Determination and Comparison of Carbonyl Stretching Frequency ofa Ketone in Its Ground State and the First Electronic Excited StateSubhajit Bandyopadhyay* and Saswata Roy

Department of Chemical Sciences, Indian Institute of Science Education and Research (IISER) Kolkata, Mohanpur, West Bengal741252, India

*S Supporting Information

ABSTRACT: This paper describes an inexpensive experiment to determine the carbonyl stretching frequency of an organic ketocompound in its ground state and first electronic excited state. The experiment is simple to execute, clarifies some of thefundamental concepts of spectroscopy, and is appropriate for a basic spectroscopy laboratory course. The experiment iscomplemented by an optional computational component.

KEYWORDS: Upper-Division Undergraduate, Physical Chemistry, Hands-On Learning/Manipulatives, Aldehydes/Ketones,Spectroscopy, UV−Vis Spectroscopy, IR Spectroscopy, Computational Chemistry, Photochemistry, Molecular Modeling

■ INTRODUCTION

Designing an experiment for an undergraduate physicalchemistry laboratory that integrates the concept of electronicenergy levels and the vibrational energy levels, coupled to theestimation of stretching frequency in the first electronic excitedstate, and correlating that with the one in the ground statemight seem like a complicated task. Generally, in spectroscopylab courses, the students carry out discrete experiments indiverse areas of spectroscopy. The “big picture” that integratesthe electronic and vibrational spectroscopy, their nature, andthe energy scale at which they operate are often missed by thestudents. Inspired by a textbook of photochemistry,1 a simpleundergraduate laboratory experiment has been designed thathelps students to link the concepts of IR and UV spectroscopy.Normally, the vibrational stretching frequency of a compoundin its first electronic excited state is determined precisely byUV−IR double resonance spectroscopy of jet cooledmolecules.2 Although the experiment described here lackssuch precision, it does provide a quick and easy means to have arough estimation of the vibrational frequency of a CO bondin its first electronic excited state. The experiment can beperformed easily with readily available resources in anundergraduate laboratory. At the end of the experiment, thestudents had better understanding of some of the fundamentalelements of molecular spectroscopy. The estimated time takento perform this experiment is 40 min, and 10−20 min shouldbe allotted for the calculations. A group of two students wouldbe ideal for conducting this experiment in the class. The prelabdiscussion (see the Supporting Information) may take up to a50 min lecture depending upon the level of the class.The prelab talk includes a discussion of the normal modes of

vibration and infrared spectroscopy and emphasizes that thevibrational frequency of a molecule in its ground electroniclevel is obtained with an IR spectrophotometer. Specialemphasis is also given to the carbonyl group and its vibrationmodes. In addition, we also stress the fact that the presentexperiment deals with vibrational frequency not only in theground electronic state of the molecule (which is obtained

directly from the IR-experiment) but also in the first electronicexcited state where the behavior of the molecule can be quitedifferent. The students record the UV−visible spectrum of anorganic ketone, which provides the excited state vibrationalfrequency for the carbonyl stretching.

■ BACKGROUND

The Structure of a Simple Organic Carbonyl Bond and ItsElectronic Transition

The carbonyl bond consists of a σ bond where the bondingelectrons are cylindrically localized between the carbon and theoxygen atom (Table 1). The π orbital spread over the sigmaplane above and below the >CO planar framework has anode and is higher in energy compared to the σ bonding orbital(see the figure for the π orbital, Table 1). The nonbondingorbital resides on the plane of the sigma framework and islocalized on the individual atoms. The correspondingantibonding π* orbital lies perpendicular to the plane of thesigma framework and has an additional node (see the figure).The σ* orbital of the carbonyl is high in energy with multiplenodes.Molecular Vibrations, the IR Spectra, and the CO Stretch

In a molecule the atoms are connected through bonds. Thus, avibration of one of the bonds is naturally dissipated to the otherones. To simplify, if it is assumed that these vibrations occur ina perfectly simple harmonic fashion, then, for a molecule withN atoms, the number of independent coordinates required todescribe the dynamically independent vibrational motions ofthe atoms is 3N − 6.3 These vibrational energies are quantized.According to the electromagnetic theory of light, light consistsof an oscillating electric and magnetic field. This oscillatingelectromagnetic field interacts with the dipole moment of themolecule (or, for the carbonyl group, the dipole moment of thebond). In the language of quantum mechanics, the probabilityof a transition from a state ψi to a state ψf is given by the square

