14-1: CHEMICAL EQUILIBRIUM & LE CHÂTELIER'S PRINCIPLE INTRODUCTION: Le Châtelier's Principle states that when a system at equilibrium is stressed the system adjusts so as to minimize the stress, if it can. In this experiment, several equilibrium systems will be studied. The stress applied to the system will be studied. The stress applied to the system will either be a temperature or concentration change. Each stress produces a change in color which allows you to determine the direction in which the equilibrium has shifted. An equilibrium condition is indicated by a double arrow (). For example, the systems studied are the following: Part A: Fe3+(aq) + SCN-(aq)  [FeSCN]2+(aq) Part B: Cr2O72-(aq) + 3H2O(l)  2CrO42-(aq) + 2H3O+(aq) Part C: [CoCl4]2-(al) + 6H2O  [Co(H2O)6]2+(al) + 4Cl-(al) * “(al)” indicates in solution in alcohol—ethanol unless otherwise indicated

When equilibrium is indicated by the double-headed arrow, ALL SPECIES SHOWN IN THE EQUATION ARE PRESENT AT EQUILIBRIUM. This is an important factor to keep in mind as you attempt to explain why a specific shift occurs when you stress the system. Before you begin the procedures you must determine the colors of the various ions. For example, when NaCl is placed in water it dissociates in the following manner: H2O

NaCl(s) (colorless)  Na+(aq) + Cl-(aq) (colorless) The reaction shifts almost totally to the right, that is, only Na+ and Cl- must be colorless. If you were to dissolve Na2CrO4 in water, the solution is yellow in color. Since the Na+(aq) ion is colorless, the CrO42-(aq) ion must be yellow. Na2CrO4(s)

(colorless)

H2O



2Na+(aq) + CrO42-(aq) (yellow)

MATERIALS: 0.1 M solutions of KCl, KSCN, BaCl2, FeCl3, K2Cr2O7, K2CrO4, NaOH, HNO3, AgNO3; ethanol, solid KCl, solid CoCl2•6H2O; graduated cylinder, distilled water, 6 test tubes, hot-water and ice-water baths.

PROCEDURE: Part A: Effect of concentration changes on the equilibrium system: Fe3+(aq) + SCN-(aq)  [FeSCN]2+(aq) 1. WEAR GOGGLES AND APRON FOR ALL PARTS OF THE LABORATORY! 2. Examine solutions of KCl, KSCN, FeCl3, and AgNO3. Write the colors for each ion in your data table. The color of the [FeSCN]2+ is determined in step 3. 3. Place 3 drops of 0.1 M FeCl3 and 3 drops of 0.1 M KSCN in a small test tube. Add distilled water and swirl until the orange-red solution is TRANSPARENT and UNIFORM in color intensity. (If you have filled the entire test tube with water, and it is still too dark, pour out half of the solution and add more water.) 4. Using 4 more test tubes, divide the solution into a total of 5 equal parts. 5. SAVE THE FIRST TEST TUBE TO BE USED AS A COLOR STANDARD. 6. To the second test tube, add enough drops of 0.1 M FeCl3 to produce a detectible color change (2-4 drops). Record the direction of the shift in you data table, and explain why the shift occurs. 7. To the third test tube, add several crystals of solid KCl. Record the direction of the shift in your data table, and explain why the shift occurs. 8. To the fourth test tube, add 2-4 drops of 0.1 M KSCN to produce a color change. Record the direction of the shift, and the reason for the shift. 9. Add 1 or 2 drops of AgNO3 to the last test tube. Record your observations. Part B: The effect of changing concentrations on the following: Cr2O72-(aq) + 3H2O(l)  2CrO42-(aq) + 2H3O+(aq) 10. Determine the colors of the Cr2O72- and CrO42- ions by examining the solutions of K2Cr2O7, K2CrO4, Na2Cr2O7, Na2CrO4. Record your findings. 11. Place about 2 ml of 0.1 M K2Cr2O7 in a test tube and add NaOH dropwise until a color change is detected. Record observations and the reason for the shift. 12. Put 2 ml of K2CrO4 in a test tube and add HNO3 dropwise until a color change is detected. Record observations and the reason for the shift. 13. To 2 ml of K2CrO4, add 1 ml of BaCl2. Allow the precipitate to settle (or centrifuge the mixture). Decant the solution and note the color of the precipitate. Explain the reason for the change. 14. To 2 ml of K2Cr2O7, add 1 ml of BaCl2. Allow the precipitate to settle (or centrifuge the mixture). Decant the solution and note the color of the precipitate in your data table. Explain the reason for the shift.

