Many chemical reactions do not go to completion. Rather, they come to a point of chemical equilibrium before the reactants are fully converted to products. At the point of equilibrium, the concentrations of all reactants remain constant with time. In this experiment, you will investigate how outside forces acting on a system at equilibrium provoke changes within the system (Le Chatelier’s principle).
Early in the study of chemical reactions, it was noted that many chemical reactions do not produce as much product as might be expected, based on the amounts of reactants taken originally. These reactions appeared to have stopped before the reaction was complete. Closer examination of these systems (after the reaction had seemed to stop) indicated that there were still significant amounts of all the original reactants present. Quite naturally, chemists wondered why the reaction had seemed to stop, when all the necessary ingredients for further reaction were still present. Some reactions appear to stop because the products produced by the original reaction themselves begin to react, in the reverse direction to the original process. As the concentration of the products begins to build up, product molecules will react more and more frequently. Eventually, as the speed of the forward reaction decreases while the speed of the reverse reaction increases, the forward and reverse processes will be going on at exactly the same rate. Once the forward and reverse rates of reaction are identical, there can be no further net change in the concentrations of any of the species in the system, at this point, a dynamic state of equilibrium has been reached. The original reaction is still taking place but is opposed by the reverse of the original reaction also taking place. In this experiment, you will study changes made in a system already in equilibrium, by the reference to Le Chatelier’s principle. Le Chatelier’s principle states that, if we disturb a system that is already in equilibrium, then the system will react so as to minimize the effect of the disturbance. This is most easily demonstrated in cases where additional reagent is added to a system in equilibrium, or when one of the reagents is removed from the system in equilibrium.
Suppose we have a solution that has been saturated with a solute: This means that the solution has already dissolved as much solute as possible. If we try to dissolve additional solute, no more will dissolve because the saturated solution is in equilibrium with the solute:
Solute + Solvent ( Solution
Le Chatelier’s principle is most easily seen when an ionic solute is used: Suppose we have a saturated solution of sodium chloride, NaCl. Then
NaCl ( Na+ (aq) + Cl- (aq)
will describe the equilibrium that exists. Suppose we then try adding an additional amount of one of the ions involved in the equilibrium: For example, suppose we added several drops of concentrated HCl solution (which contains the chloride ion at high concentration). According to Le Chatelier’s principle, the equilibrium will shift so as to consume some of the added chloride ion. This would result in a net decrease in the amount of NaCl that could dissolve. If we watched the saturated NaCl solution as the HCl was added, we should see some of the NaCl precipitate as a solid.
Complex Ion Equilibria
Oftentimes, dissolved metal ions will react with certain substances to produce brightly colored species called complex ions. For example, ion (III)reacts with the thiocyanate ion (SCN-) to produce a bright red complex ion:
Fe3+ + SCN- ( [FeNCS2+]
This is an equilibrium process that is easy to study because we can monitor the bright red color of [FeNCS2+] as an indication of the position of the equilibrium: If the solution is very red, there is a lot of [FeNCS2+]…