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The reaction quotient (Q) formula looks very similar to the formula used for the equilibrium constant, K so what is their difference? Is it just that K has to be at equilibrium or is there other differences between them?
K is the equilibrium constant and can only be solved for when the concentrations of reactants and products at equilibrium are known. Q is the reaction quotient and can be solved at any point in the reaction, regardless if it is in equilibrium. If Q<K then the reaction moves forward and favors the products. If Q>K then the reaction moves reverse and favors the reactants.
Q is calculated the same was as K. However, K is the equilibrium constant while Q is the reaction quotient at any time during the reaction. You compare Q to K in order to find out whether the reaction is at equilibrium or not. When Q<K, the forward reaction is favored since the concentration of reactants is higher than the concentration of products. When Q>K, the reverse reaction is favored. If Q=K, the reaction is at equilibrium.
They look very similar because the only difference between the two constants is when you should use them. You use K when the reaction is at equilibrium, and you use Q at any stage of the reaction. You can compare the two constants to determine which way the reaction is shifting.
Q and K are actually calculated using the same formula with products over reactants. The difference between the two lies in when they apply. K ONLY applies to when a reaction has reached equilibrium. On the other hand, Q can represent any time during a reaction. Using Q, you an determine if the reaction is at equilibrium (K=Q) or what direction the reaction is moving in depending on if K>Q or K<Q.
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