Organic Synthesis:

Solvent-free Synthesis of Coumarin (in-person experiment)

Purpose

To synthesize and characterize an organic compound.

Learning Objectives

Practice elementary organic synthesis.

Understand the importance of green organic synthesis in medicine, pharmaceuticals and industry.

Characterize the products by IR, melting point, and qualitatively by fluorescence.

Background

In this lab, you will be performing a Synthetic Organic Chemistry experiment where you will synthesize a com-

pound called 4-methylumbelliferone under solvent-free conditions.

How to draw an organic molecule?

You may have seen and drawn organic compounds following the Kekulé structures, which are similar to Lewis

structures, but may not always include lone pairs of electrons (Figure BU.1, left). Furthermore, the structures are

drawn in a linear fashion with all the Carbon and Hydrogen atoms explicitly showcased (Figure BU.1, left side).

These types of structure are very helpful when you want to look at every single atom. However, when dealing

with complex organic molecules, Kekulé representations could cause the structures to be too crowded and hard

to understand. Moreover, showing all the hydrogen atomsmakes it difficult to compare the overall structure with

other similar molecules. You can’t focus on the double bonds or additional groups. Therefore, organic molecules

are often represented in a less detailed manner by condensed structural formulas called a bond-line (or zig-zag)

formula where the chemical symbols for carbon atoms and the hydrogen atoms bonded to these carbon atoms

are omitted (Figure BU.1, right). This simplified representation allows us to easily identify C-C bonds, double

bonds, or triple bonds as well as any additional groups sticking off the main ring or chain. In addition, it is much

quicker to draw the zig-zag structure. If you take an organic chemistry course, you will learn to appreciate this

type of formula writing.

The bond-line formula will be depicted throughout this experimental procedure, you can learn more about these

representations in Section 1.8 of Organic Chemistry (Wade).

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Figure BU.1: Selected xamples of Kekulé structures (left) and their corresponding bond-line representations

(right).

Introduction

Figure BU.2: Structure and

number scheme of coumarin

(1).

Coumarins are natural products that can be found in a variety of plants.

Coumarin (Figure BU.2), was first isolated by Vogel1, in 1820, from tonka

beans, also known as Coumarou, a vernacular French name. Since then, iso-

lation, structural characterization, synthesis and biological activity of thou-

sands of natural coumarins from plants, bacteria, fungi, and chemical syn-

thesis, have been reported.2

You can recognize coumarins by their general structure, which has two six-

membered rings fused together, with one of the rings being a benzene ring

and the other containing an alkene functionality and an ester functional

group inside the ring (Figure BU.3).

Figure BU.3: General structure of coumarins.

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Since their discovery in the 19th century and given their presence in several food sources, the biological properties

of coumarins and their derivatives have been widely investigated, with different molecules showcasing different

pharmacological properties such as anticoagulant (“blood thinners”), anti-inflammatory, antifungal, antibacte-

rial, anticancer and anti-oxidant properties (a couple of selected examples are illustrated in Figure BU.4).3–6

More recently, given their optical properties – fluorescence – coumarins have been employed as laser dyes in the

blue and green regions, OLEDs (Organic Light Emitting Diodes) and bio-probes for physiological tests.7,8

Figure BU.4: Selected examples of biologically active coumarins: an anti-inflammatory

columbianadin 2, and an antibiotic novobiocin 3 (coumarin scaffold highlighted in blue)

Given the wide range of applications of coumarins, several methods have been developed over the years to syn-

thesize these compounds and their derivatives, with increasing focus on milder reaction conditions which allow

for greener paths to obtain these molecules.7

Green Chemistry

High demand for the synthesis of the compounds, which have found their application inmedicine pharmacology,

cosmetics or food industry, results in an increased generation of different waste chemicals. In order to minimize

the utilization and generation of toxic organic materials, green chemistry methods have been developed. They

cover a wide range of methods, including the application of ultrasound and microwaves, multicomponent reac-

tions, and solvent-free synthesis.

