Q<K

[Labiba Sardar 2A]

Can someone explain the relationship of why when Q<K, it means that more products are formed?

[Wilson 2E]

If Q is less than K, that means the Q ratio has a higher concentration of reactants due to the form [P]/[R], as a higher denominator produces a smaller number. To raise Q to the value of Kc, the amount of R must be reduced, which is done by creating more product, or favoring the right

[Norman Dis4C]

The reaction sits to the right when Q<K since there are more reactants than the products. To achieve an equilibrium state, more products are formed.

[Matt F]

The reaction quotient Q is similar to K since they both involve products over reactants ([P]/[R]), with the difference being that Q is taken at any point in the reaction while K is that ratio taken at equilibrium. Therefore, if you observe that Q<K, you know that there is a greater concentration of reactants (in the denominator, meaning that Q is smaller), which tells us that the forward reaction is favored to create more product

[IScarvie 1E]

When a reaction reaches equilibrium, no matter how many products or reactants you start with, you will always have the same ratio of products to reactants, K. So if you were to take a sample of the reaction before it reached equilibrium, measure the concentrations of reactants and products, and found the ratio Q was less than K, you'd know that the denominator/reactants of Q were larger than the denominator/reactants of K, and/or there is a smaller concentration of products in Q than K. Either assumption shows that the reaction still needs to convert more reactants to products to reach equilibrium, so the reaction is still moving forward (converting reactants to products).