,y=1$

On making the concentration of $CaC{l}_2$ double with no change in concentration of $AgN{O}_3$ rate of reaction also increases 2 times which implies that x=1

And on making concentration of both the reactants double rate increases by 4 times and since x=1 it can be concluded that y=1