UV-Visible spectroscopy and density functional study of solvent effect on halogen bonded charge-transfer complex of 2-Chloropyridine and iodine monochloride

4.1.1 Spectral properties of the acceptor, donor and the complex formed in solution

2-Chloropyridine and ICl both are soluble in all the five solvents used in this study. Solution of 2-Chloropyridine in all the solvents is colourless while that of ICl is light brownish. Gradual addition of 2-Chloropyridine solution to the ICl solution results in lightening of brown colour of ICl up to different extents indicating the formation of a new compound (Refat et al., 2013). Pure ICl absorbs at 453.2 nm and 2-Chloropyridine absorbs at 256.2 nm in chloroform. When these two are mixed a new band appears at 324 nm which is neither the characteristic of ICl or of 2-Chloropyridine. Appearance of this band confirms the formation of the charge transfer complex between ICl and 2-Chloropyridine (Fig. 1). Similar observations are obtained for the other four solvents (Fig. 2). From the UV-vis spectra the stoichiometry of the complex formed is found to be 1:1 using the modified Benesi-Hildebrand plot which gives straight lines for all the five solvents considered (Fig. 3).

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Figure 1

UV-vis spectra of 2-Chloropyridine and ICl in Chloroform at 25 °C for 1 cm cell. The concentrations of 2-Chloropyridine (M) are (1) zero; (2) 0.1054; (3) 0.3162; (4) 0.527; (5) 0.7378; (6) 0.9486.

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Figure 2

UV-vis spectra of 2-Chloropyridine-ICl charge transfer complex in (a) Hexane, (b) CCl₄, (c) CH₂Cl₂ and (d) Cyclohexane at 25 °C in 1 cm cell. The concentration of ICl is kept constant and that of 2-Chloropyridine is varied. The concentrations of 2-Chloropyridine (M) are (a) Hexane: (1) zero; (2) 0.06999; (3) 0.20997; (4) 0.34995; (5) 0.48993; (6) 0.62991. (b) CCl_4: (1) zero; (2) 0.0351; (3) 0.0763; (4) 0.1053; (5) 0.1755; (6) 0.3159. (c) CH₂Cl₂: (1) zero; (2) 0.1054; (3) 0.3162; (4) 0.527; (5) 0.7378; (6) 0.9486. (d) Cyclohexane: (1) zero; (2) 0.06999; (3) 0.20997; (4) 0.34995; (5) 0.48993; (6) 0.62991.

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Figure 3

Modified Benesi-Hildebrand plot for 2-Chloropyridine-ICl complex in CHCl₃.

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Table 1 lists the physical data such as CT absorption band wavelength (λ_max), Extinction coefficient (ε_max)_, Formation constant (K_C), Energy of the CT band (ΔE_CT), Ionisation Potential of the donor in the complex (I_p), Dissociation Energy (W), Resonance Energy of the complex in the ground state (R_N), and the Gibbs free energy change (ΔG) at the experimental temperature. These values are calculated with the help of established equations (Razzaq et al., 2008). The Ionisation potential values of the donor match nicely with the previously calculated values by electron impact method (Basila and Clancy, 1963).

Table 1 clearly shows the dependence of the λ_max value on the dielectric constant value of the experimental solvents. Similarly all the parameters linked with λ_max vary accordingly. It is already reported that increasing solvent polarity affects the Charge Transfer Energy by solvating the energy levels up to different extents. Lowering of the LUMO in the energy axis or stabilizing it more as compared to the HOMO, may lead to an increase in the charge transfer transition energy (ΔE_CT). It may occur as a result of more polar nature of the LUMO which leads to more solvation of the LUMO and as a result it gets stabilized to a greater extent. Moreover a blue shift of λ_max indicates n-σ^∗ transition. It establishes the formation of the 2-Chloropyridine-ICl complex through the lone pair on N atom as is found in case of Iodine forming compounds with lone pair donors (Kamlet et al., 1979; Salman et al., 2004).

4.1.2 Quantitative estimation of effect of solvent parameters upon CT band by MLRT

The solvent parameters for the solvents under study are in Table 2. Solvatochromic study of CT band is done in a large scale in varieties of compounds such as drugs, dyes, flavones (Catalan, 2009; Aggarwal and Khurana, 2014; Alimmari et al., 2015; Jovie et al., 2014; Dawoud et al., 2014). MLRT is a recent approach to analyse the dependence of any spectroscopic property (λ_max, Δ $\overline{\nu }$ etc.) of a compound on various solvent parameters (Hmuda et al., 2013; Alimmari et al., 2015; Zakerhamidi et al., 2012). In this study, MLRT is applied using Excel 2007 at the 95% confidence level. We have examined the dependence of λ_max on the solvent parameters under study. Among the solvent parameters we have selected ε, n, π^∗value, Hansen parameters and Catalan parameters. ε and n determine solute-solvent interaction considering bulk environment whereas the rest three indicates solute-solvent interaction at the molecular level. The regression plot of λ_max (exp.) against λ_max (calc.) (Fig. 4) shows that Hansen parameters have excellent correlation with the λ_max value with R² value of 1. The result of regression analysis is shown in Table 3.