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Solvent-free synthesis

Most reactions occur in solutions. When reagents are dissolved in solvents, they are able to easily mix and come

in contact with each other. As all reactions need energy, reaction mixtures in solution also allow the heat energy

to be easily distributed. However, there are some disadvantages to using solvents. Primarily, solvents make up

most of the chemical waste from a reaction, and a preventing of chemical waste is one of 12 principles of Green

Chemistry.

Avoidance of an organic solvent eliminates the associated hazards as well as reduces waste. The purification of

the compounds becomes easier.

In this experiment, youwill be synthesizing a coumarin named 4-methylumbelliferone 4 in a solvent-free reaction

and using an acidic solid resin (Dowex 50WX4) which will aid in the reaction between resorcinol 5 and ethyl

acetoacetate 6 (Scheme 1, Figure BU.5). This procedure is based on the work reported by Kenneth M. Doxsee

et. al.8 The target compound belongs to the family of coumarins named umbelliferone, or 7-hydroxycoumarin,

which are found in various plants of the family Umbelliferae, including carrots, parsley, cumin, and celery.

Figure BU.5: Synthesis of 4-methylumbelliferone (4) that you will be performing in the lab, follow-

ing an adapted procedure by Doxsee et. al. Scheme 1.

Umbelliferone and its derivatives have recently found applications in fluorometric enzyme essays13 and as blood-

brain barrier probes14, anti-inflammatory agents15, dyes16, and fluorescent pH indicators.14,17

Fluorescence

In experiment 2 (CHEM 1120), we discussed light and its properties. Recall that after absorbing energy (for

example, photon of light), an electron can be excited from a lower energy level (ground state) to a higher allowed

energy level (excited state). Excited states are short-lived with a lifetime at about 10−8 sec. After a molecule has

been promoted to an excited state by absorption of a photon, it is in a non-equilibrium state and will eventually

dissipate the energy gained and return to the ground state. The first way that energy is lost is through vibrational

relaxation, where excess vibrational energy is lost to vibrationalmodes within the samemolecule (intramolecular

process) or to surrounding molecules (intermolecular) via collisions, until the lowest vibrational level of the

electronic state is reached. Vibrational relaxation occurs rapidly, on a time scale of 10−12 – 10−10 and outcompetes

all other transitions.

At the next step, an excited molecule returns to its ground state by emitting a photon of a lower energy, which

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corresponds to a longer wavelength, than the absorbed photon. Fluorescence is one of the possible processes by

which a molecule can relax to a lower energy state, after being excited to higher energy level through absorption

of a photon of appropriate wavelength. Fluorescence occurs on a timescale of 10−10 to 10−7 s.

Watch the following videos to learn more about the fluorescence process:

• How does fluorescence work?

• Introduction to Fluorescence.

The fluorescence phenomenon is very common in molecules that contain extended 𝜋-systems (multiple double

bonds in conjugation with one another) which decreases the energy gaps between excited energy levels. Perhaps

you have heard that certain scorpions emit a “glow” when exposed to ultraviolet light. This glow is a form of

fluorescence. The coumarin compound that we will synthesize will not only fluoresce but will fluoresce with

different colors under different pHs. So, in essence, we will create a type of UV pH indicator!

Characterization of Synthesized Coumarin

After the sample isolation is complete, measure the amount of sample collected and then measure its melting

point and record an IR spectrum of the sample.

Melting point determination

The melting temperature is a characteristic physical property of a substance; it is the temperature at which solid

and liquid phases are in equilibrium. When heat is applied to a pure solid at its melting point temperature, the

temperature remains constant until the change of state is complete. However, high purity is seldom attainedwith

laboratory grade organic compounds and conventionalmethods. For this reason, amelting temperature is seldom

a point but instead it is a range. If nearly pure, the melting temperature range of an organic compound will be

small (0.5 - 1.0°C). The presence of even small amounts of impurities usually lowers the melting temperature

measurably and widens the melting temperature range.

