Chesapeake Campus – Chemistry 112 Laboratory

LAB #4 – Absorption Spectroscopy

Objectives

After completing this lab you will be able to:

• Determine the wavelength of maximum absorbance for copper solutionsusing a spectrophotometer. Create a calibration curve using absorbance from a series of standards.
• Calculate the molarity of an unknown by using the equation of best-fit fromyour calibration curve.
• Introduction Absorption Spectroscopy Absorption spectroscopy is the study of light absorbed by chemicals. Molecular absorption can occur only when a photo of sufficient energy strikes a molecule causing an electron to be moved from a lower energy level to a higher energy level. As the electron returns to its ground state, energy is reflected off that (often) corresponds to a visible region of light. In this lab, students will create absorption spectra of metal cations by shining light of specific wavelengths through a sample and monitoring the absorbance. Since absorption spectra are unique to each compound, it can be used to identify substances. More importantly, the aborbance of light by a substance is directly proportional to the concentration of the substance in solution. This allows for the calculation of the concentration for a solution of unknown concentration through the use of standards (known concentrations) and a graph.Color and Absorption of Light:
  Although absorption spectroscopy can be performed using light from any region of the electromagnetic spectrum, the spectrophotometer used in this lab operates primarily in the visible region. Therefore we will confine our discussion to visible light. The color we perceive when viewing an object is a direct response to the photons of light that interact with our eyes. When a photon of a specific wavelength reaches our retinas, we are able to view a specific color corresponding to that photon. It does not matter whether the object is itself emitting light or simply reflecting it back. If a wavelength of light reaches our eyes, we will perceive that color. So, in a way, our eyes are acting as spectrophotometers.

  White light is an equal mixture of light of all wavelengths (colors). When such light strikes an object and is completely reflected, light of every color reaches our eyes and we see equal amounts of light of all colors and perceive the object to be white. White light can be split by a prism into a rainbow. When all light striking an object is absorbed, no light is reflected to our eyes and we perceive the object to be black. A sheet of paper is white because all light striking it is reflected and none is absorbed. The print on the paper is black because all light striking it is absorbed; none is reflected. We perceive color when some wavelengths of light are reflected (or transmitted, as in the case of a solution) more than others.

Absorption spectra indicate which wavelengths of light are being absorbed or reflected by an object and are often complex patterns. The statement that “an object appears red because all red light is reflected and all other light is absorbed” is an extreme over-simplification of the truth. In general small amounts of all wavelengths are absorbed by most objects. In general the colors we see are a mixture of the photons actually reaching our eyes. Consider how there are numerous shades of blue. Some blues are warmer (closer to purple) while others are more cool (true blue or having a green tint) in comparison. We perceive color based on the mixture of wavelengths that reach our eyes. Consider how one toddler may call teal “green” while another may call it “blue.” Therefore all colors are generally a mixture of wavelengths that are interacting with our retinas.

Spectrophotometric Analysis
Absorption spectroscopy is a useful tool for performing both qualitative analysis and quantitative analysis. Here we can focus a beam of light of specific wavelengths through a sample to determine if the wavelength is absorbed or reflected. A plot is made of the amount of light absorbed as a function of the wavelength of the light. The wavelengths of maximal absorbance (where light is absorbed the strongest) are characteristic of particular molecular species. These wavelengths can be used to identify unknowns. In addition, the amount of light absorbed is directly proportional to the concentration of the absorbing species. It can be used to determine the amount of a species present in a sample. Furthermore, we can use the information from these plots to draw conclusions about the bonds and molecular geometry of the analyte. Absorption spectra are often characterized by the percent transmittance at a given wavelength.

This is defined as:

I
%T100I  (1)

0

where I is the intensity of light transmitted by the sample and I0 is the intensity of light incident on the sample. When the sample is in solution and a cuvette or sample container must be used, I is taken to be the intensity of light transmitted by the cuvette when it contains a sample solution, while I0 is taken to be the intensity of light transmitted by the cuvette filled with pure solvent (blank solution).

Another way of describing spectra is in terms of the absorbance, A, where:

I
A completely transparent sample would have %T = 100 or A = 0, while a completely opaque

sample would be %T = 0 or A = ∞.
The absorbance, A, is related to the path length, b, of the sample and the concentration,

I
Alog 0  (2)

c, of absorbing molecules by the Beer-Lambert law (frequently referred to as “Beer’s Law”): A = εbc (3)

where the proportionality constant ε is called the absorptivity or extinction coefficient. When the concentration is expressed in moles per liter, ε is the molar absorptivity coefficient. The quantity ε is a property of the absorbing material which varies with wavelength in a characteristic manner. It should be clear from equation (3) that, if you know the molar absorptivity of a species at a particular wavelength, it should be possible to determine its concentration in a solution by measuring the absorbance of that solution. Graphing the absorbance vs concentration allows us to find a line of best fit (y = mx +b) of A = mc + b. Where the absorbance is related to the concentration by the slope (the molar absorptivity).

