Data Table for Experiment 12G: Principles of Equilibrium and Thermodynamics

Directions for Expt. 12H Parts A and B (Principles of Equilibrium and Thermodynamics)

Pre-lab: You do NOT have to do the pre-lab.

Lab report: Everything you need for the lab report (data tables, questions, and calculations) are included in this document. This is for your convenience so you don’t have to flip back and forth to answer the questions.

To repeat: Once you have completed this document, you are done with the lab report. All of the calculations and questions that I want you to do have been transferred from the lab to this document.

The final report should be very neatly written with all work clearly shown. When doing a calculation, show only one step on each line and include units on every number that has one. You may need to re-write the report to achieve a professional result. A very important (perhaps most important) part of science is the clear communication of results!

DO NOT INCLUDE THIS PAGE IN YOUR REPORT!
Names: ______

Experiment 12H, Parts A and B: Principles of Equilibrium

Net ionic equation for the reaction performed in this experiment:

What is the role of HNO3 in this reaction?

Color of KSCN solution: ______

Color of Fe(NO3)3 solution: ______

Color of FeSCN2+ solution: ______

Wavelength chosen for this experiment ______

Procedure and Data

Part A. Determining the molar absorptivity of FeSCN2+

Mix the following in an empty, clean cuvette:

2.00 mL of 2.00 x 10-4 M KSCN

2.00 mL of 2.0 M HNO3

Use the solution above as a blank to set 100% T point. Set 0% T point with no cuvette in the spec.

Weight out about 0.8 g of solid Fe(NO3)3∙9H2O. Add about 1/8 of the sample to the solution above. Mix to dissolve and read absorbance (A). (Make sure solid is dissolved before you read absorbance.) Repeat until the absorbance no longer changes. (You might have to get more solid!)

Total mass of solid Fe(NO3)3∙9H2O added ______(include unit)

Part B. More Examples of LeChatelier’s Principle

Now we will do the reaction 3 times. (Each pair will do only one of the reactions.) Each time, we will start with different initial concentrations of reactants and determine the equilibrium constant.

Reaction 1: Mix the following in an empty, clean cuvette and read the absorbance:

2.00 mL of 2.00 x 10-2 M Fe(NO3)3 (dissolved in 1 M HNO3)

2.00 mL of 2.00 x 10-4 M KSCN

Class average for the equilibrium value of A = ______

Reaction 2: Mix the following in an empty, clean cuvette and read the absorbance:

2.00 mL of 2.00 x 10-2 M Fe(NO3)3 (dissolved in 1 M HNO3)

2.00 mL of 2.00 x 10-4 M KSCN

1.00 mL of 2.0 M HNO3

3.00 mL of water

(Notice that you added the same number of moles of each reactant as in reaction one but now the total volume is 8.00 mL rather than 4.00 mL.)

Class average for the equilibrium value of A = ______

Reaction 3: Mix the following in an empty, clean cuvette and read the absorbance:

2.00 mL of 2.00 x 10-2 M Fe(NO3)3 (dissolved in 1 M HNO3)

1.00 mL of 2.00 x 10-4 M KSCN

1.00 mL of 2 M HNO3

4.00 mL of water

Class average for the equilibrium value of A = ______

Add a single crystal of KSCN to your last solution and mix to dissolve. Note any changes to the absorbance. Leave the sample in the Spec-20 for a few minutes and note what happens to the absorbance.

Data Analysis

Part A. Determining the molar absorptivity of FeSCN2+

Molar mass of Fe(NO3)3∙9H2O ______(include unit)

(Note: When asked to calculate an initial concentration, pretend that no reaction has taken place.)

1. Calculate the initial concentrations of Fe3+ and SCN- after adding all of the solid Fe(NO3)3. (Ignore the volume change that occurs upon adding the solid.)

[Fe3+]o = ______

[SCN-]o = ______

Show work below:

[Fe3+]o [SCN-]o

2. After adding all of the solid Fe(NO3)3∙9H2O and allowing the reaction to occur, what is the concentration of FeSCN2+? Justify your answer. (Hint: Compare the two initial concentrations calculated above. You have forced the reaction to go to completion!)

3. Use Beer’s law (A = cl) to calculate the molar absorptivity (a.k.a. extinction coefficient) for FeSCN2+. Include units. l = 1.00 cm. (A is in your data table and you just calculated the value of c!)

Part B. More Examples of LeChatelier’s Principle

To calculate an equilibrium constant, you need to know the initial concentrations of all reactants and products but you need to know only one of the equilibrium concentrations. (The initial concentrations of products is usually zero!)

For all 3 reactions, calculate the initial concentration of both reactants ( [Fe3+]o and [SCN-]o ) and the equilibrium concentration of the product ( [FeSCN2+]eq ). To calculate [FeSCN2+]eq, use Beer’s law and the value of calculated above.

Show your work for reaction 1 below. Enter all of your answers in the table above.

For all 3 reactions, complete an ICE table and calculate Keq for the reaction at room temperature.

Questions

1. You’ve calculated the equilibrium constant for this reaction 3 times by doing the reaction with 3 different initial concentrations of reactants. How do your three calculated equilibrium constants compare? Is this what you should have expected? Explain. (Assume that two values that differ by 10% or less are actually the same. This takes into account experimental error.)

1. Notice that reaction 2 started with the same number of moles of each reactant as reaction 1. However, the total volume was 8.00 mL rather than 4.00 mL. Thus, the initial concentrations in reaction 2 were half of what they were for reaction 1. Imagine doing reaction 2 in a different way. Start by doing reaction 1 and allowing it to come to equilibrium. Then add 4.00 mL of water. Before you dilute the reaction with water, the reaction is at equilibrium so Q = Keq.

1. At the moment you add the 4.00 mL of water and before any additional reaction can occur, what is the value of Q? Show your work. (Hint: The concentrations of reactant and products are now ½ of their equilibrium values. See your ICE table for reaction 1.)

1. You should find that Q is no longer equal to Keq so the reaction is no longer at equilibrium. (You have “disturbed” the equilibrium.) Compare Q to Keq and determine which way the reaction should proceed (forward or backward) in order to re-establish equilibrium.

1. How would LeChatelier determine which way the reaction should proceed upon dilution? (Hint: Think about what happens to the TOTAL concentration of solutes when you dilute. Which way must the reaction proceed to counteract this change?)

1. At the moment you dilute reaction 1, the absorbance drops to half of its original value because the concentration of the product (the red stuff) is halved. However, since the reaction is no longer at equilibrium, the absorbance starts to change. As the reaction heads toward equilibrium, will the absorbance increase or decrease? Explain.

1. Compare the absorbance measured for reaction 1 to the absorbance measured for reaction 2. Is this result consistent with your answer to part d? Explain.

1. When you added a crystal of KSCN to your reaction, you disturbed the equilibrium. Explain what happened to the absorbance (immediately) in two ways. First compare Q to Keq at the moment the crystal was added and discuss the consequences. Second, use Le Chatelier’s principle.

1. After adding the crystal of KSCN, you left the sample in the Spec-20 for a few minutes. Explain what happened to the absorbance by comparing Q to Keq and by using Le Chatelier’s principle.. (Read “One last complication” on page one of the lab.)

1. The reaction in this experiment is exothermic. Assume the reaction is at equilibrium. Predict what should happen to the absorbance if you heat the reaction after it reaches equilibrium. Explain in terms of the effect of temperature on the equilibrium constant.

1