The standard process for determining melting points is the capillary method. In this method, a small amount

of solid is packed into a thin glass capillary tube. The capillary is then placed in a melting point apparatus (Fig-

ure BU.6) and heated slowly near an accurate thermometer. The temperature range over which the solid visibly

softens and finally liquefies is the melting point range, which is commonly abbreviated as m.p.

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Figure BU.6: SRS DigiMelt MPA160 apparatus and its keypad

Infrared Spectroscopy

Infrared Spectroscopy and Molecular Vibrations

The bonds between atoms are not static, but they are actually in a state of constant motion called molecular

vibrations; they bend, stretch, and vibrate rapidly. They can do so either symmetrically or asymmetrically (watch

a video The invisible motion of still objects by TED-Ed).

The bonds will absorb the photons of infrared (IR) radiation if they have the right amount of energy to cause a

vibration. It happens when the frequency of the IR radiation equals that of the vibrating bond. The absorbed

radiation then serves to increase the amplitude of the vibrations and thus the kinetic energy of the molecule.

Please note: different bonds vibrate with different frequencies, thus they can absorb the frequencies of radiation

that are unique to them. This is the fundamental idea behind IR spectroscopy. You should watch the video on

Labflow called Interpreting IR Spectra.

Infrared, or IR, spectroscopy is a technique to study the bonds and functional groups in a sample to confirm the

structure or identity of an organic compound. It is a highly useful technique for chemists. When IR light is shone

upon a sample, the bonds will absorb only the frequencies that are characteristic to them, while the rest of the

radiation passes onto a detector. The detector will generate a spectrum by sensing the wavelengths which are

hitting it and which aren’t.

A typical IR spectrum is a variation of a plot of absorbance versus energy or frequency. The x-axis of an IR

spectrum is usually presented as wavenumbers in units of inverse centimeters (cm−1). The y-axis of the spectrum

is usually % transmittance, which is the opposite of absorbance. Wavenumbers are simply a way of measuring

and reporting the frequency, while % transmittance refers to the amount of light that the detector is sensing at

that wavelength. At 100% transmittance, all the incident light reaches the detector. As bonds absorb radiation,

the amount of detected light decreases, generating a “peak” in the spectrum by lowering the % of transmittance.

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An example of the IR spectrumof acetaminophen (the active ingredient in pain relievers such isTylenol) is shown

in Figure BU.7. While the spectrummay look complicated at first, a trained chemist could immediately recognize

that this compound contains O-H andC=Obonds as well as other functional groups. It becomes possible because

these bonds have very different and distinct absorption frequencies.

Figure BU.7: IR spectrum of Acetaminophen Acetaminophen product information Sigma-aldrich: St-Louis,

MO, Accessed on Oct 10, 2021.

Functional Groups in IR Spectra

The bonds in a molecule vibrate, changing their length and angle when the right energy is available. IR radiation

can trigger these different vibrations in a molecule. The vibration energy depends on the atoms involved in the

bond and the type of bond between them. If the mass of one of the atoms increases, the bond vibrates at a lower

frequency and energy. Another factor that determines the energy of a vibration is the strength of the bond. In

general, a stronger bond vibrates at a higher frequency.

The general guidelines for the influences on vibrations canhelp us connect IR bands to the featureswithin organic

structures, as shown in Figure BU.8. Furthermore, the combination of multiple bond types can be used to identify

functional groups.

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Figure BU.8: IR Absorption Ranges by Bond Type

IR spectra can provide a good sense of the bonds present and thus the functional groups in a molecule. However,

an IR spectrum can be very complicated. Therefore, the best strategy, especially when starting out, is to look for

strong signature bands before looking into finer details.

