Lab 08 – Very Green Fluorescence

This lab was inspired by:

Popescu, S. A., & Peled, A. (2023). Optimized RED spectral band fluorescence of edible plants leaves extracts.
Applied Surface Science Advances, 13, Article 100385. https://doi.org/10.1016/j.apsadv.2023.100385

About Fluorescence

In labs 2 and 3, we identified absorbance of visible light as a useful way to measure concentrations in solutions using standard curves for both molecular absorption and atomic emission. Standard curves can be generated for any method that predictably relates a measurement, like absorbance at a specific wavelength, to determine a property, like concentration. Many of these happen without us even realizing it – scales relate the amount of force needed from an electromagnet to offset the pull of gravity, which is measured by the amount of current drawn for the electromagnet. So a scale actually measures the current needed to match the weight of gravity, but it is calibrated so we can make a determination of the mass.

Often the types of signals we can get in a solution are complex. When we generated a standard curve for fluorescein absorbance in Lab 2, we only had fluorescein, sodium hydroxide, and water in solution. When we generated a standard curve in Lab 4, we had a more complex indicator, but a clear zinc control. But the world is full of colored compounds – how can we make sure that signal only comes from a molecule we are interested in (analyte) instead of one of those compounds?

We generally look for relationships that are selective – giving more signal for one analyte than others – or, even better, specific – giving signal for only the analyte of interest. For example, the thin lines of atomic emissions are often fairly specific to an element if the instrument has good resolution between wavelengths. On the other hand molecular absorptions of the chelators like the catechols we used in Lab 4 are formed only when there are metals in solution but give signal for many metals making them selective towards metals but not specific.

Figure 1: Energy changes during molecular spectroscopy are dependent on electronic transitions, and to a smaller degree vibrations and rotations. (CC BY 3.0; OpenStax)

Figure 2: Jablonski diagram of absorbance, non-radiative decay, and fluorescence. (Public Domain, Jacobkhed via LibreTexts.org)

Absorbance of ultra-violet or visible (UV-Vis) light is an incredibly common phenomenon, with many compounds having absorbance at any given wavelength. This is because the UV-Visible region of light matches with the amount of energy required to move electrons between the HOMO and an unoccupied orbital of many organic molecules and inorganic compounds. The breadth of the signal in molecular UV-Vis absorptions is because there are slight changes in the energy required as a molecular bonds vibrates or rotates. Recall that the energy stored in a bond depends on the distance between atoms as shown in Figure 1. Atomic absorptions tend to look more like lines because atoms don’t have bonds to vibrate. In Figure 2, the difference in energy between the singlet state (S, indicating no unpaired electrons) HOMO and unoccupied orbital is indicated by S₀ and S₁, while each of the lines 0, 1, 2, 3… indicate the vibrational states. Typically, when a molecule absorbs energy, it can absorb energy to increase both the electronic state energy (i.e., S₀, S₁, etc.) as well as the vibrational state. When many vibrational modes are available, relaxation often occurs quickly through vibrational states, and energy is released as heat. However, when there are few vibrational states available, like in rigid circular molecules, relaxation occurs rapidly within available states, followed by relaxation through the emission of a visible photon to reach an electronic ground state (S₀), a phenomenon known as fluorescence. Since some energy is lost to vibration, the energy of the fluorescence is typically less than that of absorption and the wavelength is longer.

Luminescence is a general term used to describe the emission of light (a photon) of some wavelength as a material relaxes from an electronic excited state back to the ground state. There are many ways a material can be excited, but in the context of this experiment we are particularly interested in photoluminescence (or PL), where excitation is the result of absorption of a photon of some, usually shorter, wavelength.

Now, why does any of this matter? The usefulness of fluorescence spectroscopy stems from two main factors. First, since most molecules do not exhibit fluorescence the background fluorescence signal is often very small. Therefore, the technique can detect specific molecules with high sensitivity. Second, the spectral properties of the fluorescence emission are usually highly sensitive to local environment (solvent, pH, ionic strength, etc.). Thus, the use of fluorescence is a key technology in applications where local environments can vary quite dramatically, for example in the study of macromolecule systems and biological specimens.

Chlorophyll, well known for its light absorption properties, has fluorescent properties. There are several types of chlorophyll. We will be measuring chlorophyll A (Figure 3), from an undifferentiated standard because it absorbs light in the right range for our fluorimeters.

