Lab 4: Equilibrium and Le Châtelier’s Principlenobel.scas.bcit.ca/courses/wpmu/chem3615/files/2012/08/... · Web viewThis is usually due to the equilibrium state that is reached

Lab 5: Equilibrium and Le Châtelier’s Principle

Objectives: To explore the effect of temperature and changing the concentrations of reactants and products on the equilibrium composition of five equilibrium systems.

Introduction: Most chemical reactions do not result in a 100% yield of products based on the stoichiometry of the reaction. This is usually due to the equilibrium state that is reached when the forward rate of reaction equals the rate of the reverse reaction. In this lab, the effect of qualitative changes on a number of reactions at equilibrium will be studied.

Le Châtelier’s Principle states that “If a change in conditions is imposed on a system at equilibrium, the equilibrium position will shift in a direction that tends to reduce that change in conditions.”

For example, the change in conditions could be either the temperature or concentration of a reactant or product and the effects observed. It should be noted that for a system, there exists many equilibrium positions but only one value of the equilibrium constant at a specific temperature. Adding more of a reactant or product will increase the concentration of that species (addition of a common species). By adding a substance that will react with the reactant or product (eg via a side reaction) the concentration of the reactant or product will decrease.

In this experiment, we will study the equilibrium of five systems and observe the reaction of the equilibrium systems as predicted by Le Châtelier’s Principle. The five systems are:

Part A - The Equilibrium of Co(II) Complex Ions

The element cobalt can form compounds in two different oxidation states, +2 and +3. The +2 state is more common. The chloro complex of cobalt (II), CoCl42-, is tetrahedral while the aquo complex of cobalt (II), Co(H2O)6

2+, is octahedral. Both of these complexes exhibit different colours. Cobalt complexes are used as drying agents with the colour change indicating when the drying agent should be changed. The equilibrium reaction is:

Co(H2O)62+(aq) + 4 Cl-(aq) CoCl42-(aq) + 6 H2O() ΔH = +50 kJ/mol (1)

Part B - The Equilibrium of the thiocyano-iron(III) complex ion

When colourless aqueous solutions of iron (III) ion, Fe3+, and thiocyanate ion, SCN-, are combined, the reaction that occurs produces the thiocyanoiron (III) complex ion, FeSCN2+, which is responsible for the equilibrium mixture's deep red colour.

Fe3+(aq) + SCN-(aq) FeSCN2+(aq) (2)

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The colour of the thiocyanoiron (III) complex ion, FeSCN2+, solution will indicate how the equilibrium system is being affected.

Part C - The Equilibrium of a Mg2+ Precipitate

Reactions which form precipitates are written as an equilibrium using the solubility product. If there is a precipitate MX, then the Ksp expression is:

MX(s) Mn+(aq) + Xn-(aq) (3)

Part D: The Equilibrium of an Acid-Base Indicator

An acid-base indicator can be used to observe an equilibrium reaction. Indicators are weak acids which show one colour in the acid form, HInd, and another colour in the basic form, Ind-. At pH = pKa of the indicator there are equimolar amounts of the conjugate forms and the observed colour is a mixture of the two. Bromothymol blue is a yellow-green-blue indicator which has a pKa of 7.0. The reaction of the indicator bromothymol blue can be illustrated as follows:

HInd(aq) H+(aq ) + Ind-(aq) (4)

Part E: The Equilibrium of a Saturated Solution

A saturated solution is solution in which the dissolved substance is in equilibrium with some of the undissolved substance. One example is a salt such as in the equilibrium (3). In this case when one of the ions is added to the solution, the equilibrium will shift to the left and some of the solid will precipitate out.

Procedure: During this lab, record all observations such as colours, any liberation or absorption of heat, and any precipitates formed or dissolved.


1. Number four clean and dry test tubes. Record the initial colour of the stock CoCl26H2O solution.

2. Pour 12 mL of 0.1 M CoCl26H2O into a clean, dry 50 mL beaker. In the fume hood, add concentrated HCl (several squirts) with mixing until a colour change is observed. Record the observed colour.

3. Divide the solution equally into four test tubes. Set aside test-tube #4 to use as a reference.

4. In Test tube #1 add water with mixing until a colour change is produced. Record the observed colour.

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5. Heat Test tube #1 in a hot water bath (add boiling chips) and you should see a colour change. (If you don’t then you have likely added too much water. Try again with another sample.) Record the observed colour.

6. Cool Test tube #2 in an ice water bath and record the observations. Record the observed colour. Keep your ice bath for Part B.

7. Heat Test tube #3 in a hot water bath and record your observations. Record the observed colour. Keep your hot water bath for Part B.

8. Dispose of the cobalt solutions in the waste bottle.


You should compare the colour of each tube with the reference Test tube #1. Note the colour of the 0.1 M Fe(NO3) 3.