How to read an IR spectrum

A typical IR spectrum can be visually divided into two regions. The part above 2000 cm−1, usually contains

relatively few peaks, but they have some useful diagnostic information. One of the most distinct and easily rec-

ognizable peaks in an IR spectrum is a broad peak between 3100 and 3600 cm−1 corresponding toO−Habsorption

of alcohols and phenols (Table BU.1). The characteristic frequencies from 2800 to 2000 cm−1 are used to identify

C≡C or C≡N bonds.

In contrast, the complexity of infrared spectra in the region below 2000 cm−1 makes it difficult to assign all the

absorption bands. Two signals, which can be identified in this area are the carbonyl group (C=O), with a very

strong peak around 1700 cm−1, and a C−O bond with one or two strong peaks around 1200 cm−1. This complex

region is also known as the “fingerprint region” because compounds produce a unique pattern in this area. By

comparing absorbances in this region to a known spectrum, we can identify an unknown compound.

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How are IR spectra recorded?

(Section cited from Luong, H. CHEM 2122 Laboratory Manual)

In short, an IR spectrophotometer will measure the amount of light that is transmitted through the sample (in%

transmission). There are several commonly encountered IR sampling methods:

• Thin film or neat: A sample is sandwiched between NaCl plates. If the sample is a pure liquid then

this technique is known as “neat” or “thin film”. One of the most common problems with the thin film

technique is application of too much sample, which leads to a heavily absorbed spectrum (in the 0-10%

transmission range). To resolve this, some of the sample is wiped off. Likewise too little sample will

give weak peaks (80-90 % transmission).

• KBr pellet: A sample is ground up with KBr and compressed into a clear pellet.

• Nujol mull: If the sample is a solid, then it must be made transparent, and this is usually accomplished

by dissolving the sample in a mull. To prepare a mull, the sample is ground in an agate mortar with

somemineral oil (Nujol) until a clear paste is made. One has to be very careful with preparing the Nujol

mulls because Nujol is a mixture of hydrocarbons and consequently absorbs in the IR spectrum. You

generally want to minimize its absorption by using as little as possible. If you are given a spectrum of

a sample prepared in Nujol mull, you will need to identify which peaks are Nujol and therefore do not

correspond to the sample structure.

• ATR (AttenuatedTotal Reflectance): The IR spectrometer for use in our first-year chemistry laboratories

has anATR samplingmethod. This is a useful and convenientmethod because the sample only requires

intimate contact with a crystal. There is no need for any special sample preparation.

The method used will determine the quality of the spectrum and ease of data collection. The present units for

the IR bands are reported as wavenumbers with units of cm−1. The conversion of one unit to another should

be straightforward by examining the relationship between the numbers. The undergraduate first-year Chemistry

Laboratories have five infrared spectrometers. These spectrometers are very expensive (the price of a car!) and

therefore we hope that you will treat the instruments with respect.

The particular method that you will be using to collect the spectral data is called “ATR” (Attenuated Total Re-

flectance). The advantage of thismethod is the efficiency and high turnover rate (it only takes a couple of minutes

for the user to obtain a spectrum). In the not-too-distant past, sample preparation for recording IR spectra was

tremendously more involved and time consuming. ATRworks by crushing a sample of the analyte on a diamond

surface so that there is intimate contact. A bad quality IR spectrum is the result of a sample not completely cov-

ering the diamond surface. Bad quality IR spectra are characterized by a low signal to base-line noise ratio and

basically they appear “hairy”.

In the spectra below, can you figure out which is acceptable, and which is not? You should be able to immediately

assess whether a spectrum displayed on the instrument is acceptable or not because it will save you time in the

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end (when you go to analyze your spectrum after the lab, you may discover that it’s unacceptable and therefore

would need to run it again.) So what do you do if the spectrum quality is poor (Figure BU.9)? Chances are if

you lift up on the pressure-applying tip, you’ll see that there’s exposed ATR diamond surface. You should push

your sample around so that it covers the entire diamond surface and then recollect the spectrum. Liquid samples

typically do not show this issue since it effectively covers the sampling cell.

Figure BU.9: IR spectra comparison. Can you figure out which is good quality and which is poor?