Figure 3: Chlorophyll A has an extended network of π bonds that create a fluorescing rigid structure.

Procedure:

1. Choose a green you want to test. Use a scissors to cut 10 grams of material it into small pieces. Grind in a mortar thoroughly.

2. Add 30 mL of 95% ethanol to the mortar and briefly grind.

3. Vacuum filter using a Buchner funnel into a clean side-arm flask.

4. After the material has passed through the filter, remove from the vacuum immediately and transfer into a 50 mL beaker. Cover with parafilm when not in use.

5. Prepare a chlorophyll solution from 0.500 g of chlorophyll dissolved to 1.000 L in 95% ethanol. Calculate the concentration of this buffer in µg/cc (the units from the original literature) and record this as the stock standard concentration in your local copy of the linked Excel spreadsheet.

6. Prepare a 100-fold dilution of the stock standard into a 95% ethanol. You will probably want to pour out a small amount of the stock standard into a beaker to make measuring easier. You want your final volume to be 100 mL in a volumetric flask. Calculate the concentration of this buffer in units of µg/cc. Show your work in your Excel spreadsheet. This is your working standard.

7. Setup Vernier GoDirect SpectroVis Plus by plugging the USB into the instrument and your LabQuest. On the LabQuest, click on the red banner and select “Change Mode”. Select “Fluorescence”. For excitation wavelength, select 405 nm. For emission wavelength, select 674 nm. The SpectroVis may alter this slightly to a number near 674 nm. This is normal and not a concern, but you should make sure to take note of the actual number.

8. Obtain a plastic fluorescence cuvette, touching only the very top, and fill about ¾ full with 95% ethanol, and carefully place in the SpectroVis so that the clear sides are lined up with the light and detector. Calibrate the spec by selecting the following menus/buttons with the stylus (NEVER A PEN OR PENCIL) (Sensors / Calibrate / USB:Spectrometer / {wait for the warm up period} / Finish Calibration / OK). Before your begin to make measurements, go to the gear on the top righthand side of the window. Place your blank (either buffer or water) in to the cuvette holder, select an integration time of 100 ms (this increases the signal intensity by adding the signals over this time), and temporal averaging of 5 (this reduces the variability by taking replicate measurements automatically), and confirm excitation wavelength as 405 nm.

9. Label a white piece of paper with numbers 1-6, spaced enough that each number could have a cuvette placed on top. Using a calibrated micropipette, add the following amounts of ethanol and your working standard chlorophyll solution to create cuvettes 1-6.

10. Remove the blank cuvette. Test the fluorescence intensity of each cuvette. Record your data into your local copy of the linked Excel spreadsheet. You will have to calculate out the concentration of each cell.

11. Create a linear regression plot for the standard curve. Add a trendline and show both the equation and R-squared value.

12. Take a fluorescence measurement of your extracted plant chlorophyll. Does the fluorescence intensity fall into the range of your standards? If so, proceed to step 15. If not, dilute 10-fold and repeat this step. Continue diluting and measuring your sample until it falls within the range of your standards. Make sure you record how many dilutions you needed! Record this data under “Sample Fluorescence”

13. Now that you know how many dilutions you need, take two more samples and dilute them to the proper concentration for measurement, for a total of 3 measurements of 3 distinct samples of your chlorophyll (not the same sample 3 times). Convert each to concentrations using your standard curve, then calculate the average and standard deviation for your samples only.

14. If you diluted your sample, find the average concentration of the original, undiluted sample. Note that the standard deviation should scale by the same value. Record both in the section “Values for undiluted samples:”

15. Record your average and standard deviation on the online copy of the linked Excel spreadsheet on the “Class Data” sheet.

Clean up:

All solutions should be dumped into the large “Green Clean” collection beaker in the hood.

Post-lab questions:

1. Using the data shared by your classmates for the same plant, calculate the average value and the standard deviation. Is this technique precise?

2. Your concentrations are in µg (of chlorophyll) / cc (solution), but most analyses of extracted analyte want to know how much you had in the original mass of sample. Could you calculate this directly from your data? There is a major source of error if we tried to analyze this way. What is it?

3. Assume that you were able to recover all of the 30 mL of 95% ethanol you originally added. What is the concentration of chlorophyll per gram of your plant?

4. Can you think of a way to make sure you know the volume of solvent that has all your extracted chlorophyll? What would that look like in the experiment