Exp. 6 - Equilibrium of Fe SCN - Summer 2021 PDF

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Experiment 6

Equilibrium of Fe(SCN)2+

Introduction Many chemical reactions do not go to completion. That is, the reactants do not fully transform into products. The reason for this is that many reactions can proceed both from reactants to products (forward reaction) and from products to reactants (reverse reaction): A→2B and 2B→A When this is the case, rather than the reaction going to completion, an equilibrium is established where the rates of the forward and reverse reactions are equal. While the reaction has not stopped or gone to completion, the concentrations of the chemical species will no longer be changing – as fast as reactant is converted to product an equal amount of product is converted back to reactant. Rather than saying the reaction is finished, we say that it has reached a state of equilibrium. Reactions that can establish equilibrium are written with a double arrow: A⇌2B It is established that for systems at equilibrium, the ratio of products to reactants raised to their coefficients from the chemical equation will be a constant at a specific temperature. This is called the law of mass action. For the general reaction aA + bB ⇌ cC + dD Using the law of mass action, we can establish the equilibrium expression [฀฀฀฀฀฀ ]฀฀[฀฀]฀฀฀฀฀฀ ฀฀฀ ฀ = ฀฀ [฀฀]฀฀฀฀ [฀฀]฀฀฀฀฀฀ Kc = concentration equilibrium constant The subscript (eq) denotes that the values in this equation can only be equilibrium values. Kc is a constant for a chemical reaction at a given temperature. Initial concentrations or pressures, adding or removing reactants or product, or adding a catalyst will have no effect on the value of K. At a given temperature, the concentrations will always adjust themselves (stoichiometrically of course) to establish the appropriate value of K. This Page 1 of 6

may take time if a reaction proceeds slowly, but given enough time, equilibrium will be established and the ratio of reactants to products raised to their coefficients will be the K value. In this experiment, equilibrium will be quantitatively studied for the formation of the Fe(SCN)2+ complex ion. Fe(SCN)2+ is formed through the reversible reaction of Fe3+ with SCN-: Fe3+ (aq) + SCN- (aq) ⇌ Fe(SCN)2+ (aq) As this is an equilibrium reaction [฀฀฀฀฀฀฀฀฀฀ 2+] ฀฀฀ ฀ = [฀฀฀฀ 3+][฀฀฀฀฀฀ −] Through this experiment, you will show that the value of Kc for this reaction is independent of the initial concentrations of reactants. The initial concentrations of iron (III) and thiocyanate ions will be varied and the equilibrium concentrations and resulting equilibrium constant will be determined. If all goes as planned, the K values will be consistent within experimental error. While initial concentrations of the reactants will be easily determined based on the chemicals provided, the equilibrium concentrations must be measured. To accomplish this, it is only necessary to determine the equilibrium concentration of one of the species in the reaction as all change occurs stoichioimetrically. For this reaction, the concentration of the Fe(SCN)2+ product will be measured by spectrophotometry. Spectrophotometry Spectrophotometry is the process of determining the concentration of a substance in a sample by measuring the amount of light at a specific wavelength that it absorbs. A spectrophotometer consists of a lamp emitting light in the entire visible spectrum. The light is focused on a grating, which based on its orientation chooses a single wavelength of light to focus through the sample (or reference). The intensity of the light that passes through the sample is measured using a phototube.

Schematic of a Spectrophotometer

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The ratio of the light that passes through the sample (I) to that light that passes through a reference (Io) is known as the transmittance (T). T = I/Io % T = 100% *T The intensity of the light that will pass through the sample will be less than the light that will pass through the reference, because substances within the sample are able to absorb some of the light’s photons that are not absorbed by the reference. The transmittance has been shown to be dependent on three factors that are related to the substance: ε, the molar absorptivity coefficient, which is a unique constant for every substance at a given wavelength; b, the length of the sample that the light has to pass through; and c, the concentration of the absorbing substance in the solution. The following equation shows the relationship. T = 10-εbc Taking the logarithm of both sides one gets:

log T = - εbc

Rearrange and get:

-logT = εbc

Rearrange again and get:

log(1/T) = εbc

log (1/T) is defined as the absorbance (A):

log (1/T) = A

Combining the previous two equations one gets:

A = εbc

This is the common form of the Beer-Lambert Law, which basically states that the absorbance (A) is directly proportional to the concentration of the absorbing species (c) , the path length (b), and the absorptivity coefficient (ε). One can easily see the application of the Beer-Lambert law in determining solution concentrations as the transmittance (or absorbance) of a sample can be easily measured, and its concentration can be calculated using the above equations. Experimental Determination of the Equilibrium Constant Fe(SCN)2+ is a complex ion that appears orange in color in aqueous solutions. The reason that it appears orange is that it absorbs wavelengths in the blue region of the visible spectrum. For this reason a blue light of 450nm will be used to measure the transmittance of several solutions containing different concentrations of Fe(SCN)2+. According to the BeerLambert law, the amount of light absorbed at a particular wavelength is related to the concentration of the Fe(SCN)2+ in the solution. In order to determine the concentration of a sample, it is required to know the molar absorptivity coefficient and the path length of the sample. While the diameter of the sample tube would be relatively easy to measure and one could look up the molar absorptivity coefficient for Fe(SCN)2+, a better method will be to establish these values from a calibration curve. A calibration curve in this respect is a plot of absorbance versus concentration values for several solutions of known concentrations of Fe(SCN)2+ . The slope of the line will be equivalent to the quantity εb (A = εbc so if y = A and x = c, then εb is the slope) . As Fe(SCN)2+ is in equilibrium with Fe3+ and SCN- one may think it is difficult to know accurate Page 3 of 6

