For the reaction, A(g) + B(g) => 2 C(g), the following data were obtained at constant temperature.
Experiment  Initial [A], mol/L  Initial [B], mol/L  Initial Rate, M/min 
1  0.10  0.10  2 x 10^{4} 
2  0.30  0.30  5.4 x 10^{3} 
3  0.10  0.30  1.8 x 10^{3} 
4  0.20  0.40  6.4 x 10^{3} 
Which of the following is the correct rate law for the reaction?
1. 
Rate = k[A][B] 

2. 
Rate = k[A]^{2}[B] 

3. 
Rate = k[A]^{2}[B]^{2} 

4. 
Rate = k[A] 

5. 
Rate = k[A][B]^{2} 
Solution :
Lets calculate the order of each reactant
Calculating the order of A using the data from experiment 2 and 3
Rate 3/ rate 2 = ([A]3/[A]2)^m
1.8*10^3/5.4*10^3 = [0.10/0.30]^m
0.33 = 0.33 ^m
Log 0.33 = m * log 0.33
Log 0.33 / log 0.33 = m
1=m
So order of A = 1
Now lets calculate the order with B using the data from experiment 1 and 3
Rate 3 / rate 1 = ([B]3/[B]1)^n
1.8*10^3/2*10^4 = [0.30/0.10]^n
9=3^n
Log 9 = n* log 3
Log 9 / log 3 = n
2=n
So order of B is 2
So the rate law for the reaction is
Rate = K[A][B]^2
So the correct answer is option 5
Get Answers For Free
Most questions answered within 1 hours.