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Figure 4

Calculated λ_max Vs experimental λ_max plots w. r. t. (a) ε and n, (b) ε, n, π^∗, (c) Hansen solubility parameters, (d) Catalan parameters.

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Table 3 shows that Solvatochromic shift is largely dependent on solute–solvent interaction at the molecular level. Similar Regression analysis with the Formation Constant value is also done. Although formation constant value of a CT complex is reported to show variation with the changing polarity no specific change is reported till now (Refat et al., 2010). Table 1 shows irregular variation in the values of K_c with polarity. Thus we tried to check whether K_c is a multi parameter dependent characteristic rather than being dependent solely upon the solvent polarity or dielectric constant value. In case of Formation constant too Fig. 5 shows that it is dependent on Hansen parameters with a high R² value of 0.993 and it is completely independent of bulk parameters. Although we have performed MLRT considering a linear relation to exist between the Formation Constant of the CT complex and the solvent parameters and we are obtaining quite satisfactory result for the Hansen parameters, still we cannot put forward a conclusive remark on that until and unless we perform similar tests with other CT complexes in different media. Since it is beyond the scope of this paper to include a large number of such complexes, we propose to examine our approach in our future works.

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Figure 5

Calculated K_C Vs experimental K_C plots w. r. t. (a) ε and n, (b) ε, n, π^∗, Hansen solubility parameters, (d) Catalan parameters.

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4.1.3 Validity of Buckingham and Lippert Mataga equation

Buckingham equation relates Spectroscopic Property (Δ $\overline{\nu }$ in our case) with the function of dielectric constant f(ε) and the function of refractive index, f(n) as given in Eq. (2).

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Figure 6

3D plot for Buckingham equation.

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Combination of f(ε) and f(n) together is called as Orientation Polarizability. Lippert Mataga equation (Reichardt and Welton, 2011) shows a linear correlation between Frequency Shift (Change in position of CT band from the respective acceptor) and the Orientation Polarizability of corresponding solvents. Fig. 7 shows good linear correlation between frequency shift and Orientation Polarizability with R² value of 0.971.

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Figure 7

Graphical plot for Lippert Mataga equation.

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4.2.1 Molecular geometries

The optimised geometries of 2-Chloropyridine-ICl in all the media are shown in Fig. 8 whereas selected geometrical parameters are tabulated in Table 4. It is seen from Fig. 8 that the Cl atom of acceptor ICl molecule stays a bit away from the substituent Cl atom in the donor 2-Chloropyridine molecule with an angle greater than 120⁰. The reason may be the steric repulsion between the two Cl atoms. Moreover, it is seen from Table 4 that the <C(6)N(1)I(11) angle decreases with increase in dielectric constant of the solvent. However, this decrease in the angle is not very large. For instance, <C(6)N(1)I(11) angle in gas phase and CH₂Cl₂ solvent with dielectric constant 9.08 are 127.89° and 125.85° respectively. This observation may be attributed to slight decrease in steric repulsion between the two Cl atoms with solvent polarity. Moreover, inspection of N–I and I–Cl bond length from Table 4 indicates that N–I bond length decreases with the increasing dielectric constant of the solvent and the I–Cl bond distance increases with the increasing solvent dielectric constant. The N–I bond length in the gas phase and the increase in the I–Cl bond length due to complexation in the gas phase are close to those of previously reported values for Pyridine-I₂ CT complex (Reichardt and Welton, 2011).

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Figure 8

Optimized geometries of 2-Chloropyridine-ICl.

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Decrease in the bond length between the bridging atoms of donor and acceptor molecules with complexation and the increasing polarity of the mediums is already reported for similar compounds (Tiwary and Mukherjee, 2009; Reiling et al., 1997; Tiwary et al., 2008). From Table 4 and previous reports we thus can conclude a stronger donor acceptor interaction in polar medium.

We have tried to examine the probability of a quantitative relation between the N–I and I–Cl bond length with that of dielectric constant. Thus we made a plot of bond length against dielectric constant values for both the selected bonds under study. Satisfactorily we obtain a logarithmic relation between the bond length values with the corresponding dielectric constant values of the medium as seen in Fig. 9 with an R² value of 0.997 for both the plots.

