Determination and Comparison of Carbonyl Stretching Frequency of a

Laboratory Experiment pubs.acs.org/jchemeduc

Determination and Comparison of Carbonyl Stretching Frequency of a Ketone in Its Ground State and the First Electronic Excited State Subhajit Bandyopadhyay* and Saswata Roy Department of Chemical Sciences, Indian Institute of Science Education and Research (IISER) Kolkata, Mohanpur, West Bengal 741252, India S Supporting Information *

ABSTRACT: This paper describes an inexpensive experiment to determine the carbonyl stretching frequency of an organic keto compound in its ground state and first electronic excited state. The experiment is simple to execute, clarifies some of the fundamental concepts of spectroscopy, and is appropriate for a basic spectroscopy laboratory course. The experiment is complemented 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

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directly from the IR-experiment) but also in the first electronic excited state where the behavior of the molecule can be quite different. The students record the UV−visible spectrum of an organic ketone, which provides the excited state vibrational frequency for the carbonyl stretching.

INTRODUCTION Designing an experiment for an undergraduate physical chemistry laboratory that integrates the concept of electronic energy levels and the vibrational energy levels, coupled to the estimation of stretching frequency in the first electronic excited state, and correlating that with the one in the ground state might seem like a complicated task. Generally, in spectroscopy lab courses, the students carry out discrete experiments in diverse areas of spectroscopy. The “big picture” that integrates the electronic and vibrational spectroscopy, their nature, and the energy scale at which they operate are often missed by the students. Inspired by a textbook of photochemistry,1 a simple undergraduate laboratory experiment has been designed that helps students to link the concepts of IR and UV spectroscopy. Normally, the vibrational stretching frequency of a compound in its first electronic excited state is determined precisely by UV−IR double resonance spectroscopy of jet cooled molecules.2 Although the experiment described here lacks such precision, it does provide a quick and easy means to have a rough estimation of the vibrational frequency of a CO bond in its first electronic excited state. The experiment can be performed easily with readily available resources in an undergraduate laboratory. At the end of the experiment, the students had better understanding of some of the fundamental elements of molecular spectroscopy. The estimated time taken to perform this experiment is 40 min, and 10−20 min should be allotted for the calculations. A group of two students would be ideal for conducting this experiment in the class. The prelab discussion (see the Supporting Information) may take up to a 50 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 the vibrational frequency of a molecule in its ground electronic level is obtained with an IR spectrophotometer. Special emphasis is also given to the carbonyl group and its vibration modes. In addition, we also stress the fact that the present experiment deals with vibrational frequency not only in the ground electronic state of the molecule (which is obtained © XXXX American Chemical Society and Division of Chemical Education, Inc.

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BACKGROUND

The Structure of a Simple Organic Carbonyl Bond and Its Electronic Transition

The carbonyl bond consists of a σ bond where the bonding electrons are cylindrically localized between the carbon and the oxygen atom (Table 1). The π orbital spread over the sigma plane above and below the >CO planar framework has a node and is higher in energy compared to the σ bonding orbital (see the figure for the π orbital, Table 1). The nonbonding orbital resides on the plane of the sigma framework and is localized on the individual atoms. The corresponding antibonding π* orbital lies perpendicular to the plane of the sigma framework and has an additional node (see the figure). The σ* orbital of the carbonyl is high in energy with multiple nodes. Molecular Vibrations, the IR Spectra, and the CO Stretch

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

A

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Table 1. Molecular Orbitals of a Carbonyl Compound and the Electronic Transitions

MolCalc) may be introduced by the teacher as a tool to help the 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 the normal modes of vibration of the molecule in its electronic ground state S0. The red arrow in Figure 1 refers to the

of the integral ∫ ψf  ψi dτ (known as the transition moment integral) where  is the operator that acts on the ground state ψi to transport it to the excited state ψf. This occurs only when the energy difference between the ψi state and the ψf state is equal to the energy of the photon associated with the electromagnetic wave. According to first order perturbation theory, the operator that acts on state i is the dipole moment operator. Thus, the larger the dipole moment of the molecule (or the concerned bond), the higher the probability of such transitions, and the greater the intensity of the spectral band. This, however, is not the sole criterion for the transition to be intense. The symmetry associated with the change in the dipole moment of the particular bond is also to be considered. In the case of the CO bond, the change in the CO bond moment is high compared to the other bonds in the molecule. Hence the CO stretching appears as an intense band compared to the C−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 electronic transition, the energy required is several times higher in magnitude (typically of the order of 105 cm−1, although it is usually quoted in terms of wavelength in nm; the cm−1 unit is used here for easy comparison).