Spectrophotometers
Instruments used to determine the amount of light absorbed at a particular wavelength

are known as spectrophotometers. These are constructed so that the sample to be studied can be irradiated with light or other radiation of known wavelength and intensity. The wavelength can be varied automatically or continuously by the operator, and the amount of radiation absorbed or transmitted by the sample determined for each wavelength used. In this way it is possible to learn which wavelengths of radiation are absorbed by the sample and how effective the species in the sample is in absorbing a particular wavelength. From this information, an absorption spectrum for a species can be obtained and used to identify the species in unknown samples. For example, it is known that the Co(NH3)63+ ion absorbs wavelengths between 350 and 400 nm and between 450 and 520 nm. If this ion is present in an unknown mixture and the mixture is placed in a spectrophotometer, one would expect to see absorptions in the two wavelength regions indicated.

Spectrophotometers record the amount of light transmitted or absorbed by a sample in terms of the percent transmittance, %T, and/or the absorbance, A. The %T tells us in percentage units how much light is transmitted by (passes through) the sample. Absorbance is a more complex term; it conveys information about the amount of light absorbed in logarithmic terms. (See equations 1 & 2)

The spectrophotometer contains a lamp which emits all wavelengths of light in the range we will be studying. The light from this bulb then passes through several parts: the entrance slit, objective lens, grating, and exit slit. The combination of parts functions as a monochromator, a device which selects only one wavelength or one color of light from all the wavelengths/colors emitted by the bulb. A particular wavelength is selected, using the wavelength control, by adjusting the angle of the grating. This works because different wavelengths are reflected off the grating at different angles, much like light transmitted through a prism of glass or quartz. The wavelength we have selected is then passed through a sample. The amount of light transmitted through the sample is measured by the phototube which generates an electrical current proportional to the intensity of the light striking it. The size of this current, proportional to the transmitted light intensity, is indicated on the display at the top of the instrument.

There is one final consideration that must be made in measuring the light absorbed by a solution. Since the light passes through the sample cells, care must be taken to see that the cells do not affect the measurement. The cells must be constructed of absolutely clear glass or quartz. Since the absorbance is dependent on the path length (see Beer’s Law) they must be made to exact dimensions. Sample cells made to such rigorous specifications are known as “cuvettes”. These tubes are not exactly uniform, but they will work as long as they are reasonably clean. Dirty test tubes will adversely affect your results.

Experimental Procedure

A. Preparation of the Spectrophotometer

1. Make sure the spectrophotometer is turned on and let it warm up for at least fifteen minutes before making any measurements.
2. Set the initial wavelength to 360 nm using the buttons on the front of the instrument.
3. Make sure the instrument is set to measure %T (transmittance). You can set themode using the “mode” button.
4. Adjust the display to 100 %T. The instrument will read BLA—while calibrating. Oncethe Transmittance value appears, change the mode to Absorbance to record thisvalue in your data.
5. Every time you change the wavelength OR open the top of the spectrophotometer youMUST put the settings back to %T and calibrate to 100% T with your blank solution.
6. The instrument can hold the blank AND 3 solutions at a time. The blank should go inthe first space (lever pushed all the way into the instrument), the samples can go in the other 3 spots. When placing the cuvettes into the instrument, they should only be ~ 1⁄2 full and should have a transparent side from left to right in the instrument. Otherwise the light will not be able to go through the sample.

Absorption Spectra of Metal Cations

1. Obtain several cuvettes and make sure they are clean.
2. Fill one cuvette with deionized water. This will be your “blank.”
3. Fill cuvettes with the metal ion solutions. The FeCl3 will be used to find the absorptionspectra of Fe3+, the NiCl2 will be used to find the absorption spectra of Ni2+ and theCuSO4 will be used to find the absorption spectra of Cu2+.
4. Carefully wipe each cuvette with a Kimwipe® to remove any dirt or grease from theoutside of the tube. From this point on, handle these tubes only at the top, above theliquid level.
5. Place the blank in the sample holder and close the lid.
6. Adjust the %T value to 100%.
7. Change the mode to measure absorbance. The absorption value for your blankshould be close to 0.
8. Pull the lever on the front of the instrument to adjust so that the instrument is nowreading the first sample. Record the absorbance value on the data sheet.
9. Repeat until the you have the absorbance for all three metal cations.
10. Increase the wavelength by 20 nm and repeat steps 5-8. Continue doing this until you have values for wavelengths up to 800 nm. Graph the Absorbance vs. Wavelength for all metal cations onto a single graph (If your instructor agrees, this can be done in Microsoft Excel and attached, otherwise include a manual graph in the space provided). Make sure the graph is labeled appropriately.