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Figure 9

Dielectric constant of medium against bond length, (a) dielectric constant against r_N–I, (b) dielectric constant against r_I–Cl.

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4.2.2 Natural bond orbital analysis

A single best Lewis type description of the total electron density is obtained from Natural Bond Orbital (NBO) calculations. Therefore, to understand the bonding between 2-Chloropyridine and ICl, we have performed NBO calculation on the complex between 2-Chloropyridine and ICl. In Table 5, we have tabulated the electron occupancy of the lone pair (LP) of the nitrogen and electron occupancy of valence antibond (BD^∗) of N–I. Moreover, the interaction energy (E⁽²⁾) between LP of N and BD^∗ of N–I from second order perturbation analysis is also tabulated in Table 5. It is observed that electron occupancy of LP of N decreases with increase in solvent polarity. For instance, electron occupancy of LP of N in gas phase and in CH₂Cl₂ solvent is 1.83e and 1.73e respectively. On the other hand, electron occupancy of BD^∗ of N–I increases with increase in the solvent polarity. This behaviour indicates that electron transfer from LP of N takes place and BD^∗ of N–I increases with the increase in solvent polarity. On inspection of E⁽²⁾ values from Table 5, it is clear that stabilization interaction between LP of N and BD^∗ of N–I increases with the increase in the solvent polarity.

4.2.3 Natural resonance theory analysis

Natural resonance Theory is performed to analyse the total electron density in terms of a series of idealised resonance forms. Moreover, NRT theory provides the percentage of ionic (%ionic) and covalent (%covalent) character in a bond. The %ionic and %covalent characters are tabulated in Table 6. It is observed that in gas phase, the N–I bond is almost ionic (∼94%). With increase in solvent polarity, %ionic character decreases whereas %covalent character increases. For instance, in hexane, %ionic and %covalent characters are 92.41 and 7.59 respectively, whereas in CH₂Cl₂, %ionic and %covalent are 74.63 and 25.37 respectively. The decrease in %ionic and increase in %covalent indicates increase in stability of N–I bond with solvent polarity.

4.2.4 Natural energy decomposition analysis

To decompose the interaction energy (IE) to a sum of physically meaningful term, we have performed Natural Energy Decomposition Analysis (NEDA) by Glendenning and coworkers (Glendening et al., 2013). In Table 7, we have tabulated IE, Energy of charge transfer (E_CT) and Polarisation Energy (E_POL). It is seen from Table 7 that IE in gas phase is −47.99 kJ/mol. However, in hexane solvent with dielectric constant 1.89, IE drastically increases to −261.75 kJ/mol and in CH₂Cl₂ solvent with dielectric constant 9.08 IE increases to −1099.09 kJ/mol. This indicates that interaction energy increases with the increase in the solvent polarity. Inspecting the E_CT terms from Table 7, it is inferred that charge-transfer increases with the increase in the solvent polarity. Similar to E_CT, E_POL term also increases with the increase in the solvent polarity. However, the variation of E_POL with solvent polarity is small, as compared to E_CT. For instance, E_CT and E_POL in hexane solvent with dielectric constant 1.89 are −283.05 and −120.33 kJ/mol respectively, whereas E_CT and E_POL in CH₂Cl₂ solvent with dielectric constant 9.08 are −754.67 and −217.69 kJ/mol respectively.

4.2.5 QTAIM analysis

QTAIM analysis has been performed to analyse halogen bonded interaction between 2-Chloropyridine and ICl. The properties of the (3,−1) Bond Critical Point (BCP) between two atoms indicate the types of interaction between the two atoms. The various properties of BCP between N atom of 2-Chloropyridine and ICl are tabulated in Table 8. According to QTAIM, electron density (ρ(r)), Laplacian of electron density $({\nabla }^2\rho (r))$ and electronic energy density (H(r)) at BCP are valuable parameters for probing the nature of the bond (Bader, 1990, 1991). Generally, for covalent interactions, ρ(r) is large, while ${\nabla }^2\rho (r)$ is large and negative. On the other hand, for ionic interaction, ρ(r) is small, while ${\nabla }^2\rho (r)$ is positive. Moreover, the magnitude of H(r) reflects the covalence of interaction. A negative value of H(r) indicates a significant covalent contribution while a positive value of H(r) indicates a significant ionic contribution. It is observed from Table 8 that in gas phase, ρ(r) is small, and ${\nabla }^2\rho (r)$ is positive, which indicates that the interaction between 2-Chloropyridine and ICl is ionic. However, H(r) is negative, which indicates that there is also covalent contribution to the interaction. As expected, ρ(r) increases with increase in dielectric constant, whereas H(r) becomes more negative with increase in dielectric constant. This infers that, covalent contribution to the interaction increases with increase in dielectric constant.