Physical meaning of electrooptical parameters derived from infrared

J . Phys. Chem. 1984,88, 600-604

600 approximation, for this molecule.

Downloaded by TEXAS A&M INTL UNIV on August 30, 2015 | http://pubs.acs.org Publication Date: February 1, 1984 | doi: 10.1021/j150647a053

Conclusions We refer to the spectroscopic observations mentioned at the beginning and which require some further study: (1) For the specific case of 1,1,1,4,4,4-hexadeuteriobutane studied in this work, the calculations, which include Fermi resonances as well as electrooptical parameters for predicting intensities, account satisfactorily for the observed spectra in both the polarized and depolarized cases, for both T and G structures. (2) For the d- band near 2900 cm-’ we neglect, as already stated, the problem of the natural breadth of the band, which will be treated elsewhere. From our calculations, the d- band has to shift upward in going from T to G structures, if the unperturbed fundamentals are different from T to G. Attempting a generalization to longer hydrocarbon chains, the increase of content of G structures may generate a broader multiplet structure in the polarized spectra centered at a frequency above 2900 cm-’, which increases along with the increase of G content. This seems to agree

with the experimental observations in longer hydrocarbon chains.2,10 (3) With regard to the meaning and use of the ratio R = I2940/I2850, we consider this work as a step forward, with respect to our previous work, in disentangling the various contributions and refer to later work for a more thorough assessment of its use as structural probe for longer molecules. R has been tested and explained on CD3CH2CH2CD3and we think it is usable in compounds with no C H 3 end groups such as polyethylene and compounds with CD3end groups, from which the contribution of CH3, as pointed out by Snyder, has been removed. Moreover, this ratio should be measured in the polarized spectra of liquid samples, where the contribution of the d- mode is smaller.

Acknowledgment. This work was supported by Consiglio Nazionale delle Ricerche, Programma Finalizzato Chimica Fine e Secondaria, Sottoprogetto Polimeri. Registry No. 1,1,1,4,4,4-Hexadeuteriobutane,13183-67-0.

Physical Meaning of Electrooptical Parameters Derived from Infrared Intensities M. Gussoni, Unit6 di Spettroscopia Molecolare in Milano dell’lstituto CNR, di Chimica Applicata dei Materiali di Genoua, Milano, Italy

C. Castiglioni, and G. Zerbi* Dipartimento di Chimica Industriale e Ingegneria Chimica del Politecnico di Milano, Milano. Italy (Received: June 3, 1983)

It is shown that infrared intensities, when interpreted in terms of the model of electrooptical parameters (EOP), provide experimental atomic charges q / at the equilibrium geometry and charge fluxes aqa/aRtduring the vibrational motion; these parameters account for several physical and chemical properties of molecules. q / and dq,/dR,, calculated from the vibrational intensities of many molecules, are compared and discussed. Various properties associated with electronic charge on the atoms are defined, classified, compared, and used in predicting the chemical behavior of molecules.

Introduction

and a p k / a R t are the so-called EOP. Both APT5 and EOP6 have been shown to be useful for predicting spectra; the EOP, due to their localized meaning, have been particularly successful for predicting spectra in sophisticated applications such as the interpretation of the spectra of geometrically disordered molecules.6 A particular kind of EOP describes atomic charges qm0at the equilibrium geometry and their fluxes dq,/aR, during vibrations R,; they were introduced by van Straten and Smit’ as derived from standard EOP, but they are the same as those previously introduced in a different way by Decius; hence, we call them equilibrium charges-charge fluxes (ECCF) as proposed by Decius. The ECCF parameters can be derived from EOP by assuming’ /.Lk = qkrk where q k is the bond charge on bond AB ( q A = -qk; qB= qk) and rk is the interatomic distance between atoms A and B. The ECCF of course have the same capability as EOP for predicting intensities and give a more straightforward description of the charge distribution in the molecule, as can be inferred from observed intensities (together with the observed molecular dipole moments). In this paper we will discuss the physical meaning of ECCF. It is w e l l - k n o ~ nthat ~ ~ ’ the ~ reliability of any kind of intensity kko

