Lab Report on Le Chatelier's Principle

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LE CHATELIER’S PRINCIPLE
Castro, Lharize C.
Experiment # 1

I. Introduction:
In this experiment, using Le Chatelier’s principle, we will observe several responses of a system at equilibrium to various changes in external conditions.
The experiment aims to investigate two equilibrium systems: (a) cobalt complexes and (b) chromate-dichromate equilibrium and explain observations in light of the Le Chatelier’s principle.

II. Theory/Concepts:
In 1884 the French chemist and engineer Henry-Louis Le Chatelier proposed one of the central concepts of chemical equilibria. Le Chatelier's principle can be stated as follows: A change in one of the variables that describe a system at equilibrium produces a shift in the position of the equilibrium that counteracts the effect of this change. Le Chatelier's principle describes what happens to a system when something momentarily takes it away from equilibrium. When a system at equilibrium is disturbed, the equilibrium position will shift in the direction which tends to minimize, or counteract, the effect of the disturbance. * If the concentration of a solute reactant is increased, the equilibrium position shifts to use up the added reactants by producing more products.

* If the pressure on an equilibrium system is increased, then the equilibrium position shifts to reduce the pressure. * If the volume of a gaseous equilibrium system is reduced (equivalent to an increase in pressure) then the equilibrium position shifts to increase the volume (equivalent to a decrease in pressure) * If the temperature of an endothermic equilibrium system is increased, the equilibrium position shifts to use up the heat by producing more products. * If the temperature of an exothermic equilibrium system is increased, the equilibrium position shifts to use up the heat by producing more reactants.

III. Materials:
Reagents: Apparatus:
* 0.4 M CoCl2ő graduated cylinder
* 0.5 M HClő beaker
* 1 M HClő test tubes
* 6 M HClő test tube holder
* Ethanolő test tube rack
* 0.1 M K2CrO4ő stirring rods
* 1 M NaOHő thermometer
* 0.1 M Ba(NO3)2ő bunsen burner
* 0.1 M K2Cr2O7ő tripod w/ wire gauze

IV. Procedure:
Calibration of Medicine dropper
A 100 ml graduated cylinder was filled with water until the 1 ml mark. Using the dropper, we added water to the cylinder drop-wise, until it reached the 2 ml mark. We recorded the number of drops until the 2 ml mark and found out that 20 drops of water was added.

Therefore, 1 ml is equivalent to 20 drops on our medicine dropper.

Cobalt Complexes
We prepared solutions by dissolving 2 g of CoCl2 in a flask containing 100 ml of water. In another flask, we dissolved 2 g of CoCl2 in a flask containing 100 ml of ethanol. We took down notes on the initial color of the solutions in each flask.

We’ve poured 2 ml of ethanol-solution into a test tube and added just enough water to cause a change in its color. After we’ve noted the number of drops, we then added a concentrated HCl solution until the color had changed again. We note down again the number of drops. We’ve repeated this same procedure in a water-solution, 0.5 M HCl, 1 M HCl, 6 M HCl and 0.5 M NaCl solution.

Chromate-Dichromate Equilibrium
In different test tubes, we’ve prepared approximately 1 ml (10 drops) of 0.1 M K2CrO4 (aq) solution, (10 drops) of 0.1 M K2Cr2O7 (aq) solution. We’ve recorded the initial color of each solution. Then, we added 1 M HCl drop by drop to each test tube until at least one of the tube changes color (maximum of 1 ml).

We repeated the first step using fresh solutions and we added 1 M NaOH drop by drop to each test tube until at least one of the tube changes color (maximum of 1 ml).
In another set of test tubes, we have prepared approximately 1 ml 0.1 M K2CrO4 (aq) solution and added 2 drops of 1 M NaOH and 0.1 M Ba(NO3)2 ions drop by drop until a change in color was...
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