The Solvation of Urea, Tetramethlyurea, TMU, in some Organic Solvents

G. Rezaei Behbehani, M. Hamedi, F. Hoseinpour Rajabi.

Department of Chemistry,
Imam Khomeini International University,
Qazvin, Iran

Abstract

Introduction

Infrared spectra of methyl acetate, [1] acetone,[2] and several amides [3-4] including N, N-dimethylacetamide, in aqueous DMSO mixtures, have been reported previously. In mixed solvents, these spectra commonly show several carbonyl stretching bands, assigned to differently hydrogen bonded species.

The spectroscopic studies of –C=O groups in a range of amides [3-4] in mixed aqueous and non-aqueous solvent systems have shown clearly the existence of di-, mono- and non-hydrogen bonded examples of protic solvent (such as water or methanol) interaction with the –C=O group. The frequency of this band depends on the nature and strength of the intermolecular interactions (such as a hydrogen bond). Thus, when the aprotic solvent is dissolved as a dilute solute in a second solvent, the solvent induced frequency shift, SIFS, may be used to characterize the intermolecular interactions.

In the following paper, one single band was found for –C=O stretch for all of the solvents considered, but the position of the band shifted to higher values in the more aprotic solvents. These measurements provide information about the solvation of the oxygen atom of –C=O group and also other surfaces of the solute molecules as well as the whole solute molecules.

Experimental

Spectra were collected using a Bomem Research Series FTIR spectrometer with a resolution of 0.5 cm^-1. All experiments were carried out at 25°C. The solvents used were HPLC grade, at least 99.9% pure with a residual water level less than 0.005%. Solutions were made up with a constant molar ratio of 1 mole of solution to 12 moles of solvent. They were handled in a dry box under a nitrogen atmosphere to prevent contamination by water. The spectra were analyzed by subtracting the spectrum of the pure solvent from that of the solution. The data were analyzed using Quattro Pro software.

Results and Discussion

The –C=O stretch band in urea and trimethyl urea were observed in 13 pure solvents namely water, methanol, ethanol, 1-propanol, 2-propanol, t-butanol, DMSO, MeCN, Nitromethane, Diethylether, CCl₄, Benzene and Hexane.

Studies on solvation in these systems have shown that the acidity and basicity of the solvent have significantly effects the SIFS. The Gutmann donor number, DN [7], and the Gutmann acceptor number, AN, were chosen for the solvent basicity and the solvent acidity [8] respectively. Detailed arguments for the choice of these scales were given previously [5, 6].

Equation 1 can be written for the p_DN solvents AN and DN as follows:

(1)

The overall preferential solvation parameter, p, for the whole solute molecule can be defined through:

(2)

This parameter has been determined from previously calorimetric studies [9-11]. The partial preferential solvation factors, p_AN, ¾ the specific preferential solvation of the oxygen atom of the –C=O group for the solvents considered and, p_DN, ¾ an indicator of preferential solvation to other sites on the solute molecule can be defined as follows:

(3)

(4)

If we divide the values calculated from Equations 2-4 by that of a reference solvent (such as water in the cases considered), we can deduce the relative preferential solvation factors in binary solvent systems as follows:

(5)

(6)

Where α₁(R) and α₂(R) are the coefficients describing the response of the band to the values of AN_R and DN_R of the reference solvent respectively. The wave number values have been measured for the –C=O group of urea and trimethyl urea in the solvents selected. The equation for the least squares fit is

(7)

with a standard deviation of 8.17 cm^-1 and correlation coefficient of 0.956. When the electron density associated with the –C=O group interacts with a protic solvent, the frequency of the –C=O stretching band is decreased. The fit to the experimental data is improved somewhat by adding the solvent DN as a second variable. There is a clear correlation between the frequency of this band and the solvent’s AN as can be seen in Table 4. The description of the solvent effect is then

(8)

resulting in a standard deviation of 0.58 cm^-1 and correlation coefficient of 0.9998. On the basis of the partial regression coefficients solvent acidity accounts for 62% of the explained variation in ω₁ and solvent basicity for 38%. In other words, the probability that the solvents interact with the oxygen atom of the –C=O group is 62% and the probability that the solvent interaction occurs with other sites on the urea molecule (i.e. – NH2 groups) is 38%. The variance indicates that interaction of the solvent with the electronegative –C=O group in urea results in a reduction in the bond frequency. The relative overall preferential solvation factor for the whole urea molecule, pR, and the relative partial preferential solvation factors, p_AN, p_DN, using water as a reference solvent for urea in the considered solvents is listed in Table 2. It should be stated that the value of unity for the preferential solvation factors of the solute in the reference solvent (water in these cases) does not mean that random solvation occurs. p<1 or p>1 indicate preferential solvation of the solute by the reference solvent or by the co-solvent respectively; p=1 (with the exception of the reference solvent) indicates random solvation. We can presumably relate the preferential solvation parameters for the solutes in pure solvents to those in mixed binary solvent systems. For example the preferential solvation factors (Table 2) indicate that in the aqueous MeCN solvent system, both the –C=O and –N– H protons of the urea are preferentially solvated by water (both p_AN and p_DN are less than one indicating preferential solvation by water), but that in the aqueous DMSO solvent system, preferential solvation of the –C=O is by water (as p_AN=0.352) whilst that of the –N– H protons of urea is by DMSO (as p_DN=1.656) . In the aqueous MeCN solvent system, it is possible to predict that urea residues in the aqueous structure as both sides of the urea molecule are preferentially solvated by water.

