Question

The rate law for the reaction, \(\text{A}+\text{B}\rightarrow \text{P}\) is rate \(=k[\text{A}][\text{B}]\). The rate constant of the reaction is \(4.5\ \text{M}^{-1}\text{s}^{-1}\) at \(25^{\circ }\)C. Calculate the rate of the reaction when \([\text{A}]=0.1\) M and \([\text{B}]=0.2\) M.

• A. \(9.25\times 10^{-2}\ \text{M s}^{-1}\)
• B. \(9.0\times 10^{-2}\ \text{M s}^{-1}\)
• C. \(9.50\times 10^{-2}\ \text{M s}^{-1}\)
• D. \(9.75\times 10^{-2}\ \text{M s}^{-1}\)

Hint:

Substitute the rate constant and the two concentrations into the rate law.

Answer 1

The Correct Option is B

Step 1: Understanding the Concept
The rate law is rate \(=k[\text{A}][\text{B}]\) with \(k=4.5\ \text{M}^{-1}\text{s}^{-1}\).

Step 2: Substitute
\([\text{A}]=0.1\) M and \([\text{B}]=0.2\) M.

Step 3: Calculation
Rate \(=4.5\times0.1\times0.2=9.0\times10^{-2}\ \text{M s}^{-1}\).

Final Answer:
The rate is 9.0 times 10 to the power minus 2 M per second. \[ \boxed{\text{(B) }9.0\times10^{-2}\ \text{M s}^{-1}} \]