Equilibrium constant when adding more of a reactant Suppose I have the reversible reaction:

$$\ce{A +B⇌ C} $$

The reaction is at equilibrium with equilibrium constant $K$

I am told that if I increase the concentration of $\ce{B}$, the rate for the forwards reaction will exceed the backwards one. Fair enough. 

I am also told that $K$ will necessarily increase. Why though? Its true that Forward reaction > back ward reaction until we reach a new equilibrium such that more of $\ce{C}$ is produced but I don't see why this implies in any way that the final quotient $\frac{[\ce{C}]}{\ce{[A][B]}}$ will necessarily be any greater. 

Certainly if we have a simple reaction:

$$\ce{A⇌ C} $$

and we add more of $\ce{A}$, then the equilibrium constant for the new final final state will remain as it was, ceteris paribus. 

What am I getting wrong, because my textbook suggests that $K$ will always increase, no matter what type of reaction I am dealing with (of course, as long as all reactants are in a suspended form, e.g. dissolved or gaseous and the reaction is subject to Le Châtelier's principle).

Please, if possible, keep the answer as a level intelligible to a high-school student. I don't know much about advanced university chemistry.
You're correct.  The equilibrium constant will not change.  The equilibrium constant is only defined at equilibrium.  Your analysis of the situation was flawless.