15. To 4 ml of K2Cr2O7, add 2 ml of HNO3 AND 1 ml of BaCl2. Record any color change or precipitate formation, or both. Explain any change. Part C:

Effect of temperature on the equilibrium: [CoCl4]2-(al) + 6H2O  [Co(H2O)6]2+(al) + 4Cl-(al) 16. Place a rice-grain-sized amount of CoCl2•6H2O in a test tube and add 2 ml of ethanol. Stir with a stirring rod until all of the solid is dissolved. If the solution is not pink, add water DROPWISE until the solution just turns pink. 17. Heat your solution using the hot-water bath provided until you observe a change. 18. Cool the test tube in the ice-water bath provided, and note the color change.

SUGGESTED DATA FORMAT: Ion K+ ClSCNFe3+ H3O+ NO3-

Color

Ion [FeSCN]2+ [FeCl4]Cr2O72CrO42[CoCl4]2[Co(H2O)6]2+

Color Colorless

Blue Red (Pink)

Fe3+(aq) + SCN-(aq)  [FeSCN]2+(aq) Species Added to

Color

Direction of

Equilibrium Mixture

Change

Shift

Fe3+ from FeCl3 Cl- from KCl SCN- from KSCN Ag+ from AgNO3

Reason for the shift

Equilibrium: Cr2O72-(aq) + 3H2O(l)  Original solution Cr2O72CrO4

2-

Species added to soln

2CrO42-(aq) + 2H3O+(aq)

Color

Direction of

Color

change

shift

of ppt

Explanation for change

OH- from NaOH H3O+ from HNO3

CrO42Cr2O72-

Ba2+ from BaCl2 Ba2+ from BaCl2

Acidified

Ba2+ from

Cr2O72-

BaCl2

Equilibrium: [CoCl4]2-(al) + 6H2O  [Co(H2O)6]2+(al) + 4Cl-(al) Procedure Rm-temp control

Color

Direction of shift (right or left) -----

Warm Solution Cool Solution

ANALYSIS: A. Discussion: When a stress is placed on a system at equilibrium the reaction shifts so as to establish new equilibrium conditions, if it can (Le Châtelier's Principle). In this experiment, the concentrations of reacting species were affected in one or more of the following ways: 1. Increasing the concentration of one of the reacting species by adding excess of that particular ion. 2. Decreasing the concentration of one of the reacting species by transforming it into a molecule. 3. Decreasing the concentration of one of the reacting species by forming a complex ion or by consuming the reacting species by forming a precipitate.

The Collision Theory tells us that reacting species must collide with enough energy (activation energy) and with the proper geometric orientation before a reaction can occur. If the concentration of a particular ion is increased, there are more collisions taking place and a higher probability that a successful collision will occur. For example, if you increase the Fe3+ and SCN- ions more [FeSCN]2+ is formed and the reaction shifts to the right (more orange). (Remember there are 2+ unsuccessful collisions between SCN and [FeSCN] , between Fe3+ and [FeSCN]2+, between SCN- and SCN-, and between Fe3+ and Fe3+.) If a precipitate like BaCrO4 forms, the concentration of the CrO42- ion is decreased. This results in more collisions between Cr2O72- and H2O. More CrO42is formed and the shift is to the right (lighter yellow). Possibly you noted that in Part B, step 15, you applied two stresses. You added acid (represented by H3O+ or H+) and Ba2+ (from BaCl2). You must determine if both stresses affect the direction in the same way or in opposite ways. Your observations should provide this answer. Other complex ions can form which also decrease concentrations of reacting species. The following reactions may help you to determine why a particular shift occurred: a) Fe3+(aq) + 4Cl-(aq)  [FeCl4]-(aq) b) Ba2+(aq) + CrO42-(aq)  BaCrO4(s) B. Terms to know: equilibrium Le Châtelier's Principle 

collision theory complex ion activation energy

C. Complete the data tables. Be sure to explain the reaction shifts in specific detail! (That is, what concentrations are increased or decreased, what complex ions are formed, or what precipitates result.)

D. Discussion questions: (Instructor will assign.) 1. Given the following hypothetical equilibrium reaction: A+B  C+D (yellow) (green) a) Describe a test that could be used to show that A is a participant in the equilibrium reaction. b) Describe a test that could be used to show that D is participant in the equilibrium reaction. 2. Explain how you know a system is at equilibrium. 3. Given the following information, determine the color of EACH ion. Give a written argument for your choice. i) XY2(s) ii) XR(s) iii) SY(s)

H2O

 H2O

 H2O



X2+(aq) + 2Y-(aq) (colorless) X2+(aq) + R2-(aq)

(green)

S+(aq) + Y-(aq)

(purple)

4. In Part C, you investigated the temperature effect on a system at equilibrium. Write a heat (HEAT) term on the proper side of the equation. Is the forward reaction exothermic or endothermic? CoCl42-(al) + 6H2O  Co(H2O)62+(al) + 4Cl-(al)