Considerable effort has been made in the interpretation of infrared intensities in the last decade.’ Use is generally made of parameters derived from infrared intensities of small molecules both for elucidating the charge distribution within the molecule and for predicting the intensities of larger ones. The parameters used most often are the atomic polar tensors (APT)Z and the electrooptical parameters (EOP).3 The former are derivatives of the molecular dipole moments with respect to atomic Cartesian displacements; the latter are bond dipole moments at the equilibrium geometry and derivatives of bond dipole moments with respect to internal coordinates (Le., vibrational coordinates as defined in ref 4). The definition of electrooptical parameters2tarts from the assumption that the molecular dipole Foment M can be expressed as a sum of bond dipole moments, M = &.& ( z k = unit vector of the k-th bond). Thus, the derivative of M with respect to the internal vibrational coordinate R, can be written as ah?/aRt = C(zko(aMk/aR,) -k

/.LkO(aak/aR,))

k

(1) W. B. Person and G. Zerbi, Eds., “Vibrational Intensities in Infrared and Raman Spectroscopy”, Elsevier, Amsterdam, 1982. (2) W. B. Person, ref 1, Chapter 4. (3) M. Gussoni, ref 1, Chapter 5. (4) E. B. Wilson, J. C. Decius, and P. C. Cross, “Molecular Vibrations”, McGraw-Hill, New York, 1955.

(5) W. B. Person, ref 1, Chapter 14. (6) G.Zerbi, M. Gussoni, S . Abbate, and P. Jona, ref 1, Chapter 15. (7) A. J. van Straten and W. M. A. Smit, J . Mol. Spectrosc., 62, 291 (1976). (8) J. C. Decius, J . Mol. Spectrosc., 57, 384 (1975).

0022-3654/84/2088-0600$01.50/00 1984 American Chemical Societv

EOP Derived from IR Intensities

1.0

a.

The Journal of Physical Chemistry, Vol. 88, No. 3, 1984 601

2.0 I

'I

*


F

Figure 1. Experimental atomic equilibrium of halogens in CH3aplotted vs. (a) difference of electronegativity between carbon and halogen, (b) stretching force constant of the bond aC,and (c) interatomic distance aC.

1:O

112

1:4

1:6

R:clA1

Figure 4. Experimental atomic equilibrium charges1* of hydrogen in hydrocarbons plotted vs. interatomic distance of the adjacent CC bond.

parameter has severe limitations caused both by the large errors in intensity measurements an$ by the lack of knowledge (from experiments) of the sign of 8M/8Qi (derivative of the molecular dipole moment with respect to the normal coordinates). These limitations are drastically reduced (and sometimes removed) when the data of several isotopic species and of related molecules are used together with those of the parent m o l e c ~ l ea;nd ~ ~when ~~ various criteria for the determination of the sign of 8A4/8Qi are used." In this paper use will be made only of ECCF derived from such a "complete" analysis of the data. Let us stress that all the q: and aq,/aR, reported in this paper are "experimental", namely, are derived from observed infrared intensities and observed molecular dipole moments.

Atomic Equilibrium Charges

Figure 2. Experimental atomic equilibrium chargesi6*" of halogens in H a plotted vs. (a) difference of electronegativity between hydrogen and halogen, (b) stretching force constant of the bond, (c) interatomic distance, and (d) bond dissociation energy. C3H4 cZ H4

CaHc

C@

H$

CtgCH CH,CI CH3Br CHgl CH,F

The atomic equilibrium charge q: represents the gain (q: < 0) or loss (q: > 0) of electronic charge on atom a resulting from the bond of a with another atom p; for diatomic molecules qo = M/R0 (AI" = equilibrium molecular dipole moment; Ro = equilibrium interatomic distance). From its very definition qao provides a measure of the fractional ionic character12 of the bond cup; it must therefore be related to the difference Ix, - xBI of ele~tronegativity'~ between a and p. It is pleasing to see that the quo derived from observed infrared intensities (together with the molecular dipole moments for polar molecules) are linearly related to Ix, - xBI. The value of x, for many atoms has been derived by Paulingi3 in such a way that Jx, - xBlzis proportional to DaB, the ionic part of the bond energy. In a model based on point charges,14 D,, can be expressed as proportional to (qa:)Z/ra2 (where; ,4 = 1q21 = Iq;] and therefore Ix, - xBl is proportional to qM0. In Figure l a the :4 of CH3a( a = F, C1, Br, I) are plotted vs. Ix, - xcl; the plot is strictly linear. The same is found for q: in H a ( a = F, C1, Br, I) as shown in Figure 2a. The plot of qHo vs. lxH - xBl ( p = Se, S, I, C, Br, N, C1,0, F), shown in Figure

