K equilibrium vs Q equilibrium
Moderators: Chem_Mod, Chem_Admin
-
Rebekah_Park_2C
- Posts: 21
- Joined: Fri Sep 25, 2015 3:00 am
K equilibrium vs Q equilibrium
What is the difference between the K equilibrium and Q equilibrium? What do each variable stand for?
-
Joyce Xiong 4C
- Posts: 22
- Joined: Fri Sep 25, 2015 3:00 am
Re: K equilibrium vs Q equilibrium
Hey there!
In a chemical equilibrium, K stands for the equilibrium constant of a reaction, and Q stands for the reaction quotient. You calculate Q just like how you would calculate K, but the Q value is different (or else it would be called K). What happens here is that when you haven't waited long enough for the reaction to reach chemical equilibrium, the reaction is not in equilibrium and so it would have a different value when you try to solve for K. To further explain this, put yourself in this scenario: you are waiting for a reaction to reach equilibrium and you are given value K to begin with. After a while, you think it's about time that the reaction has reached equilibrium, so you calculate the ratio of products and reactions (you calculate "K") to confirm that the reaction has reached equilibrium. You correctly go through the procedure to find K, but you get a different value than you were given. Assuming you did everything right, this different value is Q, and it is different from K because this is the ratio of products over reactants at a instant when the reaction is not yet at equilibrium.
Furthermore, by knowing Q and K values, you can compare them to find out what direction a reaction will go towards (what reaction is favored/whether the reaction is going towards the reactants or the products). If Q<K, that means [Reactants]>[Products], and so in order to go towards equilibrium, a forward reaction toward the products is favored. On the other hand if Q>K, that means [Products]>[Reactants], and therefore a reverse reaction toward the reactants is favored, so that more reactants are formed to gradually reach the equilibrium ratio between reactants and products.
Hope this helps!
In a chemical equilibrium, K stands for the equilibrium constant of a reaction, and Q stands for the reaction quotient. You calculate Q just like how you would calculate K, but the Q value is different (or else it would be called K). What happens here is that when you haven't waited long enough for the reaction to reach chemical equilibrium, the reaction is not in equilibrium and so it would have a different value when you try to solve for K. To further explain this, put yourself in this scenario: you are waiting for a reaction to reach equilibrium and you are given value K to begin with. After a while, you think it's about time that the reaction has reached equilibrium, so you calculate the ratio of products and reactions (you calculate "K") to confirm that the reaction has reached equilibrium. You correctly go through the procedure to find K, but you get a different value than you were given. Assuming you did everything right, this different value is Q, and it is different from K because this is the ratio of products over reactants at a instant when the reaction is not yet at equilibrium.
Furthermore, by knowing Q and K values, you can compare them to find out what direction a reaction will go towards (what reaction is favored/whether the reaction is going towards the reactants or the products). If Q<K, that means [Reactants]>[Products], and so in order to go towards equilibrium, a forward reaction toward the products is favored. On the other hand if Q>K, that means [Products]>[Reactants], and therefore a reverse reaction toward the reactants is favored, so that more reactants are formed to gradually reach the equilibrium ratio between reactants and products.
Hope this helps!
Return to “Equilibrium Constants & Calculating Concentrations”
Who is online
Users browsing this forum: No registered users and 2 guests