How to Interpret an IR spectrum

IR spectroscopy provides information on the functional groups present in the molecule since certain functional

groups (OH, NH and C=O) have characteristic absorptions. The OH group is probably one of the most promi-

nent features in an IR spectrum. Alcohol and phenols will absorb strongly between 3000 and 3700 cm−1 while

carboxylic acid OH absorbs in a broader range due to hydrogen bonding. Amine NH’s are similar in structure to

OH’s and so will normally be found between 3100 and 3500 cm−1. Carbonyls are also characteristic groups in the

IR spectrum because they yield strong and sharp peaks within 1640-1820 cm−1. A short list of common organic

groups that are observed is given in Table BU.1.

Experimental Procedure

Please read BEFORE coming to the lab! You will receive printouts of the procedure in the lab.

Safety Precautions

Standard laboratory safety procedures should be followed. Resorcinol may be harmful if it is absorbed

through open wounds or ingested. Ethyl acetoacetate, Dowex 50WX4 (acid form), and the product, 4-

methylumbelliferone, can act as irritants; use of gloves and eye protection is advised. Ethanol is flammable,

and hydrochloric acid solutions is corrosive and should be handled with care.

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Table BU.1: Common IR frequencies

Bond Functional Group Frequency (cm−1) andIntensity

C=C Alkene 1680-1600 (m-w)

C=C Aromatic 1600 and 1475 (m-w)

C=O Aldehyde, Ketones, Esters, Amides, Carboxylic Acids,Anhydrides

1600-1800 (s)

C-O Alcohols, Ethers, Esters, Carboxylic Acids, Anhydrides 1300-1000 (s)

O-H Alcohols, Phenols 3600-3400 (m)

O-H Carboxylic Acid 3400-2400 (m)

N-H Primary and Secondary Amines 3500-3100 (m)

N-H Amides 1640-1550 (m-s)

C-N Amines 1350-1000 (m-s)

C�N Nitriles 2260-2240 (m)

N=O Nitro (R–NO2) 1550 and 1350 (s)

Intensity abbreviations: s (strong), m (medium), and w (weak).

Activity 1. How to use a digital pipette

Figure BU.10: Digital pipette.

This part must be completed before starting the experiment – all students need to be able to use a digital pipette.

These are delicate and expensive instruments, so follow instructions carefully.

1. Obtain a 10 mL digital pipette and make sure that there is a plastic tip attached to it (Figure BU.10). Liquid

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must never be drawn into the digital pipette itself – this will damage it and contaminate future samples.

Liquid will be drawn only into the plastic tip. The digital pipette tips that you will use in the first year labs

are not disposable – they must be rinsed and reused.

2. Locate the digital pipette’s volume readout (shown in Figure BU.10; some digital pipettes will have the vol-

ume readout on the side). Some digital pipettes have a volume readout that includes a decimal point. For

others, the first two digits indicate volume in mL, the third digit indicates the tenth’s place. For example, a

reading of 055 indicates 5.5 mL.

3. Set the volume you need by gently twisting the plunger (Figure BU.10):

• Clockwise to decrease the volume.

• Counterclockwise to increase the volume.

4. Take the digital pipette, a beaker with about 50mL of distilled water, and a clean and dry beaker to a balance.

5. Place the empty beaker on the balance and tare it.

6. Practice drawing up water into the plastic tip of the digital pipette, then delivering the water to the beaker

on the balance:

a. Tare (zero) the balance tare by pressing the“tare” or “0/T” button before each trial.

b. Set the volume readout to the volume you want to draw up and deliver. For this practice exercise, try

the following three volumes: 1.0 mL, 5.3 mL, and 10.0 mL.

c. While holding the digital pipette vertically, depress the plunger to the first stop, then immerse the tip

deep enough into the water to be able to draw up the desired volume without sucking up air. Slowly

allow the plunger to return to its original position to start filling the tip. Once the plunger reaches the

original position wait 1 to 2 seconds for the tip to finish filling and then remove the tip from the water.