HF

eNH, HCl

I

1,o

' 2.0

1 % - X I

a

Figure 3. Experimental atomic equilibrium chargesl6-I8 of hydrogens bonded to different atoms a plotted vs. difference of electronegativity between a and H .

(9) M. Gussoni in 'Advances in Infrared and Raman Spectroscopy", Vol. 6, R. J. H. Clark and R. E. Hester, Eds., Heyden, London, 1980. (10) G. Zerbi, ref 1, Chapter 3. (1 1) S. Abbate and M. Gussoni, Chem. Phys. 40, 385 (1979). (12) R. McWeeny, 'Coulson's Valence", Oxford University Press, London, 1979. (1 3) L. Pauling, "The Nature of the Chemical Bond", Cornell University Press, Ithaca, NY, 1960. (14) M. Gussoni, C. Castiglioni, and G. Zerbi, Chem. Phys. Lett., 99, 101 (1983). (1 5) M. Gussoni, S. Abbate, and G. Zerbi, to be submitted for publication. (16) C. Castiglioni, Thesis, Universitl di Milano, Milano, Italy, 1982. (17) R. E. Bruns and R. E. Brown, J . Chem. Phys., 68, 880 (1978). (18) M. Gussoni, P. Jona, and G. Zerbi, J . Chem. Phys., 78,6802 (1983).

602

Gussoni et al.

The Journal of Physical Chemistry, Vol. 88, No. 3, 1984

TABLE I: Experimental Atomic Charges 4', Experimental Effective Charges x, and Net Charges Qn Calculated by Mulliken Population Analysis with "ab Initio" Methodsa 9O

0.42 0.33 0.27 0.22 0.21 0.13 0.10 0.07 0.06 0.09 0.09 0.08 -0.23 -0.17 -0.15 -0.12


0.10

a Units: e. ence 22.

Ixlb

Qnc

0.39 0.25 0.17 0.21 0.20 0.10 0.17 0.09 0.09

0.39 0.33 0.26 0.27 0.23 0.13 0.09 0.12 0.1 1

Q'

0.18 0.23 0.23 0.24

0.07 0.08 0.08 0.56

-0.34

0.31 0.21 0.12

-0.14 -0.15 -0.13

References 15 and 20.

Reference 21.

Figure 5. Experimental atomic equilibrium charges16,18in several organic compounds, plotted vs. stretching force constant of C H bond.

Refer-

0 HF

0 H,O

\

\

0.0

0.2

0.1

0.3

0.4

$(el

Figure 7. Experimental atomic equilibrium chargesl6-I8 of hydrogen hydrogens with "acidic" behavior; (0) atoms in different molecules: (0) hydrogens with "neutral" behavior. AE is the formation energy of the dimer (for sources of the data see ref 25).

Figure 6. Experimental atomic equilibrium atomic charges'6i18of hydrogens in several organic compounds plotted vs. interatomic distance CH.

3, is even more impressive; here the bond HP belongs to quite different molecules, but the general trend is found to be still roughly linear, as if qHowere a measure of the electronegativity of the partner (in all cases presented in Figure 3 is more electronegative than H itself). The large scattering of qHovalues for hydrogen atoms bonded to carbon atoms suggests an electronegativity of C which varies according to its state of hybridization (more electronegative in sp than in spz and even more than in sp3); a plot (Figure 4) of qHovs. Rccoof the adjacent C-C bond supports this statement. More support for the intrinsic meaning of q," is provided by a comparison of qaowith the bond constant Kat (Figures lb, 2b, and 5). Kas is expected to increase with q," within a homogeneous series (same a);the plots of qao vs. Kas are linear. Also the comparison of 4," with R,: is encouraging (Figures IC, 2c, and 6). The plot is linear, fulfilling the expectation that the interatomic distance decreases when the charge increases. Also the observed dissociation energy A E of diatomic molecules H a (a = F, C1, Br, I) increases with increasing 4010, as expected (see Figure

