Q.If $K_b$ and $K_f$ for a reversible reaction are $0.8 \times 10^{-5}$ and $1.6 \times 10^{-4}$ respectively, the value of the equilibrium constant is,

For a general reversible reaction $xA+yB\rightleftharpoons lC+mD$, the equilibrium constant is the ratio, at a fixed temperature, of the product of the active masses of the products (each raised to its own stoichiometric coefficient) to the product of the active masses of the reactants (likewise raised to their coefficients):

When every species in the reaction is gaseous, the same equilibrium can equally be described using partial pressures instead of concentrations:

$K_C$ (or $K_P$) is a genuine constant at a given temperature -- it does not change if the starting concentrations are changed (doubling the initial concentrations of A and B in $A+B\rightleftharpoons C$ changes the equilibrium composition the system settles into, but not the numerical value of $K_C$ itself, since $K_C$ depends only on temperature). It changes only if the temperature changes (the relationship is made quantitative by the Van't Hoff equation).

This expression is applied directly to compute $K_C$ from a set of measured equilibrium concentrations -- for example, for $2A(g)\rightleftharpoons 2B(g)+C_2(g)$ with $[A]=1\times10^{-4}$, $[B]=2.0\times10^{-3}$, $[C_2]=1.5\times10^{-4}$ M at 400 K, $K_C=\dfrac{[B]^2[C_2]}{[A]^2}=\dfrac{(2.0\times10^{-3})^2(1.5\times10^{-4})}{(1\times10^{-4})^2}=0.06$. The same expression also lets a balanced equation be reconstructed purely from a stated $K_C$ formula, by reading off which species (and what power) sit in the numerator (products) and denominator (reactants) -- for instance $K_C=\dfrac{[NH_3]^4[O_2]^5}{[NO]^4[H_2O]^6}$ corresponds to the balanced equilibrium $4NO(g)+6H_2O(g)\rightleftharpoons 4NH_3(g)+5O_2(g)$.

Equilibrium-constant algebra of this kind extends naturally to a heterogeneous solubility-type equilibrium too -- for $Fe(OH)_3(s)\rightleftharpoons Fe^{3+}(aq)+3OH^-(aq)$, the constant $[Fe^{3+}][OH^-]^3$ stays fixed at a given temperature, so if $[OH^-]$ is deliberately lowered to one-quarter of its previous value, $[Fe^{3+}]$ must rise by a factor of $4^3=64$ to keep the product constant.