Figure 1. Schematic representation of the vibrational levels in the S0 and S1 electronic states. Note that the red arrow corresponds to the ground state stretching frequency ν(CO) obtained from the IR spectrum directly. The numbers with the “prime” sign refer to the vibrational levels of the S1 state. The double headed arrows between the vibrational levels in the S1 state correspond to the vibrational stretching frequencies ν(CO)* in the first electronic excited state and are determined in this experiment using UV spectroscopy.

transition between the 0 and 1 vibrational states of the ground electronic state. This corresponds to an energy difference of approximately 1745 cm−1. The students can easily identify the intense CO stretch in the IR spectrum recorded for this experiment and note its exact wavenumber. In the UV−visible spectroscopy, the molecule initially resides in the lowest electronic state (S0) and in the lowest vibrational state, v = 0 (assuming the low temperature approximation). Upon absorption of a photon, it undergoes a transition to the next electronic level (S1). The associated vibrational states for the S1 are shown in Figure 1. For simplicity, the shifts in the equilibrium position for the vibrations have not been shown in Figure 1. These aspects are lucidly described in spectroscopy text 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 two vibrational energy levels of the S1 state as shown in Figure 1. Thus, the values for the CO stretching in the S1 state, ν(CO)*, can be determined in class and compared with the one in the S0 electronic state, i.e., ν(CO), obtained directly from the IR spectrum. In addition to the experiment, students are encouraged to use a computational software to visualize the molecular orbitals. A highly recommended web based application, the Molecule Calculator (MolCalc), was used to visualize the molecular orbitals and determine the orbital energies.5 This application uses an ab initio Hartree−Fock method with a STO-3G basis set to perform its calculations. It comes with a word of caution that the numerical values of the stretching frequency obtained

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EXPERIMENTAL OVERVIEW This experiment can be introduced as a part of any basic spectroscopy laboratory course, or even as an extension of a theory class. Additionally it can be taught as a follow-up of an infrared spectroscopy lab experiment. A discussion on ground and excited state energy levels (S0 and S1) along with the associated vibrational levels, the Franck−Condon principle, and how it determines the shapes and the finer structures of the absorption spectra can be included in the prelab talk before performing the experiment.1 Basic shapes of n, π, and π* orbitals and the nature of n−π* and π−π* transitions can be discussed as well.4 Additionally, a computation software (e.g., B

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batch of 24 students working in pairs. The introductory lecture was followed by a hands-on demo of the computational software. The students finished the experiments and generated the MO in class. The UV−visible spectrum of cyclopentanone in cyclohexane is shown in Figure 2. In the spectrum, the finer

from the application do not match with the experimentally determined values from the IR experiments although they follow the same order of sequence. The energy values provide an idea of the energy levels relative to each other and the nature of the orbitals. Additionally, with the same application, the vibrational modes of a molecule in its ground state can be visualized.

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MATERIALS Spectro-grade cyclopentanone and cyclohexane are required for this experiment. The experiment needs to be performed with meticulous cleanliness. Cuvettes used in the experiments should be clean and dried.

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EXPERIMENTAL METHODS The experimental methods are provided in detail in the Supporting Information. The infrared spectrum of cyclopentanone is recorded with a liquid film smeared on a salt plate. The most intense band in the IR spectrum that corresponds to the CO stretching of cyclopentanone is noted. To record the UV−visible spectrum, a few milligrams of cyclopentanone6 is dissolved in cyclohexane such that the absorption bands are not too weak (2.5). The spectrum should be recorded with short intervals (1 nm or less) of the data points such that the finer vibrational structures become noticeable. Note that some of the finer lines of spectra are not clearly seen because of spectral broadening by solvent interactions. However, these finer bands are clearly visible 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 then the sample should be diluted to an optimum intensity.

Figure 2. UV−visible spectrum of cyclopentanone in cyclohexane. The transitions (peaks) correspond to the gap of the successive vibrational levels in the excited state for the CO bond. (Also see Figure 1.)