Laboratory Experiment

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of the integral ∫ ψf A ψi dτ (known as the transition momentintegral) where A is the operator that acts on the ground stateψi to transport it to the excited state ψf. This occurs only whenthe energy difference between the ψi state and the ψf state isequal to the energy of the photon associated with theelectromagnetic wave. According to first order perturbationtheory, the operator that acts on state i is the dipole momentoperator. Thus, the larger the dipole moment of the molecule(or the concerned bond), the higher the probability of suchtransitions, and the greater the intensity of the spectral band.This, however, is not the sole criterion for the transition to beintense. The symmetry associated with the change in the dipolemoment of the particular bond is also to be considered. In thecase of the CO bond, the change in the CO bond moment ishigh compared to the other bonds in the molecule. Hence theCO stretching appears as an intense band compared to theC−C stretch or the C−H stretch.For a vibrational transition to occur from one vibrational

state to another, the energy required is small (typically ∼102−103 cm−1 in an organic molecule), whereas, for an electronictransition, the energy required is several times higher inmagnitude (typically of the order of 105 cm−1, although it isusually quoted in terms of wavelength in nm; the cm−1 unit isused here for easy comparison).

■ EXPERIMENTAL OVERVIEWThis experiment can be introduced as a part of any basicspectroscopy laboratory course, or even as an extension of atheory class. Additionally it can be taught as a follow-up of aninfrared spectroscopy lab experiment. A discussion on groundand excited state energy levels (S0 and S1) along with theassociated vibrational levels, the Franck−Condon principle, andhow it determines the shapes and the finer structures of theabsorption spectra can be included in the prelab talk beforeperforming the experiment.1 Basic shapes of n, π, and π*orbitals and the nature of n−π* and π−π* transitions can bediscussed as well.4 Additionally, a computation software (e.g.,

MolCalc) may be introduced by the teacher as a tool to helpthe students to visualize the orbitals and calculate their energy.The vibrational bands of a molecule can be obtained directly

from the IR spectrum where the bands correspond to thenormal modes of vibration of the molecule in its electronicground state S0. The red arrow in Figure 1 refers to the

transition between the 0 and 1 vibrational states of the groundelectronic state. This corresponds to an energy difference ofapproximately 1745 cm−1. The students can easily identify theintense CO stretch in the IR spectrum recorded for thisexperiment and note its exact wavenumber.In the UV−visible spectroscopy, the molecule initially resides

in the lowest electronic state (S0) and in the lowest vibrationalstate, v = 0 (assuming the low temperature approximation).Upon absorption of a photon, it undergoes a transition to thenext electronic level (S1). The associated vibrational states forthe S1 are shown in Figure 1. For simplicity, the shifts in theequilibrium position for the vibrations have not been shown inFigure 1. These aspects are lucidly described in spectroscopytext books1 and also in the Supporting Information.The energy difference between the two consecutive peaks in

the spectrum corresponds to the difference in energy of the twovibrational energy levels of the S1 state as shown in Figure 1.Thus, the values for the CO stretching in the S1 state,ν(CO)*, can be determined in class and compared with theone in the S0 electronic state, i.e., ν(CO), obtained directlyfrom the IR spectrum.In addition to the experiment, students are encouraged to use

a computational software to visualize the molecular orbitals. Ahighly recommended web based application, the MoleculeCalculator (MolCalc), was used to visualize the molecularorbitals and determine the orbital energies.5 This applicationuses an ab initio Hartree−Fock method with a STO-3G basisset to perform its calculations. It comes with a word of cautionthat the numerical values of the stretching frequency obtained

Table 1. Molecular Orbitals of a Carbonyl Compound andthe Electronic Transitions

Figure 1. Schematic representation of the vibrational levels in the S0and S1 electronic states. Note that the red arrow corresponds to theground state stretching frequency ν(CO) obtained from the IRspectrum directly. The numbers with the “prime” sign refer to thevibrational levels of the S1 state. The double headed arrows betweenthe vibrational levels in the S1 state correspond to the vibrationalstretching frequencies ν(CO)* in the first electronic excited state andare determined in this experiment using UV spectroscopy.

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from the application do not match with the experimentallydetermined values from the IR experiments although theyfollow the same order of sequence. The energy values providean idea of the energy levels relative to each other and the natureof the orbitals. Additionally, with the same application, thevibrational modes of a molecule in its ground state can bevisualized.