2d). Since atomic equilibrium charges represent the gain or loss of charge on an atom, they should be comparable with the so-called Mulliken population analysis19 of molecular orbitals. A comparisonzOof q," with Q," in a number of cases where the Qao have been computedz1from molecular orbitals built on different basis sets has revealed that Q," values approach Q," with increasing completeness of the basis set. The values of experimental q," and the corresponding Qao calculated by high-level "ab initio" methods are reported in Table I for several cases and show a remarkable agreement. Also the effective charges xa are reported in Table I; xa is defined by Xa

= 1/3C(pUua)z uu

u, u = x, Y , z

where PUuaare the components of P",the APT of atom a. xa is a very interesting quantity because it can be directly calculated by the sum of the i n t e n s i t i e ~ when , ~ ~ data for isotopic species are also available. Kingz3has tried to compare xa and Qao for several atoms with results that are not completely satisfactory. Indeed, we can now state that xa can be compared with Q," only when xa 9010 because qao has the same meaning of Q," while xa (19) R.S. Mulliken, J . Chem. Phys., 23, 1833, 1841,2338,2343 (1955). (20) M. Gussoni, C. Castiglioni, and G.Zerbi, Chem. Phys. Lett., 95, 483 (1983). (21) S. Marriott and R. Topsom, J. Mol. Struct. (Theochem), 89, 83 ( 1982). (22) T. B. Lakdar, E. Tallandier, and G . Berthier, Mol. Phys. 39, 881 (1980). (23) W. T. King, ref 1, chapter 6 .

EOP Derived from IR Intensities

The Journal of Physical Chemistry, Vol. 88, No. 3, 1984 603

TABLE 11: Calculated and Observed Formation Energy AE (kcal mol-’) for Some Molecular Complexes of Known Geometry

EDa H,O H2O HCN HF HCN a


14.

EAb HF HOH HI: HF HCN

AEcalcdC

aEobsdd

-7.9 -3.7 -5.4 -3.9 -3.9

-6.2 t 1.8 -5.2 i 1.5 -4.4 i 0.3 -6.0 t 1.0 -3.3

ED = electron donor. EA = electron acceptor. For sources of these data see ref 14.

