If k = 0.3 1/Mxs and the concentration [A] = 0.035 M, what is the calculated rate for the reaction?

If k = 0.3 1/Mxs and the concentration [A] = 0.035 M, what is the calculated rate for the reaction?

• A. -3.675 x 10^-4 M/s
• B. -3.65 x 10^-4 M/s
• C. -3.70 x 10^-4 M/s
• D. -3.680 x 10^-4 M/s

Answer: (A)

To calculate the rate of the reaction, we use the rate law formula. The rate of a reaction is often expressed in terms of the rate constant \( k \) and the concentration of the reactant \( [A] \). In this case, the relationship given by the rate law can be expressed as follows: \[ \text{Rate} = k \cdot [A] \] Given that \( k = 0.3 \, \text{1/Mxs} \) and \( [A] = 0.035 \, \text{M} \), we can substitute these values into the equation: \[ \text{Rate} = 0.3 \, \text{1/Mxs} \times 0.035 \, \text{M} \] Carrying out the multiplication: \[ \text{Rate} = 0.3 \times 0.035 = 0.0105 \, \text{M/s} \] However, we must account for the units of the rate constant and the concentration to find the rate in M/s correctly. Given that the rate law typically involves negative signs when the concentration of a reactant decreases (assuming a decrease in

To calculate the rate of the reaction, we use the rate law formula. The rate of a reaction is often expressed in terms of the rate constant ( k ) and the concentration of the reactant ( [A] ). In this case, the relationship given by the rate law can be expressed as follows:

[

\text{Rate} = k \cdot [A]

]

Given that ( k = 0.3 , \text{1/Mxs} ) and ( [A] = 0.035 , \text{M} ), we can substitute these values into the equation:

[

\text{Rate} = 0.3 , \text{1/Mxs} \times 0.035 , \text{M}

]

Carrying out the multiplication:

[

\text{Rate} = 0.3 \times 0.035 = 0.0105 , \text{M/s}

]

However, we must account for the units of the rate constant and the concentration to find the rate in M/s correctly. Given that the rate law typically involves negative signs when the concentration of a reactant decreases (assuming a decrease in