■ MATERIALSSpectro-grade cyclopentanone and cyclohexane are required forthis experiment. The experiment needs to be performed withmeticulous cleanliness. Cuvettes used in the experimentsshould be clean and dried.

■ EXPERIMENTAL METHODSThe experimental methods are provided in detail in theSupporting Information.The infrared spectrum of cyclopentanone is recorded with a

liquid film smeared on a salt plate. The most intense band inthe IR spectrum that corresponds to the CO stretching ofcyclopentanone is noted.To record the UV−visible spectrum, a few milligrams of

cyclopentanone6 is dissolved in cyclohexane such that theabsorption bands are not too weak (<0.01) or too strong(>2.5). The spectrum should be recorded with short intervals(1 nm or less) of the data points such that the finer vibrationalstructures become noticeable. Note that some of the finer linesof spectra are not clearly seen because of spectral broadening bysolvent interactions. However, these finer bands are clearlyvisible at the higher wavelengths of the UV−visible spectrum.The λmax values of each of the bands are recorded in a tabular

form (Table 2). If the absorption bands are too intense thenthe sample should be diluted to an optimum intensity.

■ HAZARDSCyclohexane is a highly flammable solvent. It is harmful andmay cause lung damage if swallowed. It is irritating to skin, andthe vapors may cause drowsiness and dizziness. Cyclo-pentanone is flammable and is irritating to eyes and skin.Refer to the MSDS for specific hazard information. Gloves, labcoat, and protective eyewear should be worn for thisexperiment.

■ RESULTS AND DISCUSSIONThe experiment has been run several times with the third yearundergraduate students at IISER Kolkata, each time with a

batch of 24 students working in pairs. The introductory lecturewas followed by a hands-on demo of the computationalsoftware. The students finished the experiments and generatedthe MO in class. The UV−visible spectrum of cyclopentanonein cyclohexane is shown in Figure 2. In the spectrum, the finer

bands at lower wavelengths (230−280 nm) lack sharp peaks,and assigning their exact wavelength is difficult. The values ofthe wavelengths corresponding to the prominent peaks of thefiner bands are recorded in nm units which are converted to thecorresponding wavenumbers (cm−1). The IR spectrum ofcyclopentanone recorded with a liquid film of the neat sampleon an IR plate provides the carbonyl stretch in the ground statewith an intense band in the spectra at 1745 cm−1.Calculation

The wavelengths are converted to the corresponding wave-numbers. Sets of two adjacent peaks (in cm−1) are taken, andtheir differences are calculated. These numbers corresponds tothe vibrational levels of the CO bands in the S1 (excited)electronic state (Figure 1). Typical results of the experiment areshown in Table 2 as an example.The average vibrational frequency in the first electronic

excited state thus obtained with the data provided in Table 2 at30 °C is 1197 ± 24 cm−1. This value is in agreement with theones reported for the carbonyl compounds in the literature1,7,8

and is ca. 550 cm−1 less than the ground state stretchingfrequency of same unit.As a result of the electronic transition, there is a pronounced

change in the polarity, the bond order, and the length of thecarbon−oxygen bond (see Supporting Information). Theweaker C−O bond in the excited state causes a decrease inthe carbonyl stretching frequency. The value of the COstretch in the excited state obtained from this experimentsupports it.Suggested Add-On to the Experiment: MO CalculationUsing MolCalc

Using the Web-based application Molecule Calculator(MolCalc), the molecular orbitals of cyclopentanone can bevisualized easily and the energies of the orbital can bedetermined.5 For the sake of simplicity, a beginner can startwith the formaldehyde molecule, where the number of orbitalsis lower. Please refer to the MolCalc Web site for tutorials.5

The shapes of the orbitals using a simple model carbonylcompound have been presented in Figure 1. Note that thenumbers quoted in Table 3 are obtained from the application

Table 2. Wavelengths of the Prominent Peaks and TheirCorresponding Wavenumbersa

Wavelength, nm Wavenumber,b cm−1 Difference,c cm−1

323.2 30941311.2 32134 1193299.8 33356 1222289.6 34530 1175

aThe data was acquired in an undergraduate laboratory class. The dataprovided here was obtain by an individual in the class (S.R.). Althoughthis is the data from one of the students, the results of the entire classagreed well among each other. bWavelength (nm) to wavenumber(cm−1) conversion: x nm = 107/x cm−1. c“Difference” refers to thedifference in wavenumbers between two consecutive peaks. Refer toFigure 1.

