Experiment #7. Determination of an Equilibrium Constant

Introduction

It is frequently assumed that reactions go to completion, that all of the reactants are converted into products. Most chemical reactions do not go to completion because they are equilibrium systems where the reaction proceeds in both directions. As the reactants are used up, the rate of the forward reaction decreases. Conversely, as the concentrations of the products increase, the rate of the reverse reaction increases. Eventually, the rate of the forward reaction equals the rate of the reverse reaction and the concentrations of the reactants and the products stay constant. The system has reached a state of dynamic equilibrium. At equilibrium, both the forward and reverse reactions are occuring, but no net change is observed. Consider the general reaction:

aA bB ⇌ cC dD where a,b,c and d are the stoiochiometric coefficicents.

( 1 )

Experimental evidence shows that the ratio of products to reactants (with each product and reactant expressed as a molar concentration and raised to its stiochiometric coefficient) is a constant for a reaction that has reached equilibrium. This constant, which is different for each chemical reaction, is known as the equilibrium constant and is designated with the letter K. There is a separate value of K for each temperature at which the reaction occurs. Thus, at equilibrium, the equilibrium constant K is equal to:

K =

C D

A B

c d

a b

( 2 )

where the brackets [ ] imply molarity and the exponents are the stoichiometric coeffients of the balanced chemical equation. The equilibrium constant measures the extent to which a chemical reaction occurs. The larger the value for K, the greater the tendency for the reaction to go to completion is and the more products will be formed relative to the reactants. In this experiment you will determine the equilibrium constant for the following reaction: (Spectator ions are not shown.)

Fe 3+ (aq) + SCN- (aq) ⇌ FeSCN 2+ (aq) ( 3 )

K

SCNFe

FeSCN3

2

( 4 )

Solutions of Fe3+ and SCN- will be mixed and will react to form some FeSCN 2+. The initial amounts of Fe3+ and

SCN- can be calculated. The equilibrium concentration of FeSCN 2+ will be found using its spectroscopic properties – how much light it absorbs at a specific wavelength. FeSCN 2+ is a blood red complex that absorbs the blue-green wavelengths of visible light. Its absorbance is directly proportional to its concentration. The absorbance (a measure of the amount of light absorbed) will be measured by a spectrophotometer. Solutions to be measured are placed in cuvettes; these are square tubes have minimal absorbance in the wavelength range of the spectrophotometer.

A cuvette for measuring light absorption. It is usually made of quartz or plastic. On some, two

of the sides are frosted. For a solution placed in a 1 cm cuvette, the absorbance, A, is equal to the “extinction

coefficient”, ε (epsilon), times the molar concentration, C. The value of ε is can be determined experimentally for each substance from solutions of known concentration at a particualr wavelength. ε varies with the wavelength of light. In this experiment we will be measuring absorbtion at 470 nm.

A = ε C ( 5 )

We will use a best-fit line equation to find the concentration from the absorbance measurement for each sample. Using an equilibrium (ICE) chart, the equilibrium concentrations of Fe 3+ and HSCN are then calculated. Finally, the equilibrium concentrations are put into equation ( 4 ) to find the equilibrium constant, K. Note: All of the solutions are made in 1.0M HNO3(aq), so be cautious and wear gloves.

Equipment

4 small beakers 5 cuvettes stirring rod thermometer Venier equipment and colorimeter

Chemicals

0.0020 M KSCN(aq), potassium thiocyanate (source of SCN-) 0.200 M Fe(NO3)3(aq), iron (III) nitrate dissolved in 1.0M HNO3(aq). (Source of Fe3+) 0.0020 M Fe(NO3)3(aq), iron (III) nitrate dissolved in 1.0M HNO3(aq). (Source of Fe3+)

Spill/Disposal The contents of all test tubes, cuvettes, and beakers may be disposed of in the sink. Flush with a large

volume of water.

Prelab Calculations Note: you MUST have these calculations done before coming to lab. For all 4 of the solutions you will prepare in Step 3 of the procedure, calculate [FeSCN2+]. For these solutions, assume that moles of SCN- = moles of FeSCN2+ (a large excess of Fe3+ makes the SCN- react completely). Use M1V1 = M2V2. V1 should be the volume of SCN- in that beaker, and V2 should be the total volume of everything in that beaker.