1. In a 100 mL beaker, combine: 1.5 mL of 0.1 M Fe(NO3)3, 1.5 mL of 0.1 M KSCN, 50 mL H2O.

2. Pour 5 mL of the solution into nine numbered test tubes.

3. Add two drops of H2O to Test tube #1, which will serve as reference for colour. Record your observations.

4. Add two drops of 1 M Fe(NO3)3 to Test tube #2. Record your observations.

5. Add two drops of 1 M KSCN to Test tube #3. Record your observations.

6. Add 8 drops 6 M NaOH to Test tube #4. The precipitate Fe(OH)3 will take a few minutes to form. Record your observations.

7. Add 4 drops of AgNO3 to Test tube #5. The precipitate is AgSCN. Record your observations.

8. Add 4 drops of 0.1 M NaCl to Test tube #6. Record your observations.

9. Place Test tube #7 in an ice water bath. Record observations.

10. Place Test tube #8 in a boiling water bath. Record observations.

11. Add 1 mL of distilled water to Test tube #9. Record your observations.

12. Now add an additional 4 mL of water and record your observations.

13. Dispose of the solutions into the waste bottle.

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Part C - The Equilibrium of a Mg2+ precipitate

1. Into a test tube add: 1 mL water, 10 drops of 0.1 M Mg(NO3)2, 10 drops of 6 M NH3 (note it maybe labelled as NH4OH)

Record your observations.

2. Add a small amount (1/4 spatula) of solid NH4Cl to the test tube and mix to dissolve. Record your observations. What is the product? Consult a solubility table.


1. Obtain a pH 7 buffer solution and pour 3 mL into a 50 mL beaker. Add 5 drops of bromothymol blue indicator. Record your observations.

2. Add 1 M HCl drop wise with mixing until the solution is acidic and the indicator shows a colour change. Record your observations.

3. Add 1 M NaOH drop wise to return to the original colour and continue until the solution is basic and a new colour is reached. Record your observations.

Part E: The Equilibrium of a Saturated Solution

1. To 3 mL of saturated sodium chloride (5.4 M) add 1-2 mL of concentrated HCl in the fumehood.. Record your observations.

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Lab 5: Equilibrium and Le Châtelier’s Principle Date: _______________

Name: __________________________ Partner: _____________

Purpose:

Data and Results

Use the following key to answer some of the questions following the observations for each step.

Equilibrium shifts left (towards reactants) Increase in amount

No noticeable shift in equilibrium No change in amount

Equilibrium shifts right (towards products) Decrease in amount


CoCl26H2O solution – Colour ________________________

Reagents Observations

CoCl26H2O solution + HCl

Equilibrium Reaction:


Test Tube #1 + H2O

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The colour of the solution is due to the presence of which ion? ____________

Direction of equilibrium shift (circle answer)

This shift (if any) is due to (circle response) a) common species (name it) ___________________

b) side reaction (show below)


Test Tube #1 + H2O + hot water bath



Which statement is true? (circle response) a) heating favours an exothermic process

b) heating favours an endothermic process


Test Tube #2 + ice water bath



Which statement is true? (circle response) a) cooling favours an exothermic process

b) cooling favours an endothermic process


Test Tube #3 + hot water bath

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Heating favours the equilibrium __________? (circle response) a) reactants

b) products

Write the equilibrium reaction with heat as a product or reactant



Test Tube #1 - Fe(NO3)3 + KSCN + H2O

Equilibrium Reaction:



Test Tube #2 + Fe(NO3)3




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Test Tube #3 + KSCN





Test Tube #4 + NaOH





Test Tube #5 + AgNO3

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Test Tube #6 + NaCl


Provide an explanation of your observations


Test Tube #7 + ice water bath




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Test Tube #8 + hot water bath



Write the equilibrium reaction with heat as a product or reactant.


Test Tube #9 + 1 mL water

Test Tube #9 + 5 mL water

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Write the equilibrium constant expression.


Using the equilibrium constant expression, explain the observations when water is added.

Part C - The Equilibrium of a Mg2+ precipitate


Mg(NO3)2 + NH3

Equilibrium reaction (written as a Ksp):


Tube + NH4Cl

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Buffer + bromothymol blue (Hind)

Equilibrium reaction (written as a Ka):


Solution + HCl

On addition of this reagent the [H+] has (circle the answer)





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Solution + NaOH

On addition of this reagent the [H+] has (circle the answer)




Part : The Equilibrium of a Saturated Solution


Solution + HCl

Equilibrum Equation (written as Ksp)




What is the identity of any precipitate formed? ____________________

Conclusions:

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Lab 4: Equilibrium and Le Châtelier’s Principlenobel.scas.bcit.ca/courses/wpmu/chem3615/files/2012/08/... · Web viewThis is usually due to the equilibrium state that is reached

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