Experiment 3: Le Châtelier’s Principle

Cabajar, Jairus B.,

Dionisio, Nicole Anna Marie H.Group 3

INTRODUCTION

“If an external stress is applied to a system at equilibrium, the system adjusts in such a way that the stress is partially offset as it tries to reestablish equilibrium”

—Henri Le Châtelier.

At equilibrium, both products and reactants are present in the reacting system.

INTRODUCTION stressors

Changes in Concentrations

An increase in the concentration of reactants, or decrease in the concentration of products, causes a right shift. On the other hand, an increase in the concentration of products, or decrease in concentration of reactants, causes a left shift.

Basically, increased concentrations are used up while decreased concentrations are filled up.

INTRODUCTION stressors

Changes in Pressure and VolumeChanges in pressure usually only affect reacting

species in the gas phase. By the ideal gas law, pressure is inversely proportional to volume.

Generally, an increase in pressure (decrease in volume) causes the reaction to shift to the side with fewer moles of gas; on the other hand, a decrease in pressure (increase in volume) causes it to shift to the side with more moles of gas. Whenever there is equal number of moles on both sides, a change in pressure (or volume) has no effect.

INTRODUCTION stressors

Changes in Temperature Only a change in temperature can alter the equilibrium constant.

In general, a temperature increase favors an endothermic reaction, and a temperature decrease favors an exothermic reaction.

∆H0 ∆T Shift

+ + + (forward)

+ - - (reverse)

- + -

- - +

INTRODUCTION stressors

The Effect of CatalystA catalyst affects the activation energy of the

forward and reverse reaction to the same extent. Adding a catalyst to a reaction will simply cause the mixture to reach equilibrium sooner (in the case of a + catalyst).

The presence of a catalyst does not alter the equilibrium constant, nor does it shift the position of an equilibrium system.

EXPERIMENTATION and RESULTS

H2O

0.1 M Cu(NO3)2

0.1 M NH4OH

0.1 M K4Fe(CN)60.1 M HCl

0.1 M NaOH

0.1 M NaNO3

1ml of solution + 0.50ml water subjected to water

bath

DISCUSSION

For the first solution in part A, this equilibrium was formed:

Cu+2(aq) + 4NH3(aq) Cu(NH3)4+2(aq)

sky blue colorless dark blue

DISCUSSION

For the systems A and B, similar concepts may be applied. By increasing the concentration of reactants, the reaction shifts to the right forming more products, causing the solution to become darker. In the A depression, when Cu+2

dissociated from the added reagent, Cu(NO3)2, the number of effective collisions between the reactants increased. The same is true in B depression where NH3, from the reagent, was added to the solution. Both of these favor the formation of products.

DISCUSSION

In the system C, a brownish-red precipitate was formed after adding 0.1 M K4Fe(CN)6. In this system, as the reagent dissociated to K+1 and Fe(CN)6

-4, the ferrocyanide ion reacted with the Cu+2. By consuming the copper ion, the reaction counteracts by shifting to the reactant’s side. The formation of Cu2Fe(CN)6 caused the solution’s color.

DISCUSSION

For system F, based on the experiment, the solution became darker after adding 0.1 M Na(NO)3. Theoretically, the said reagent must not affect the reactant. After dissociating to Na+1 and NO3

-1, none of these ions may significantly disturb the equilibrium. A contamination may account to the obtained results.

DISCUSSION

For system D, HCl, an acid was added to the original solution. H+ from the dissociation will react with NH3, forming NH4

+, decreasing the reactant. In order to conform to this, the reaction shifts to the reactant side, resulting to a lighter color. For system E, a base, NaOH, was added. When OH- dissociated from the reagent, it reacted with Cu+2, like that of system D, the reaction shifted to the left.

DISCUSSION

In the case of system G, the results suggest that it is endothermic because it became lighter. But since the reaction is known to be exothermic, decreasing the temperature must cause it to shift forward, thus turning the solution darker. Again, a contamination may account to this.

DISCUSSION

Summary of how the reaction will theoretically adjust to the different applied stressors.

System Applied stress Reaction shift

A 0.1 M Cu(NO3)2 forward

B 0.1 M NH40H forward

C 0.1 M K4Fe(CN)6 Reverse/backward

D 0.1 M HCl Reverse/backward

E 0.1 M NaOH Reverse/backward

F 0.1 M Na(NO)3 No shift

G Decrease in temp. Forward

DISCUSSION

For the part B of the experiment, the K2CrO4 turned from yellow to orange when HCl was added to it. And when NaOH was added, the solution became lighter, almost the same as that of the original. For this part of the experiment, this equilibrium was formed:

2CrO4-2 + 2H+ Cr2O7

-2 + H2O yellow orange

DISCUSSION

When HCl was added, the solution became acidic. As HCl dissociated to H+ and Cl-, the hydrogen ion reacted with CrO4

-2 and Cr2O7-2

(an orange solution) was formed. It can be concluded that Cr2O7

-2 is predominant in an acid solution due to the H+ from the dissociation of an acid. Next, a base, NaOH, was added. When OH- dissociated from the base, it reacted with H+, consuming it and at the same time forming more water which is part of product’s side. By this, the reaction shifts to the left, in a basic medium, Cr04-2 is dominant.

DISCUSSION

For the last part of the experiment, when Pb(NO3)2 was added to K2CRO4, a yellow precipitate was formed. This equilibrium was formed:

CrO4-2 + Pb+2 PbCrO4 yellow yellow precipitate

And when HCl was added to the solution, it became darker—from yellow to orange. The precipitate also lessened.

DISCUSSION

When the reaction was subjected to HCl, similarly to system H, H+ reacted with Cr04-2 forming Cr2O7

-2 causing it to change color. This can be verified by the decrease of the precipitate, the CrO4

-2 was consumed causing the reaction to shift to the reactant’s side.

CONCLUSION

Several factors affect the shift of an equilibrium. Changes in concentration, pressure and volume can alter the equilibrium concentrations of the reacting mixture, but not the equilibrium constant as long as the temperature remains constant. Only a change in temperature changes the value of the equilibrium constant.

RECOMMENDATION

We recommend changing the system being studied in this experiment. Observing the changes in color—which is basically the core procedure of the experiment—caused confusion. The light blue color of the solution made it hard to see if the solution became lighter or not. Systems that are expected to be darker, on the other hand, were equally confusing due to the minimal darkening of the solution.

Proper handling of the reagents, too, must be observed, several deviations from the theoretical results is probably due to contaminations.

RECOMMENDATION

Finally, to fully understand the principle, the effect of pressure (and volume) as a stressor must also be conducted. But since these kinds if experiments are in need of tremendous amount of effort and experience, at least allow the students to see how the concept works for the sake of better understanding.

SOURCES

Chang, Raymond. (2008). General Chemistry: The Essential Concepts, Fifth edition. New York: The McGraw-Hill Companies, 2008.

Committee on General Chemistry (2007) Learning Modules in General Chemistry 2. University of the Philippines Manila.

"Henri Louis Le Châtelier." Microsoft® Encarta® 2007 [DVD]. Redmond, WA: Microsoft Corporation, 2006.

Silberberg, M.S. (2006). Chemistry: The molecular nature of matter and change (4th edition). New York: The McGraw-Hill Companies, 2008