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

The wavelengths are converted to the corresponding wavenumbers. Sets of two adjacent peaks (in cm−1) are taken, and their differences are calculated. These numbers corresponds to the vibrational levels of the CO bands in the S1 (excited) electronic state (Figure 1). Typical results of the experiment are shown 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 at 30 °C is 1197 ± 24 cm−1. This value is in agreement with the ones reported for the carbonyl compounds in the literature1,7,8 and is ca. 550 cm−1 less than the ground state stretching frequency 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 the carbon−oxygen bond (see Supporting Information). The weaker C−O bond in the excited state causes a decrease in the carbonyl stretching frequency. The value of the CO stretch in the excited state obtained from this experiment supports it.

Table 2. Wavelengths of the Prominent Peaks and Their Corresponding Wavenumbersa Wavelength, nm

Wavenumber,b cm−1

Difference,c cm−1

323.2 311.2 299.8 289.6

30941 32134 33356 34530

1193 1222 1175

a

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

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HAZARDS Cyclohexane is a highly flammable solvent. It is harmful and may cause lung damage if swallowed. It is irritating to skin, and the vapors may cause drowsiness and dizziness. Cyclopentanone is flammable and is irritating to eyes and skin. Refer to the MSDS for specific hazard information. Gloves, lab coat, and protective eyewear should be worn for this experiment.

Suggested Add-On to the Experiment: MO Calculation Using MolCalc

Using the Web-based application Molecule Calculator (MolCalc), the molecular orbitals of cyclopentanone can be visualized easily and the energies of the orbital can be determined.5 For the sake of simplicity, a beginner can start with the formaldehyde molecule, where the number of orbitals is lower. Please refer to the MolCalc Web site for tutorials.5 The shapes of the orbitals using a simple model carbonyl compound have been presented in Figure 1. Note that the numbers quoted in Table 3 are obtained from the application

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RESULTS AND DISCUSSION The experiment has been run several times with the third year undergraduate students at IISER Kolkata, each time with a C

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(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 II Carbonyl Compounds. J. Chem. Educ. 1969, 46, 217−226. (5) Jensen, J. H.; Kromann, J. C. The Molecule Calculator: A Web Application for Fast Quantum Mechanics-Based Estimation of Molecular Properties. J. Chem. Educ. 2013, 90, 1093−1095. Molecule Calculator (MolCalc) Website: http://molcalc.org/ (accessed Aug 2014). (6) The reason we have chosen cyclopentanone and not acetone or benzophenone as the source for a CO group was simple. At room temperature, acetone does not give rise to finer structures in the electronic spectra because of the dissipation of energy by the C−CH3 rotation 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 for easier 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 states of 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. The photophysical properties of 2-naphthol: A physical chemistry experiment. J. Chem. Educ. 1985, 62, 630.

MolCalc and they are close to the experimentally obtained values. Table 3. MO of the Carbonyl Compound and Their Relative Energies Calculated Using MolCalca

a†

Note that the wavelengths (nm) cannot be negative. The negative sign appears because of the conversion of the difference of the energy to wavelength. ‡Note for the advanced learners: The transition energies are not simply equal to these differences since they do not take into account changes in coulomb and exchange interactions accompanying the transitions. However, the differences in the orbital energies usually give the correct order of the transitions.

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

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ASSOCIATED CONTENT

S Supporting Information *

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

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AUTHOR INFORMATION

Corresponding Author

*E-mail: [email protected]. Notes

The authors declare no competing financial interest.

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ACKNOWLEDGMENTS The author thanks Prof. Sanjib Bagchi for his suggestions during the design of the experiment. S.R. is an undergraduate student funded by a KVPY fellowship from the Department of Science and Technology, Government of India. The authors acknowledge Ankan Bag, Suman Pal, Joydev Hatai, V. Siva Rama Krishna, and Mousumi Samanta for their help with the experiments and IISER Kolkata for support.

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REFERENCES

(1) Principles of Molecular Photochemistry: An Introduction, 1st ed.; Turro, N. J., Ramamurthy, V., Sciano, J. C., Eds.; University Science Books: 2009. (2) (a) Steinfeld, J. I.; Houston, P. L. In Laser and coherence spectroscopy; Steinfeld, J. I., Ed.; Plenum Press: New York, 1978. (b) Ito, M. Electronic spectra in a supersonic jet as a means of solving vibrational 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. D

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