Beaker [FeSCN2+]

1

2

3

4

*You will need to enter these concentrations into Excel to make graphs for the postlab, so have these ready when you come to lab. Also, copy them into the first data table in your postlab.*

Procedure

1. Obtain 5 cuvettes. Fill one (to the mark) with distilled water. This will serve as a “blank” for calibration.

Notes for using cuvettes:

Handle cuvettes toward the top or bottom, so that the middle (where measurements are taken) stays free of fingerprints and scratches.

If there are bubbles in a filled cuvette, gently tap or knock the cuvette until the bubbles rise to the top (so they don’t interfere with the measurement).

Position the cuvette so that the measuring light passes through the clear sides (many cuvettes have frosted sides, which are bad for measuring). If you look inside the well of the colorimeter you will see two sides with holes where light passes through.

Before measuring, gently wipe the outside of each cuvette with a Kimwipe.

Use a lid to prevent spilling inside the colorimeter. Lids can be rinsed and reused.

Remember to remove the last cuvette before returning the spectrometer!

2. Obtain a Vernier Colorimeter, and connect it to a Labquest machine. Calibrate the Colorimeter with the blank:

a. Open the Colorimeter lid, insert the blank, and close the lid. b. Press the <or> button on the Colorimeter to select the 470 nm wavelength. c. Press and hold the CAL button until the red light begins to flash. Release the button. When the light stops

flashing, calibration is complete.

3. (Making standard solutions for a calibration line) Obtain 4 small beakers. Mix solutions in each beaker as shown in the table below, measuring all volumes with large, non-disposable pipets. Mix each solution thoroughly. Note: the products of this reaction begin forming immediately, but they degrade in light, so be sure to measure the absorbance of each solution within 2-5 minutes of mixing the chemicals. Beakers 1-4 are made from 0.200 M Fe(NO3)3. Beakers A,B, C are made from 0.0020M Fe(NO3)3. Do not mix these up!

Beaker 0.200 M Fe(NO3)3 (mL)

0.0020 M SCN- (mL)

Distilled H2O (mL)

1 5.0 4.0 41.0

2 5.0 3.0 42.0

3 5.0 2.0 43.0

4 5.0 1.0 44.0

4. For each beaker, rinse a cuvette with 1 mL portions of that beaker’s solution twice (discard the rinsings), then fill the cuvette to the mark with the solution. You should now have 4 filled cuvettes, one for each beaker. 5. Measure the absorbance of each cuvette solution, using the Colorimeter:

a. Open the Colorimeter lid, insert the cuvette for Beaker 1, and close the lid. As before, if the cuvette has frosted sides, be sure you insert it so that the beam passes through the clear sides.

b. Write down the absorbance for that cuvette in your first data table. (If the number never totally stops changing, then write down the number after 10-15 seconds, and move on.)

c. Repeat steps (a) and (b) for the cuvettes for Beakers 2, 3, and 4. 6. Measure and record the temperature of one of the solutions (since all the solutions are room temperature, we assume they all have the same temperature). 7. Either now, or after the experiment, plot the data from your first data table in Excel (or any other graphing program), with [FeSCN2+] on the x-axis, and absorbance on the y-axis. Add a linear trendline for the data, display the trendline equation, and write down the trendline equation below your first data table. You will need this equation for your calculations later.

8. (Making solutions for finding K) Obtain 3 more small beakers (or rinse out the previous beakers and reuse them). Mix solutions in each beaker as shown in the table below. Mix each solution thoroughly. Note: as before, the products of this reaction begin forming immediately, but they degrade in light, so be sure to measure the absorbance of each solution within 2-5 minutes of mixing the chemicals. *Please note that you are using 0.0020 M Fe(NO3)3 for these solutions, NOT the 0.200 M used previously.