concentrations of Fe(SCN)2+ in a solution if we do not know K. (We do not know K, we are trying to determine K in this experiment). To overcome any limitations that the equilibrium may present, we take advantage of LeChatelier’s principle with respect to the concentration of one of the reactants. Calibration samples will be prepared from 0.2 M Fe3+ and 0.002 M SCN- solutions. Due to the much larger concentration of Fe3+, the equilibrium will shift largely to product meaning essentially that all of the SCN- will be converted to product. So, the concentration of Fe(SCN)2+ at equilibrium can be assumed to be the concentration of SCN- in the solution before equilibrium is established. Once a calibration curve is established, the initial concentrations of Fe3+ and SCN- will be varied and the transmittance of the sample will be measured. From the calibration curve, the value of εb will be determined and with the transmittance data, the concentration of Fe(SCN)2+ in each sample can be readily calculated. Using this equilibrium concentration and the initial concentrations of the reactants, the equilibrium constant for this reaction can be determined.

Procedure

Part 1 – Making Standard solutions. 1.

Into a clean, dry beaker combine the solutions from the table for each calibration solution using the appropriate pipettes. Pour the contents of the beaker into a provided cuvettes (do not fill the cuvettes to the top). Pour any excess from the sample beaker into a waste beaker, rinse the sample beaker, and continue making solutions until you have the blank and four solutions for the calibration. Calibration Solution Blank (Reference) #1 #2 #3 #4

0.2 M Fe(NO3)3 in 0.5M HNO3 5.00 mL 5.00 mL 5.00 5.00 5.00

0.002 M KSCN in 0.5M HNO3 0 0.50 mL 1.00 mL 1.50 2.00

0.5M HNO3 5.00mL 4.5 mL 4 mL 3.5 3

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Part 2 – Establishing the calibration curve. 1. Plug in LabQuest to AC outlet. 2. Turn on the LabQuest, 3. Plug in the Spectrometer into the USB outlet – the LabQuest will show USB: Abs 4. Fill a cuvette ¾ full with the blank (reference) solution and place in cuvette holder 5. Touch the red area on the Labquest and choose “Change Wavelength” 6. Enter 450 nm in the space for the wavelength. Choose “OK” 7. Touch the red area on the Labquest and choose “Calibrate” 8. When warm up is complete choose “Finish Calibration”, after a few seconds choose “OK”. 9. Fill a cuvette ¾ full with your first solution. Make sure the outside of the cuvette is clean and dry. 10. Place the cuvette in the cuvette holder and press the collect button

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11. After a few seconds, press the stop button (or collect button again). 12. Record the absorbance from the box on the right hand side of the screen. (Double check that the wavelength is at or near 450nm in the box at the lower right) 13. Place your next sample in the cuvette holder and press collect to determine its absorbance. Choose “discard” when prompted to save or discard data. Repeat the process to determine the absorbances of each solution. Part 3 –Equilibrium mixtures. Equilibrium Solution Blank (Reference) #1 #2 #3 #4 #5

0.002 M Fe(NO3)3 in 0.5M HNO3 0 5 mL 5 5 5 5

0.002 M KSCN in 0.5M HNO3 0 1 mL 2 3 4 5

0.5M HNO3 10 4 mL 3 2 1 0

1. Into a clean dry beaker combine the solutions from the table for each equilibrium solution using the appropriate pipettes. Pour the contents of the beaker into a provided test tube (do not fill the test tubes to the top!) and place the tube into the test tube rack. Pour any excess from the sample beaker into a waste beaker, rinse the sample beaker and continue making solutions until you have the six solutions (including Blank) in test tubes (or cuvettes). 2. Recalibrate the spectrometer as you did in the previous section; however, you should use your new blank, nitric acid, for the calibration. 3. Record the absorbance of each solution as you did with the calibration solutions. In this case use the new Blank (reference) solution. 4. Pour your waste beaker into the large waste beaker in the hood.

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Data Record all data with appropriate labels and units. The following data tables need to be transcribed into your lab notebook for data collection during lab. Transmittance of Equilibrium Solutions: Transmittance of Equilibrium Solutions: Calibration Solution Absorbance Equilibrium Solution Absorbance #1 #1 #2 #2 #3 #3 #4 #4 Analysis Calibration solutions 1.

Concentration of Fe(SCN)2+ in calibration Solutions –Recall that the concentration of Fe(SCN)2+ in the calibration solutions is essentially the concentration of the SCN- initially in the solution. So, determine the concentration of SCN- initially in each of the 10.0 mL calibration solutions. This is a simple dilution!

2.

Make a scatter plot (on the computer) of absorbance versus concentration for the calibration samples. On the plot, add a linear trendline and display the equation on the chart. Take your time to label the axes. Attach the chart to the end of this experiment.

3. From your plot, what is the value of εb? ________________ Equilibrium solutions 4. For each sample, account for the dilution effect and determine the initial concentrations of the reactants in the 10 mL samples: Equilibrium Solution #1 #2 #3 #4 #5 5.

[KSCN]o /M

For each sample, calculate the equilibrium concentration of Fe(SCN)2+ using the value of εb that you found. Equilibrium Solution #1 #2 #3 #4 #5

6.

[Fe(NO3)3]o /M

Absorbance

Fe(SCN)2+ eq

For each sample, set-up an ICE chart with the information from analysis 4 and 5. Fill in all the blanks and calculate values of Kc for each solution.

7. How well do the values of Kc agree with one another? Should they agree? Do the higher or lower concentrations not fit as well? etc. Questions: 1. In the past, a value of εb was given to the students rather than them having to find it. Aside from the intellectual exercise involved in finding it, why do you believe the value you found is better than using a previously determined value? Page 6 of 6...