In the case urea in aqueous DMSO, one presumes that urea is accommodated at the interface between water and DMSO aggregating with one side of the urea (–C=O) preferentially hydrated while the other surface of the molecule (–N– H protons) is preferentially solvated by DMSO. Therefore, solvent-solvent bond disruption by urea in aqueous DMSO is more than that of in aqueous MeCN. The overall preferential solvation factors for urea (p_R on Table 2), indicate the preferential solvation of the whole urea molecule by water.

The wave numbers for the –C=O stretching band of TMU in the solvents considered have been measured and a correlation with the solvent’s acceptor number may be deduced as follows:

(9)

with a standard deviation of 0.399 cm^-1 and a correlation coefficient of 0.9987. The fit is considerably improved by adding DN to the description of the solvent effect. The result is then

(10)

generating then a standard deviation of 0.183 cm^-1 and an excellent correlation coefficient of 0.9998. No further improvement is found when the solvent polarity, Y, is added as a third variable. The role of each individual parameters is computed by calculating the partial regression coefficients. Thus, the solvent acidity accounts for 94% of the variation in ω₂ and solvent basicity for 6%. The agreement between the experimental values of ω₂ and the estimated values by equation 10 is excellent and are tabulated in Table 4. Due to the steric effect of the –CH3 groups, all interactions are related to the solvent oxygen of –C=O group and interaction with other sites on the TMU molecule (i.e. –CH3 groups) are almost negligible. Significant contribution to AN (-0.784) versus a small one from DN (+0.047) in equation 10 is excellent support for this interpretation.

In those solutes considered, the interaction of the –C=O group with a protic solvent leads to a decrease in the frequency of the –C=O stretching mode. This is attributed to the formation of mono- and di-hydrogen-bonded species involving the protic solvent with the –C=O group of the solutes. It is possible to relate the solvation parameters of the TMU molecule in pure solvents to those of for this molecule in mixed binary solvent systems. For example, the preferential solvation factors in Table 3 indicate that in the all the aqueous solvent systems both the –C=O and –N– H protons of the TMU are preferentially solvated by water (both p_AN(R) and p_DN(R) p_R for all solvents are less than one indicating preferential solvation by water). The overall preferential solvation factors for TMU (p_R on Table 3), indicate the preferential solvation of the whole TMU molecule by water.

References

1. K. B. Patel, G. Eaton, and M. C. R. Symon, J. Chem. Soc. Faraday Trans. I, (1985), 85, 2775.
2. M. C. R. Symon and G. Eaton, J. Chem. Soc. Faraday Trans. I, (1985), 81, 1963.
3. G. Eaton, and M. C. R. Symon, J. Chem. Soc. Faraday Trans. I, (1988), 84, 3429.
4. G. Eaton, and M. C. R. Symon, and P. P. Rastogi, J. Chem. Soc. Faraday Trans. I, 85, (1989), 3257.
5. W. R. Fawcett, in Theoretical and Computational Chemistry: Quantitative Treatments of Solute/Solvent Interactions, P. Politzer and J. S. Murray, editors, Elsevier, Amsterdam, Volume 1, p. 183.
6. W. R. Fawcett, J. Phys. Chem., 97 (1993) 9540.
7. V. Gutmann, Coord. Chem. Rev., 19 (1976) 225.
8. U. Mayer, V. Gutmann, and W. Gerger, Monatsch. Chem., (1975), 106 , 1235.
9. D. Feakins, C. O’Duinn and W. E. Waghorne, J. Solution Chem., (1987), 16, 907.
10. A. Costigan, D. Feakins, I. McStravick, C. O’Duinn, J. Ryan and W. E. Waghorne, J. Chem. Soc., Faraday Trans. (1991), 87, 2443.
11. D. Feakins, P. Hogan, C. O’Duinn, and W. E. Waghorne, J. Chem. Soc., Faraday Trans. (1992), 88, 423.

Received 14th December 2001, received in revised format 18th January 2002, accepted 19th January 2002.

REF: G. Rezaei Behbehani, M. Hamedi, F. Hoseinpour Rajabi. Int. J. Vib. Spect., [www.irdg.org/ijvs] 5, 6, 5 (2001)