Beaker 0.0020 M Fe(NO3)3 (mL)

0.0020 M SCN- (mL)

Distilled H2O (mL)

A 3.00 3.00 4.00

B 3.00 4.00 3.00

C 3.00 5.00 2.00

9. Wash out and reuse 3 of the cuvettes from the previous trials. Wash them out with distilled water. For each beaker, rinse a cuvette with 1 mL portions of that beaker’s solution twice (discard the rinsings), then fill the cuvette to the mark with the solution. You should now have 3 filled cuvettes, one for each beaker. Remember the cuvette tips at the beginning of the procedure. 10. Measure the absorbance of each solution:

d. Open the Colorimeter lid, insert the cuvette for Beaker A, and close the lid. As before, if the cuvette has frosted sides, be sure you insert it so that the beam passes through the clear sides.

e. Write down the absorbance for that cuvette in your second data table. f. Repeat steps (a) and (b) for the cuvettes for Beakers B and C. g. Use your trendline equation from step 7 to find the equilibrium concentration of [FeSCN2+]eq in beakers A, B and

C: for each beaker, plug in the absorbance as y, and solve for the x value, which is the [FeSCN2+]eq value for that

beaker.

11. Measure the temperature of one of the solutions, as before.

Disposal

The contents of all beakers, cuvettes, and pipets may be disposed of in the sink. Flush with a large volume of water.

Name_________________________

CHM112 Lab – Determination of an Equilibrium Constant – Grading Rubric

Criteria Points possible Points earned

Lab Performance

Printed lab handout and rubric was brought to lab

3

Initial concentrations completed before coming to lab.

2

Safety and proper waste disposal procedures observed

2

Followed procedure correctly without depending too much on instructor or lab partner

3

Work space and glassware was cleaned up

1

Lab Report

ICE tables and K calculations complete and shown in detail.

5

Question 1

1

Question 2

1

Question 3

1

Question 4

1

Total

20

Subject to additional penalties at the discretion of the instructor.

Determination of an Equilibrium Constant: Data and Calculations Name ____________________________

Raw Data Tables:

Standard Solutions: Solutions for finding K:

Beaker [FeSCN2+]

(copy this from your

prelab calculations)

Absorbance

1

2

3

4

Trendline equation:____________________________

ICE Charts (Fill in charts and show all the calculations clearly)

Find the initial [Fe3+] and [SCN-] by calculating from the amount of Fe(NO3)3 and KSCN added to each beaker, using M1V1

= M2V2 (much like in the prelab calculations). There is no product at the beginning of the reaction, so initial [FeSCN2+] =

0. The equilibrium [FeSCN2+] is the [FeSCN2+]eq you calculated for each beaker in the second data table. Use the ICE

charts to find the equilibrium concentrations of every chemical, and use the equilibrium concentrations to find K for

each beaker.

1. From Beaker A

Fe3+ + SCN- ⇌ FeSCN2+

Initial

0

Change

Equilibrium

K=

Beaker Absorbance [FeSCN2+]eq

A

B

C

Determination of an Equilibrium Constant: Calculations Name ____________________________

2. From Beaker B

Fe3+ + SCN- ⇌ FeSCN2+

Initial

0

Change

Equilibrium

K=

3. From Beaker C

Fe3+ + SCN- ⇌ FeSCN2+

Initial

0

Change

Equilibrium

K=

Average K= __________________

Determination of an Equilibrium Constant: Post Lab Name ____________________________

1. Consider the following hypothetical reaction at equilibrium:

2A (g) + B (g) ⇌ 2C (g)

At 400 K, the value of the equilibrium constant Kc = 4.15 × 107

Based on this Kc, does the equilibrium favor the reactants or the products? Why?

2. The reaction you studied was Fe3+ (aq) + SCN-(aq) ⇌ FeSCN2+ (aq) Based on your calculated K value, calculate the value of K for the following reaction

2 FeSCN2+ (aq) ⇌ 2 Fe 3+ (aq) + 2 SCN- (aq)

3. Suppose you had prepared Beaker D, where the initial concentration of [SCN-] = 0.0015 M and [Fe3+] = 0.0015 M. Use your average K to predict the equilibrium [FeSCN2+] of the resulting solution.

4. Which, if any, of the following would affect the value of the equilibrium constant found in this lab? It is enough to say whether they would change K at all, or not.

a) adding a catalyst to the system

b) changing the temperature.