Figure 2. UV−visible spectrum of cyclopentanone in cyclohexane.The transitions (peaks) correspond to the gap of the successivevibrational levels in the excited state for the CO bond. (Also seeFigure 1.)

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MolCalc and they are close to the experimentally obtainedvalues.

To demonstrate the difference in behavior of a molecule inits electronic ground state and excited state, this experimentwas followed up by an experiment for the determination ofexcited state pKa of an organic phenol.9

■ ASSOCIATED CONTENT*S Supporting Information

Instructor notes and student handout. This material is availablevia the Internet at http://pubs.acs.org.

■ AUTHOR INFORMATIONCorresponding Author

*E-mail: sb1@iiserkol.ac.in.Notes

The authors declare no competing financial interest.

■ ACKNOWLEDGMENTSThe author thanks Prof. Sanjib Bagchi for his suggestionsduring the design of the experiment. S.R. is an undergraduatestudent funded by a KVPY fellowship from the Department ofScience and Technology, Government of India. The authorsacknowledge Ankan Bag, Suman Pal, Joydev Hatai, V. SivaRama Krishna, and Mousumi Samanta for their help with theexperiments and IISER Kolkata for support.

■ REFERENCES(1) Principles of Molecular Photochemistry: An Introduction, 1st ed.;Turro, N. J., Ramamurthy, V., Sciano, J. C., Eds.; University ScienceBooks: 2009.(2) (a) Steinfeld, J. I.; Houston, P. L. In Laser and coherencespectroscopy; Steinfeld, J. I., Ed.; Plenum Press: New York, 1978.(b) Ito, M. Electronic spectra in a supersonic jet as a means of solvingvibrational problems. In Vibrational Spectra and Structure; Durig, J. R.,Ed.; Elsevier: Amsterdam, 1986. (c) Ito, M.; Ebata, T.; Mikarni, N.Laser Spectroscopy of Large Polyatomic Molecules in Supersonic Jets.Annu. Rev. Phys. Chem. 1988, 39, 123.

(3) Wilson, E. B.; Decius, J. C., Cross, P. C. Molecular Vibrations;McGraw-Hill: New York, 1955.(4) Swenton, J. S. Photochemistry of Organic Compounds IICarbonyl Compounds. J. Chem. Educ. 1969, 46, 217−226.(5) Jensen, J. H.; Kromann, J. C. The Molecule Calculator: A WebApplication for Fast Quantum Mechanics-Based Estimation ofMolecular Properties. J. Chem. Educ. 2013, 90, 1093−1095. MoleculeCalculator (MolCalc) Website: http://molcalc.org/ (accessed Aug2014).(6) The reason we have chosen cyclopentanone and not acetone orbenzophenone as the source for a CO group was simple. At roomtemperature, acetone does not give rise to finer structures in theelectronic spectra because of the dissipation of energy by the C−CH3rotation mode. Benzophenone, on the other hand, has strong π−π*transitions of the phenyl rings where often the carbonyl π−π*transition gets hidden under. Therefore, cyclopentanone was used foreasier assignment of the peaks.(7) Brand, J. C. D. The electronic spectrum of formaldehyde. J. Chem.Soc. 1956, 858−872.(8) Moule, D. C.; Walsh, A. D. Ultraviolet spectra and excited statesof formaldehyde. Chem. Rev. 1975, 75, 67−84.(9) (a) Halpern, A. M.; Reeves, J. H. Experimental Physical Chemistry;Scott, Foresman and Company: Boston, 1988. (b) Boyer, R.; Deckey,G.; Marzzacco, C.; Mulvaney, M.; Schwab, C.; Halpern, A. M. Thephotophysical properties of 2-naphthol: A physical chemistry experi-ment. J. Chem. Educ. 1985, 62, 630.

Table 3. MO of the Carbonyl Compound and Their RelativeEnergies Calculated Using MolCalca

a†Note that the wavelengths (nm) cannot be negative. The negativesign appears because of the conversion of the difference of the energyto wavelength. ‡Note for the advanced learners: The transitionenergies are not simply equal to these differences since they do nottake into account changes in coulomb and exchange interactionsaccompanying the transitions. However, the differences in the orbitalenergies usually give the correct order of the transitions.

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