In my textbook it is said that changing concentration, pressure and temperature all cause a shift in equilibrium. I can understand that a change in pressure and temperature can cause a change in proportion of reactants to products, however i am unsure of the effect of changing concentration on equilibrium. In the explanation of the effect of concentration on equilibrium position a reaction A + B (goes to) C + D is used. It is said that if we add some extra A, the concentration of A increases and the rate of the forward reaction increases and equilibrium shifts to the right. Which is all fine. However, the textbook then says that "We end up with a greater proportion of products in the reaction than before we added A". I cannot believe this is correct, the way I see it, when A is added there is a greater proportion of reactants to products, and equilibrium shifts to the left, but the equilibrium then shifts to the right, until it reaches where it was before A was added and the proportion of reactants to products is the same as before A was added. Am I correct in believing the textbook is at error?
Thanks a lot
Le Chatelier's Principle Watch
- Thread Starter
- 15-05-2013 14:50
- 15-05-2013 15:14
i think it is getting at the fact that the higher initial concentraton of reactant means there is a higher ratio of reactants to products, so more products are made.
that leads to more producs that reactants, so the equiliburium shifts back to the left,
so eventually you get the same ratios as before, just the concentratiuons of each chemical/compound is greater
that or that an equilirium will more to the side with less mols of chemicals used
H2O < > H+ + OH-
this is how water acts as a weak acid,
there are more mols of product (H+ and OH-)=2 than reactant (H2O)=1
so the equilibrium shifts to the left sore more H20 is produced
hope this helps