Quantitative Spectroscopy (Creating a Calibration Curve)

1. The 0.20 M CuSO4 solution used in Part B will be referred to as “Copper Standard 1” 2. Prepare 4 more copper standards as follows:

1. Copper Standard 2: Obtain additional 0.20 M CuSO4 solution from the tray. Pipet 10 mL of 0.20 M CuSO4 solution into a 50 mL Erlenmeyer flask. Pipet 10 mL of DI water into the same tube and mix well.
2. Copper Standard 3: Pipet 10 mL of Copper Standard 2 into another 50 mL Erlenmeyer flask. Add 10 mL of DI water and mix well.
3. Copper Standard 4: Pipet 10 mL of Copper Standard 3 into another 50 mL Erlenmeyer flask. Add 10 mL of DI water and mix well.

3. Calculate the concentration of all copper standards made with dilution and record the values in your data section.
4. Transfer samples of copper standards to cuvettes. Fill your last cuvette with a solution of copper sulfate of unknown concentration.
5. Use your data for copper in Part B, find the wavelength with the highest absorbance for copper and copy the data to this section.
6. Set the wavelength on the instrument to 10 nm less than the wavelength with the highest absorbance. Place the blank and copper Standard 1 in the spectrophotometer and set %T to 100% as you did before. Obtain and record the absorbance.
7. Repeat using the wavelength 10 nm higher than the wavelength selected from part A. [For example, if your copper sample had its highest absorbance at 640 nm, you will take new measurements at 630 nm and 650 nm.] If your wavelengths are below 600 nm.
8. Look at the data for your three wavelengths and set the instrument to the wavelength which produced the highest absorbance. Once again place the blank in the chamber and adjust the %T reading to 100%.
9. Record both absorbance and %T values for all five copper solutions. (The four standards and the unknown.) Remember that when you open the instrument you must recalibrate to 100 %T with your blank.

10.The unknown concentration will be determined using a graph (calibration curve).

Pre-lab Assignment/Questions

N o t e – this pre-lab must be finished before you come to lab. (Please see syllabus for how to submit this assignment.)

1. Explain why “roses are red.”
2. You have a sample of unknown concentration. The species you are evaluating has a molar absorbtivity is 11.02, the path length is exactly 1 centimeter and the absorption is 0.621, calculate the concentration using the equation A = εbc.
3. You create a calibration curve of Absorbance vs. Concentration using a set of standards. You calculate the trendline to be y = 1.23x + 0.024. the 0.024 is the y intercept or the absorbance when the concentration is 0. What could have caused this value to be greater than 0?
4. Is the molarity of a solution independent of temperature? Why or why not? Give a better concentration value for use. Defend your choice.

Experimental Data and Results:

Part B. Absorption Spectra of Metal Cations
Record the absorbance for the metal cations at the wavelengths.

Wavelength

    360
    380
    400
    420
    440
    460
    480
    500
    520
    540
    560
    580

Cu2+

Fe3+

Ni2+

Wavelength

600 620

640 660

    680
    700
    720
    740
    760
    780
    800

Cu2+ Fe3+ Ni2+

Attach a graph of absorbance vs. wavelength using the data for all four solutions. (Absorbance should be on the y-axis.) Clearly label the graph. Include a legend and titles for your graph.

Experimental Data and Results: Part C. Quantitative Spectroscopy

Wavelength used for Analysis in Part C: ___________________

*Show an example calculation obtaining the molarity of the standards.

Before continuing, clean all equipment and glassware. Return all materials to their proper location, and turn off the spectrophotometer. Check this box when this is done and all items in your drawer are accounted for.

Instructor Signature_______________________________________________________________

**Using the absorbance graph on the next page you can determine the concentration of the unknown. Take the absorbance value for the unknown and find that value on the y-axis. Draw a horizontal line from that point on the y-axis to the plotted line. From the point where the horizontal line intersects the plotted line, draw a vertical line straight down to the x-axis. Record the value from the point where you hit the x-axis as the unknown concentration in the table.

Experimental Data and Results

Prepare a graph of Transmittance vs. Concentration using the data for the copper standards. Create it in the space provided or in Microsoft Excel and attach in the space provided (depending on your instructor’s directions).

Prepare a graph of Absorbance vs. Concentration using the data for the copper standards. Create it in the space provided or in Microsoft Excel and attach in the space provided (depending on your instructor’s directions).

Post Lab Questions

*See this syllabus for instructions on how to turn in this section of the lab handout.

1. The presence of a dirty fingerprint on the cuvette during measurement of the sample solution resulted in the absorbance being reported incorrectly. Do you think the number reported was too high or too low? Explain why.

2. List the important sources of error in this experiment and what effect each would have on the results. Discuss in particular any errors that you may have made. Include at least 3 (with specific implications) for full credit.

3. A student makes a solution for analysis by mixing 5.172 grams of Cu2SO4 into 500. grams of water. The density of water at the temperature of the lab at the time the solution was made is 0.9926 g/mL. Calculate the molarity, molality, mole fraction, and mass percent of the copper I sulfate. Show all work.