The color of chemistry pre lab answers

CHEM 404 Exp 5 Experiment 5: CHEMICAL EQUILIBRIUM – LE CHÂTELIER’S PRINCIPLE* When working in the laboratory, one often makes observations that at first sight are surprising and hard to explain. One might add a reagent to a solution and obtain a precipitate. Addition of more of that reagent to the precipitate causes it to dissolve. A violet solution turns yellow on addition of a reagent. Subsequent addition of another reagent brings back first a green solution and then the original violet one. Clearly, chemical reactions are occurring, but how and why they behave as they do is not at once obvious. In this experiment, you will examine and attempt to explain several observations of the sort just mentioned. Central to your explanation will be recognition of the fact that chemical systems tend to exist in a state of equilibrium. If one disturbs the equilibrium in one way or another, the reaction may shift to the left or right, producing the kinds of effects mentioned above. If one can understand the principles governing the equilibrium system, it is often possible to see how one might disturb the system, such as by adding a particular reagent or heat, and so cause it to change in a desirable way. Before proceeding to specific examples, a general system can be examined, noting the key principle that allows controlled changes to be made to a system in equilibrium. Consider the reaction A(aq) ⇌ B(aq) + C(aq) (1) where A, B, and C are molecules or ions in solution. If a mixture of these species is in equilibrium, their concentrations are related. There is a condition that those concentrations must meet, namely that Kc = [B] x [C] [A] (2) where Kc is a constant, called the equilibrium constant for the reaction. For a given reaction at any given temperature, Kc has a particular value. For example, for a given solution in which Reaction 1 can occur, substituting the equilibrium values for the molarities of A, B, and C into Equation 2, could give a value of 10 for Kc. Now, suppose more of species A is added to that solution. What will happen? Remember, Kc can’t change. If a new higher molarity of A is substituted into Equation 2, the value would be smaller than Kc. This means that the system is not in equilibrium, and must change in some way to get back to equilibrium. How can it do this? It can do this by shifting to the right, producing more B and C and using up some A. It must do this, and will, until the molarities of C, B, and A reach values that, on substitution into Equation 2, again equal 10. At that point, the system is once again in equilibrium. In this new equilibrium state, [B] and [C] will be greater than they were initially, and [A] will be larger than its initial value, but smaller than if there had been no forced shift to the right. The conclusion you should reach on reading the last paragraph is that one can always cause a reaction to shift to the right by increasing the concentration of a reactant.