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Laboratory 1: Chemical Equilibrium^1
Reading: Olmstead and Williams, Chemistry , Chapter 14 (all sections)
Purpose: The shift in equilibrium position of a chemical reaction with applied stress is determined.
Introduction
Chemical Equilibrium No chemical reaction goes to completion. When a reaction stops, some amount of reactants remain. For example, although we write
2 CO 2 (g) →← 2 CO (g) + O 2 (g) (1)
as though it goes entirely to products, at 2000K only 2% of the CO 2 decomposes. A chemical reaction reaches equilibrium when the concentrations of the reactants and products no longer change with time. The position of equilibrium describes the relative amounts of reactants and products that remain at the end of a chemical reaction. The position of equilibrium for reaction (1) is said to lie with the reactants, or to the left, because at equilibrium very little of the carbon dioxide has reacted. On the other hand, in the reaction
H 2 (g) + ½ O 2 (g) →← H 2 O (g) (2)
the equilibrium position lies very far to the right since only very small amounts of H 2 and O 2 remain after the reaction reaches equilibrium. Since chemists often wish to maximize the yield of a reaction, it is vital to determine how to control the position of the equilibrium. The equilibrium position of a reaction may shift if an external stress is applied. The stress may be in the form of a change in temperature, pressure, or the concentration of one of the reactants or products. For example, consider a flask with an equilibrium mixture of CO 2 , CO, and O 2 , as in reaction (1). If a small amount of CO is then injected into the flask, the concentration of CO 2 increases. Here the external stress is the increase in concentration of CO. The system responds by reacting some of the added CO with O 2 to yield an increased amount of CO 2. That is, the position of equilibrium shifts to the left, yielding more reactant and less CO. Reaction (1) also shifts with changes in pressure. Starting with reaction (1) at equilibrium, an increase in pressure causes the position of equilibrium to shift to the side of the reaction with the smaller number of moles of gas. That is, by shifting the equilibrium position to the left, the reaction decreases the number of moles of gas, thereby decreasing the pressure in the flask. In so doing, some of the applied stress is relieved. On the other hand, an increase in pressure for reaction (2) shifts the equilibrium position to the right to decrease the number of moles of gas. The response of a reaction at equilibrium to changes in conditions is summarized by LeChâtelier’s Principle :
A system perturbed from equilibrium shifts its equilibrium position to relieve the applied stress.
For an increase in temperature, the reaction shifts in the endothermic direction to relieve the stress. The decomposition of CO 2 , reaction (1), is endothermic in the forward direction. Upon an increase in temperature, the equilibrium position shifts in the forward direction to minimize the temperature increase. The formation of ammonia is exothermic:
N 2 (g) + 3 H 2 (g) →← 2 NH 3 (g) (3)
Upon an increase in temperature, the equilibrium position shifts to the left, which is the endothermic direction.
The Iron-Thiocyanate Equilibrium When potassium thiocyanate, KNCS, is mixed with iron(III) nitrate, Fe(NO 3 ) 3 , in solution, an equilibrium mixture of Fe3+, NCS–, and the complex ion FeNCS2+^ is formed:
Fe3+^ (aq) + NCS– (aq) →← FeNCS2+^ (aq) (4) yellow colorless red
The solution also contains K+^ and NO 3 –^ ions, but these are spectator ions and do not participate in the reaction. The relative amounts of the various ions participating in the reaction can be judged from the color of the solution. In neutral or slightly acidic solutions, Fe+3^ is light yellow, NCS–^ is colorless, and FeNCS2+^ is red. If the solution is initially reddish, and the equilibrium shifts to the right (more FeNCS2+), the solution becomes darker red, while if the equilibrium shifts to the left (less FeNCS2+), the solution becomes lighter red or perhaps straw-yellow.
Experimental Procedure For each of the external stresses described below, necessary information is provided regarding the manner in which one or more of the chemical species is affected. You will use a spot plate containing multiple wells and use a different well for each of the operations described, recording your observations of the color change of the solution. In a table, summarize your observations for each of the reactions that you perform on the iron- thiocyanate equilibrium. As an example, if you added a drop of concentrated HCl to the standard solution, the blood-red color lightens or perhaps disappears altogether. This change in color indicates that the FeNCS2+^ concentration decreases. To explain this result, it is necessary to know that in the presence of a large excess of Cl–, Fe3+^ forms complex ions:
Fe3+^ (aq) + 6 Cl–^ (aq) →← FeCl 6 3-(aq) (9)
The increase in Cl–^ reduces the Fe3+^ concentration, so in accord with Le Chatelier’s Principle, some FeNCS2+^ dissociates to replace some of the Fe3+^ removed by reaction with Cl–. This result is summarized in the table as follows:
Stress Observation Reactions of Interest Explanation +1 drop HCl
sol’n turned yellow (^) Fe3+^ + 6 Cl–^ → ← FeCl 6 3–^ Equilibrium shifted left in response to a decrease in [Fe3+] caused by reaction
with Cl–.
A. Operations to Introduce an External Stress - Record your observations in your data table in the format shown above.
- Add one drop each of 1 M Fe(NO 3 ) 3 and 1 M KNCS to 25 mL of distilled water. Mix well.
- Add a few drops of this solution to each of seven wells of a spot plate. One well serves as a color standard against which to judge color changes in the other wells. The other six wells are for performing your operations to introduce an external stress.
Results: The experimental stresses are listed below. Use your table from Part I to discuss the observed impact on the position of equilibrium and include a description of the underlying principle that explains the shift of the iron-thiocyanate reaction in terms of a response to the particular stress. For each case did the applied stress increase or decrease the amount of Fe3+^ or NCS–? Make sure to give the reactions that changed the Fe3+^ concentration for the addition of SnCl 2 , AgNO 3 , Na 2 HPO4, and NH 3. Finally determine if the complex formation reaction is endothermic or exothermic, as written.
Addition of extra Fe(NO 3 ) 3 Addition of extra KNCS Addition of SnCl 2 Addition of AgNO 3 Addition of Na 2 HPO 4 Addition of NH 3 Increase and decrease in temperature
Mention any student errors that may have caused problems in the determinations. All tables and plots should have a table or figure number and a caption. Refer to this Table or Figure with an explicit reference (e.g. see Figure 3) in the text of this section.
Discussion : (a). Purpose accomplished: Restate the purpose of the experiment, but as a completed goal. (b). Write the equilibrium reaction that you are studying. (c). Give a general summary statement about the agreement between the observed shifts of the position of equilibrium and the predictions based on LeChatelier’s principle. (d). Suggest one source of systematic error. Remember that student mistakes are neither random nor systematic errors; student mistakes are just student mistakes. For this experiment examples of systematic errors include errors in the concentrations of the reagents, the volumes of the reagents added, or the ability to judge the change in intensity of the red color of the solutions if the solution starts off too darkly colored. What effect does the source of systematic error have on the final results? (For example, does the systematic error cause an increase or decrease of the observed shift?) (e). To summarize the experiment answer the following question: how the outcome of chemical reactions at equilibrium be controlled?
Literature Cited : Give all literature cited, numbered according to the references in the body of your report.
Attach your data table if it is not included in the bulk of the Report.
Literature Cited:
- This experiment was modified from J. P. Birk, R. Bauer, and D. Sawyer, Laboratory Inquiry in Chemistry , Brooks/Cole, 2001.