Do not draw liquid into the tip rapidly as it can cause some of the liquid to en-

ter the shaft of the digital pipette, which can damage it and contaminate future

samples.

d. If any droplets have adhered to the external surface of the tip, remove them by sliding the tip against

the interior surface of the beaker from which you drew the liquid.

e. Dispense the water by holding the tip at an angle of 10‒40𝑐𝑖𝑟𝑐 against the side of the beaker on the

balance and then depressing the plunger to the first stop. Wait 1 to 2 seconds and then depress the

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plunger to the second stop to blow-out any remaining water in the tip.

Since water has a density of 1.00 g/mL and the digital pipette has an accuracy of ± 0.05 g, you will have

successfully delivered the correct volume if the mass of water in grams is the same as the volume in mL of

water you delivered, within ± 0.05 g. For example, when delivering 5.3 mL of water, the mass should be

between 5.25 g and 5.35 g. Practice this technique until you can do it consistently.

Activity 2. Synthesis

1. In a clean 50- or 100-mL beaker, combine 1.0 mL of ethyl acetoacetate, 1.0 gram of resorcinol and 1.0 gram

of Dowex. ADD TO YOUR LAB NOTEBOOK.

2. In a 100-mL beaker, add 50 mL of DI water.

3. In a 50-mL beaker, add 25 mL of ethanol.

4. Place all the glassware on a hot plate set to the lowest setting. You do not want to boil the water or ethanol,

but you do need to heat it up. If your samples start to boil, remove them from the hot plate and turn the heat

down and then replace them on the hot plate.

5. Occasionally mix the reaction mixture with the glass stir rod. The reaction will take up to 45 minutes.

6. When the reaction is complete, add around 15 mL of the hot ethanol from the 50- mL beaker to dissolve the

solid in the reaction flask.

7. Make sure that all the grey is dissolved. If it has not, then add more ethanol until it is dissolved.

8. Decant the solution from the Dowex beads to the clean beaker. Repeat steps 6 and 7 up to two more times

by adding small amount of hot ethanol (around 2-3 mL). Keep the beaker warm.

9. Place the beaker on the hot plate.

10. While heating, carefully add hot water to the ethanol solution until the solution starts to become cloudy.

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11. Turn off the hot plate, and allow the reaction mixture to cool to room temperature.

12. Weigh a watch glass with a piece of filter paper and record the mass. ADD TO YOUR LAB NOTEBOOK.

13. Collect the white precipitate from the flask by vacuum filtration.

14. Wash the collected solid with some DI water and allow the sample to dry for 10 minutes.

15. Weigh and record the mass of your product by transferring your solid and filter paper onto the watch glass

you massed previously. ADD TO YOUR LAB NOTEBOOK.

Activity 3. Observing Fluorescence

1. Prepare three clean dry test tubes,

2. Add a small amount of your collected solid to each tube and dissolve the sample in 1.0 mL ethanol.

3. To the second test tube, add 1.0 mL of 6.0 M HCl. To the third test tube, add 1.0 mL of 6.0 M NaOH.

4. Turn on a UV lamp, record your observations. ADD TO YOUR LAB NOTEBOOK.

Activity 4. Melting point

Youwill determine themelting point for 4-methylumbelliferone. Before coming to the lab, lookup and cite in your

lab notebook two reliable literature values for each of these melting points. The University of Manitoba Science

Library’s Chemistry website is a good source of reliable information (Science Library Chemistry website).

1. Partially fill a melting point capillary tube with sample by repeatedly pushing the open end of the tube into

the powder.

2. Pack the material to the bottom of the tube by tapping the closed end of the capillary on the bench top. Aim

for the layer of solid at the bottom of the tube to be 1 – 2 mm deep after packing.

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3. Clean the outside of the melting point tube with a dry towel and bring to a DigiMelt apparatus (Figure 5).

