What you are getting from a statement might not be necessarily true. Let's burst some of the common misconceptions.

If we change the concentration of either reactants or products, the value of equilibrium constant also changes.
Suppose you have an equilibrium established between four substances A, B, C and D

Let's say we decrease the concentration of 'C', technically the value of $$K_c$$ will reduce according to the above expression.
But, if you know Le Chatelier's principle, it states that when a system experiences a disturbance (such as concentration, temperature, or pressure changes), it will respond to restore a new equilibrium state.
This means that the system subjected to a change will try to move in a direction that reduces the effect of the change (image below).
When all this takes place, the value of $$K_c$$ does not change.

$$ \Delta n$$ in the expression $$K_p= K_c (RT)^{\Delta n}$$ is equal to Sum of coefficient of product - Sum of coefficient of reactants
This misconception leads to mistakes in solving problems related to $$K_p$$ and $$K_c$$. So, let's clear it for once and all.
Consider this reaction,
$$P_4(s) +6Cl_2 (g) \rightleftharpoons 4PCl_3(g)$$
According to the expression, $$\Delta n$$ = Sum of coefficient of product - Sum of coefficient of reactants
$$ \Delta n$$= 4-(6+1) = -3
This is not the correct value. But why?
Actually, $$\Delta n$$= Sum of coefficient of gaseous product - Sum of coefficient of gaseous reactants
Let's calculate the value of$$ \Delta n$$ again
As $$PCl_3$$ and $$Cl_2$$ are the gaseous components in the reaction, we will consider their coefficients which are 4 and 6 respectively.
Therefore, the correct value of $$ \Delta n$$ = 4 - 6= -2

Once a reaction reaches equilibrium, it stops.
Let's tackle this problem graphically.

When we study the graph, we see that the product concentration and reactant concentration becomes constant at equilibrium. Does it mean mean that the reactants are not changing into products and vice versa?

If there wasn't any reaction taking place, then rate of reaction would've been zero. But we can see that it is not. Instead we find the rate of forward reaction to be equal to the rate of backward reaction. This means as soon as reactants convert into products, they react amongst themselves to give back the reactants.
This makes it seem like the reaction is not taking place, when in reality there is a continuous reaction taking place.

Equilibrium