LeChatlier's principle states that if a system is in a state of dynamic equilibrium and that system is consequently perturbed by an external stress, the system will adjust its equilibrium to compensate for that stress.
In the lecture demonstration we observed the effects of the temperature, along with adding and removing various species on the equilibrium distribution of the cobalt complexes (indicated by the reaction in the top left frame of this web page.)
With this simulation, you can continue our investigation by not only observing the colors of the solutions, but by actually observing the concentration changes of the various species.
While performing any virtual activities, you should adjust the upper left frame to show the reaction being studied, and interpret all your results in terms of this chemical equation.
| [CoCl4]-2
In the presence of chloride ions, a blue tetrahedral complex forms, the [CoCl4]-2 complex ion. These complexes are very colorful and we will be studying them in chapter 20.
In the simulation we are using 1M CoCl2, while in the photo we have 1 M Co(NO3)2. The solutions are both pink. Can you explain this observation? If not, use the virtual lab to transfer some 1M Cobalt(II) chloride exp. sol from the stock room to the lab bench. Which species is present in the greatest concentration?
The color of the solution depends on the relative concentration of these two respective complex ions as they exist in a state of dynamic equilibrium.
This color difference allowed us to demonstrate LeChatlier's principle in lecture. As the system shifted its equilibrium to the right (according to the above equation), it would turn blue, and as the system shifted its equilibrium to the left, it would turn pink. |
| What happens when you cool the system? |
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| What happens when you heat the system? |
| Assignment: |
| This exercise is a nongraded simulation of an in class lecture demonstration and should assist you in preparing for the second hour exam. Use the equilibrium concentrations after each step to determine K for the above equation. If you wish to print this page, left click over the frame and select "open frame in new window". Now you can print it. |
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1. Add 25 mL of [Co(H2O)6]+2 to an empty Erlenmeyer flask. Now add 12 M HCl in 1mL increments until the equilibrium color has changed. Hint, type in 1mL for the volume to be transferred, and then keep clicking "pour" until you see a change - counting clicks to determine total volume added. |
2. Predict the effect of removing chloride ions. Now remove some of the free chloride ions by adding some silver nitrate. Hint, add 1 mL amounts of the silver nitrate successively until the equilibrium color has changed - instead of a whole bunch at once.
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AgNO3(aq)--> Ag+ + NO3-
Ag+ + Cl- --> AgCl(s)
AgNO3(aq) + Cl- --> AgCl(s) + NO3-
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Note by the above equation how Ag+ scavenges free chloride ions by tying them up in a precipitate, and thus removes them from the solution. You can determine the mass AgCl formed by clicking "Solid" on the "Solution Info" window of the virtual lab.
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3.Predict the effect of adding HCl. Now add HCl in very small incremental steps until the equilibrium color has changed. Discuss how adding HCl effects both the Cl- and the H2O concentrations. |
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4. Is the reaction as written endothermic or exothermic? Right click on the flask and choose "thermal properties". You can now change the temperature between 0 and 99 deg C. Heat or cool the system until you have observably perturbed the equilibrium. Then apply LeChatlier's principal to determine if it is exothermic or endothermic |
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5. Allow the system to reach thermal equilibrium (constant temperature). Use the concentration values to determine K. Now go to the thermal properties, change the temperature and click on the thermally isolated system option. Determine the new K at the new temperature. From the new K at the new temperature, determine if the system is endothermic or exothermic.
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