Aim The purpose of this experiment is to find out how a system in equilibrium responds to a change in concentration of components in the mixture. Introduction Iron(III) ions and thiocyanate ions react in solution to produce thiocyanatoiron(III), a complex ion, according to the equation: Fe3+(aq) + SCN–(aq) → Fe(SCN)2+(aq) Pale yellow + colourless → blood-red The colour produced by the complex ion can indicate the position of equilibrium. Requirements
Procedure
Fe3+(aq) + 4Cl–(aq) → FeCl4–(aq) Interpretation of results Having made three observations, suggest a cause for each colour change (in terms of the concentrations of the coloured species) and then suggest what can be inferred about a shift in the position of equilibrium. If a pattern has emerged, then you can make a prediction based on the results of the experiment.
Questions
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Consider the exothermic equilibrium system below. The [CuCl4]–2(aq) ion is light green while the [CuBr4]2(aq) ion is dark brown. Originally the equilibrium below was a dark green. [CuCl4]–2(aq) + 4 Br–(aq) ⇋ [CuBr4]–2(aq) + 4 Cl–(aq) Predict the color of the solution after the system has re-established equilibrium. Adding a few drops of Ag+(aq) solution (Ag+ reacts with Cl– ions). Answer: COLOR OF SOLUTION? Adding a few drops of colorless HCl to the solution. Answer: COLOR OF SOLUTION? Consider the following system under equilibrium: \[ \underbrace{\ce{Fe^{3+}(aq)}}_{\text{colorless}} + \underbrace{ \ce{SCN^{-}(aq)}}_{\text{colorless}} \rightleftharpoons \underbrace{\ce{FeSCN^{2+}(aq)}}_{\text{red}} \nonumber \] If more \(Fe^{3+}\) is added to the reaction, what will happen? According to Le Chatelier's Principle, the system will react to minimize the stress. Since Fe3+ is on the reactant side of this reaction, the rate of the forward reaction will increase in order to "use up" the additional reactant. This will cause the equilibrium to shift to the right, producing more FeSCN2+. For this particular reaction, we will be able to see that this has happened, as the solution will become a darker red color. There are a few different ways to state what happens here when more Fe3+ is added, all of which have the same meaning:
What changes does this cause in the concentrations of the reaction participants?
How about the value of Keq? Notice that the concentration of some reaction participants have increased, while others have decreased. Once equilibrium has re-established itself, the value of Keq will be unchanged.
What if more FeSCN2+ is added? Again, equilibrium will shift to use up the added substance. In this case, equilibrium will shift to favor the reverse reaction, since the reverse reaction will use up the additional FeSCN2+.
How do the concentrations of reaction participants change?
Concentration can also be changed by removing a substance from the reaction. This is often accomplished by adding another substance that reacts (in a side reaction) with something already in the reaction. Let's remove SCN- from the system (perhaps by adding some Pb2+ ions—the lead(II) ions will form a precipitate with SCN-, removing them from the solution). What will happen now? Equilibrium will shift to replace SCN-—the reverse reaction will be favored because that is the direction that produces more SCN-.
How do the concentrations of reaction participants change?
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