Chemical equilibrium constant We define chemical equilibrium as the point in a reaction where the rates of the forward and reverse reactions become equal. We also write 

$$K_c = \frac{k_f}{k_r}$$

If at equilibrium the rate of the forward reaction and the rate of the reverse reaction become equal then their rate constants should also be same. If rate constants $k_f$ and $k_r$  become the same, then $K_c$ should always be equal to 1. However, the value of $K_c$ is not always 1. Why is it so?
For an example reaction $\ce{A + B &lt;=&gt; C + D}$

the forward rate is given by

$$r_\mathrm{f} = k_\mathrm{f}[\ce{A}][\ce{B}]$$

and the backward rate is given by

$$r_\mathrm{b} = k_\mathrm{b}[\ce{C}][\ce{D}]$$

For the rates to be equal one simply needs

$$k_\mathrm{f}[\ce{A}][\ce{B}] = r_\mathrm{f} = r_\mathrm{b} = k_\mathrm{b}[\ce{C}][\ce{D}]$$

Note that this does not imply that the two rate constants are the same

$$k_\mathrm{f} \neq k_\mathrm{b} \quad \text{(in general)}$$