Equilibrium Reactions and Le Chatelier's Principle

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  • Topic: Oxalic acid, Color, Acid
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Lab: 5
Experiment:13

Pre-Lab
The purpose of this experiment is to observe an equilibrium reaction counteracting changes to it’s system all in accordance to Le Chatelier’s principle. An equilibrium reaction can be pushed toward products or reactant based on changes in temperature or concentration. The reversibility of reaction will also be looked at.

Pre-Lab questions
1. The concentrations of products and concentrations of reactants remain constant but both reactions are still going on, just at the same rate. 2. Le Chatelier’s principle when an equilibrium has been established in a system and is altered, the systems will counteract the change. 3. a) it will go to the left

b) to the left
c) to the right
4. Remove the CO2 gas being produced
5. concentrations and totals in equilibrium change over time. 6.
a) The solution will be yellow.
K2CrO4 ( K2+(aq) + CrO42-(aq)
b) The solution will be orange. The hydrogen ion will react with chromate to form an equilibrium and produce some dichromate and water. This reaction will cause the amount of chromate to decrease in order to form more products and the equilibrium will shift to the right. c) The NaOH, specifically the hydroxide, would react with the dichromate and push the equilibrium toward reactants. This would increase the amount of reactants and decrease the amount of dichromate being formed in the equilibrium. The solution would be yellow because there would be more chromate. 7. The common ion effect is the amount of dissociation a weak electrolyte undergoes with a common ion is reduced.

Review Questions
Complete questions # 1 - #4
1. The equilibrium would shift towards the right and the color of the solution would be orange. 2. The equilibrium would shift to the left. More acetate ion has been added to the equilibrium already containing this ion and making the solution more basic. Since the equilibrium has become more basic the solution will turn yellow from the methyl orange equilibrium. 3. The addition of sodium hydroxide would cause some calcium to form a precipitate with hydroxide. The amount of calcium and calcium oxalate in equilibrium 1 would decrease and oxalate anion would increase. Equilibrium 1 would shift to the left. Equilibrium two and three would also shift to the left because the excess oxalate anion would cause both reactions to increase their amount of products and they would have to counteract by shifting to the left. 4. The Hydrogen concentration effects the precipitation of the weak acid because the weak acid is at a complete equilibrium unlike the strong acid that dissociates completely. In the strong acid we can add as much hydrogen ions as we want but it cannot change the direction of reaction 2 since that reaction has gone to completion. Compare this to the equilibrium reactions of acetic acid from problem three with the addition or subtraction of hydrogen we can change the direction of the reaction.

Part 1: Chromate/Dichromate Ion Equilibrium
2 CrO42-(aq) + 2 H+ (aq)   Cr2O72-(aq) + H2O (l)  
.
1. Yellow to orange.
2. Orange to yellow.  
3. Yellow to orange. The reaction is the same as in step one. 4. Orange to yellow. The reaction was the same as step two.

5.
H+  (aq) +  OH-(aq)  H2O (l)
2 CrO42-(aq) + 2 H+ (aq)   Cr2O72-(aq) + H2O (l)   When sodium hydroxide was added to the dichromate solution, the orange color turned yellow. This was caused by hydroxide ions reacting with hydrogen ions forming water, shifting the equilibrium to the right. OH- ions remove H+ ions by neutralizing them and the system acts to stabilize the change again following Le Chatelier's principle and further shifting the color. The hydroxide ions stimulate the...
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