Given the data from an experiment, determine the rate constant (k_f) for the reaction.

Step 3: Calculate the value of $k_f$.
The rate law is $Rate = -k_f C_A^m C_B^n$.
From the previous steps, we found $m=1$ and $n=\frac{1}{2}$.
So, the rate law is $Rate = -k_f C_A^1 C_B^{1/2}$.
We can use the data from any experiment to find $k_f$. Let's use Experiment 1:
$C_A = 1 \text{ mol/L}$
$C_B = 1 \text{ mol/L}$
$Rate = -2.45 \times 10^{-4} \text{ mol/L.s}$

Substitute these values into the rate law:

$$-2.45 \times 10^{-4} \text{ mol/L.s} = -k_f (1 \text{ mol/L})^1 (1 \text{ mol/L})^{1/2}$$ $$-2.45 \times 10^{-4} \text{ mol/L.s} = -k_f (1 \text{ mol/L}) (1 \text{ mol/L})^{1/2}$$ $$-2.45 \times 10^{-4} \text{ mol/L.s} = -k_f (1 \text{ mol/L})^{3/2}$$

Now, solve for $k_f$:

$$k_f = \frac{2.45 \times 10^{-4} \text{ mol/L.s}}{(1 \text{ mol/L})^{3/2}}$$ $$k_f = 2.45 \times 10^{-4} \text{ mol}^{-1/2} \text{ L}^{1/2} \text{ s}^{-1}$$

The value of $k_f$ is $\boxed{2.45 \times 10^{-4} \text{ mol}^{-1/2} \text{ L}^{1/2} \text{ s}^{-1}}$.