This is an incomplete question, here is a complete question.
Consider the following equilibrium at 100°C.
Concentration at equilibrium:
![[COBr_2]=1.58\times 10^{-6}M](https://tex.z-dn.net/?f=%5BCOBr_2%5D%3D1.58%5Ctimes%2010%5E%7B-6%7DM)
![[Co]=2.78\times 10^{-3}M](https://tex.z-dn.net/?f=%5BCo%5D%3D2.78%5Ctimes%2010%5E%7B-3%7DM)
![[Br_2]=2.51\times 10^{-5}M](https://tex.z-dn.net/?f=%5BBr_2%5D%3D2.51%5Ctimes%2010%5E%7B-5%7DM)
If a system has a reaction quotient of 2.13 × 10⁻¹⁵ at 100°c, what will happen to the concentrations of COBr₂, Co and Br₂ as the reaction proceeds to equilibrium?
Answer : The concentrations of Co and Br₂ decreases and the concentrations of COBr₂ increases.
Explanation :
Reaction quotient (Q) : It is defined as the measurement of the relative amounts of products and reactants present during a reaction at a particular time.
The given balanced chemical reaction is,
The expression for reaction quotient will be :
![Q=\frac{[CO][Br_2]}{[COBr_2]}](https://tex.z-dn.net/?f=Q%3D%5Cfrac%7B%5BCO%5D%5BBr_2%5D%7D%7B%5BCOBr_2%5D%7D)
In this expression, only gaseous or aqueous states are includes and pure liquid or solid states are omitted.
Now put all the given values in this expression, we get
The given equilibrium constant value is,
Equilibrium constant : It is defined as the equilibrium constant. It is defined as the ratio of concentration of products to the concentration of reactants.
There are 3 conditions:
When 
that means product > reactant. So, the reaction is reactant favored.
When 
that means reactant > product. So, the reaction is product favored.
When 
that means product = reactant. So, the reaction is in equilibrium.
From the above we conclude that, the that means product < reactant. So, the reaction is product favored that means reaction must shift to the product (right) to be in equilibrium.
Hence, the concentrations of Co and Br₂ decreases and the concentrations of COBr₂ increases.
