For the reaction A + B ⇌ C + D, if K is 0.1, what does this imply regarding the concentrations of C and D compared to A and B?

When the equilibrium constant (K) for the reaction A + B ⇌ C + D is given as 0.1, it reflects the ratio of the concentrations of the products (C and D) to the concentrations of the reactants (A and B) at equilibrium. The expression for the equilibrium constant in this case is:

K = [C][D] / [A][B]

A value of K = 0.1 indicates that the ratio of the products' concentrations ([C][D]) is much smaller compared to that of the reactants' concentrations ([A][B]). This suggests that, at equilibrium, the concentrations of the reactants (A and B) will be significantly higher than those of the products (C and D). Therefore, it can be concluded that the reaction favors the formation of the reactants under the given conditions, leading to the scenario where concentrations of A and B are greater than those of C and D.

This understanding of the equilibrium constant provides insight into the position of equilibrium for a reaction, indicating whether the reaction favors products or reactants. In this case, the low value of K strongly suggests that the equilibrium lies towards the left, favoring the reactants.