4. Carefully follow the Quick Start Instructions provided on the DigiMelt apparatus.

5. Watch the sample through the observationwindow, paying special attention to it near the theoreticalmelting

point.

6. The melting point range recorded is from the first sign of physical change (darkening, sintering, shriveling,

etc.) of the sample to final formation of completely clear liquid. Record all observations.

7. Allow the apparatus to cool between samples.

Usedmelting point capillary tubes must be disposed of into the broken glass bin. Throw-

ing it into any other garbage bin can lead to injury.

Activity 5. IR spectra Recording

Your TA will explain you how to use an IR instrument in the lab.

References

1. Vogel, A. Darstellung von Benzoesäure Aus Der Tonka-Bohne Und Aus DenMeliloten - Oder Steinklee

- Blumen. Ann. Phys. 1820, 64 (2), 161–166. Vogel.

2. O’Kennedy, R.; Thornes, R. D. Coumarins: Biology, Applications, and Mode of Action; JohnWiley and

Sons: New York, 1997.

3. Stefanachi, A.; Leonetti, F.; Pisani, L.; Catto, M.; Carotti, A. Coumarin: A Natural, Privileged and Ver-

satile Scaffold for Bioactive Compounds. Mol. A J. Synth. Chem. Nat. Prod. Chem. 2018, 23 (2).

Stefanachi et al.

4. Garg, S. S.; Gupta, J.; Sharma, S.; Sahu, D. An Insight into the Therapeutic Applications of Coumarin

Compounds and Their Mechanisms of Action. Eur. J. Pharm. Sci. 2020, 152, 105424. Garg et al.

5. Annunziata, F.; Pinna, C.; Dallavalle, S.; Tamborini, L.; Pinto, A. An Overview of Coumarin as a Versa-

tile and Readily Accessible Scaffold with Broad-Ranging Biological Activities. Int. J. Mol. Sci. 2020, 21

(13), 1–83. Ammimzoata et al.

6. Bouhaoui, A.; Eddahmi, M.; Dib, M.; Khouili, M.; Aires, A.; Catto, M.; Bouissane, L. Synthesis and

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Biological Properties of Coumarin Derivatives. A Review. ChemistrySelect 2021, 6 (24), 5848–5870.

Bouhaoui et al.

7. Cao, D.; Liu, Z.; Verwilst, P.; Koo, S.; Jangjili, P.; Kim, J. S.; Lin, W. Coumarin-Based Small-Molecule

Fluorescent Chemosensors. Chem. Rev. 2019, 119 (18), 10403–10519. Cao et al.

8. Sun, X.; Liu, T.; Sun, J.; Wang, X. Synthesis and Application of Coumarin Fluorescence Probes. RSC

Adv. 2020, 10 (18), 10826–10847. Sun et al.

9. Skoog, D.; Holler, F. J.; Nieman, T. Principles of Instrumental Analysis, 5th ed.; 1998.

10. Jablonski Diagram | What is it? | Edinburgh Instruments Jablonski Diagram (accessed Oct 26, 2021).

11. Gulati, S.; Singh, R.; Sangwan, S. A Review on Convenient Synthesis of Substituted Coumarins Using

Reuseable Solid Acid Catalysts. RSC Adv. 2021, 11 (47), 29130–29155. Gulati et al.

12. Young, D. M.;Welker, J. J. C.; Doxsee, K. M. Green Synthesis of a Fluorescent Natural Product. J. Chem.

Educ. 2010, 88 (3), 319–321. Young et al.

13. Kämpfer, P.; Kulies, I.; Dott, W. Fluorogenic Substrates for Differentiation of Gram-Negative Nonfer-

mentative and Oxidase-Positive Fermentative Bacteria. J. Clin. Microbiol. 1992, 30 (6), 1402–1406.

Kampfer et al.

14. Sundt, T. M.; Anderson, R. E. Umbelliferone as an Intracellular PH-Sensitive Fluorescent Indicator and

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