AEcalcdd -7.7 -5.9 -3.6 -5.6 -3.7 Reference

includes also charge flux contributions. Indeed xolz q2 only when the charge is roughly undeformable, as discussed in the next section. The experimental qaoshould also be very useful in the interpretation of intermolecular interactions which are mainly due to electrostatic forces24(weak hydrogen bonds and molecular complexes). Indeed, it has been remarkedZSthat the values of qHo for hydrogens in different molecules characterize their capability to form hydrogen bonds: only when qHo> 0.2 e is the hydrogen “acidic”; otherwise, it has a “neutral” behavior with respect to an electron donor (Figure 7). To be more precise, we may state that, when a hydrogen is bonded to a more electronegative atom, it effectively releases “part of its electron” in forming the bond; this results in a qHo > 0, where qHo increases with increasing electronegativity of the partner (Figure 3). Only when qHo> 0.2 e does the hydrogen “feel the attraction” of an electron-donating atom in another molecule and form a hydrogen bond. The dissociation energy AE of the dimer increases with increasing qHo (Figure 7). Hydrogens capable of forming hydrogen bonds also show peculiar characteristics in the charge fluxes, namely, in the deformability of their charges during the motion; this will be discussed in the next section. An attempt to evaluate AE from the knowledge of experimental q2 for several molecular complexes having known geometry has proven to be successf~l’~ (Table 11), thus adding more reliability to experimental charges. For some molecular complexes both molecules can act in principle as electron donor or electron acceptor; the experimental qaoare a good tool for predicting the energetically favored ~tructure’~ when the geometry of the complex is not known. A last successful application of experimental q2 can be found in the interpretation of infrared circular dichroism spectra.26 Chiral molecules exhibit small, but significant, differences in their spectra when light is circularly polarized in the two opposite senses; these differences are important data for the determination of the intramolecular conformation but their interpretation depends on the knowledge of equilibrium charges and charge fluxes. Charge Fluxes We have discussed in the previous section the meaning and reliability of the experimental equilibrium atomic charges. We wish now to discuss similar correlations among the fluxes of charge during the vibration. However, the ideas in this field are not yet as clear as in the case of the equilibrium charges. Two or three remarks, emerging from the data, can be discussed here. The first and most general one concerns the deformability of the charge. When all the charge fluxes are small with respect to the equilibrium atomic charge, then the charge is undeformable.’7J8 It has been shownls that a good estimate of the undeformability of the charge on atom a is provided by the ratio Fa//3J where pa is the average moment2 and ,f3, is the anisotropy* of the polar tensor Pa. When F/pa1 > 1, the charge is almost undeformable and the polar tensor approximates a scalar matrix qol0I, as it should be for an atom which carries a fixed charge. In these cases it is also true that xa E qao. (24) P. Schuster, W. Zundel, and C . Sandorfy, Eds., “The Hydrogen Bond“, North-Holland Publishing Co., New York, 1976. (25) M. Gussoni, C. Castiglioni, and G. Zerbi, J . Chem. Phys., in press. (26) S. Abbate, L. Laux, J. Overend, and A. Moscowitz, J. Chem. Phys.,

75, 3161 (1981).

TABLE 111: Experimental Values of Atomic Equilibrium Charges qCYo (e), Effective Charges xCY(e), Principal Charge Flux

aqa/arCYo(e K ’ ) ,and Undeformability Parameter’n tP,/&for CY

= H, 0, F, CI in Different Compounds Go

HBr HC1 C2Hz HCN CH,OH HZO HF

1x1

aqiar

Hydrogens (Undeformable) 0.12 0.11 -0.01 0.18 -0.00 0.18 0.21 0.20 -0.02 0.22 0.22 -0.02 0.25 0.20 -0.11 0.33 0.25 -0.22 0.41 0.38 -0.11

lisiol

6.00 13.00 9.57 19.74 1.31 1.19 3.90

Hydrogens (Deformable) 0.06 0.09 -0.19 0.05 0.10 -0.21 0.05 0.10 -0.21 0.08 0.10 -0.19 0.13 0.10 -0.22 0.12 0.10 -0.20 0.08 0.08 -0.16 0.09 0.08 -0.17 0.09 0.07 -0.18 0.06 0.09 -0.21 0.05 0.11 -0.17 0.16 0.17 -0.11 0.05 0.04 -0.04 0.10 0.16 -0.32 0.27 0.17 -0.28

0.02 0.14 0.15 0.13 0.1 3 0.13 0.03 0.01 0.00 0.07 0.04 0.21 0.5 1 0.15 0.70

so2

-0.67 -0.55 -0.31 -0.23

Oxygens 0.49 0.27 0.59 0.11 0.56 -0.39 0.50 -0.41

1.49 1.04 1.18 0.86

NaF LiF HF CH,F CH,I., C*H,F,

-0.88 -0.84 -0.42 -0.23 -0.17 -0.08

Fluorines 0.89 0.86 0.39 0.58 0.61 0.60

-0.02 -0.04 0.1 1 -0.47 -0.55 -0.45

27.80 12.29 3.90 0.71 0.59 0.46

-0.78 -0.73 -0.18 -0.17

Chlorines 0.85 0.83 0.18 0.31

-0.07 -0.13 -0.00 -0.17

5.34 3.04 13.07 0.94

CH, ‘ZH6

C,H, C3H6

C2H4 6‘

H6

CH,I CH,Rr CH,C1 CH,F CH, F, C2H2F2 HI CH20 “3

H*O CH,OH CH,O

NaCl LiCl HC1 CH,Cl

a

a

b c

c c a

b b b

e b b d d d d e f’ a

e c

C

C

C

R a U

a

d e

f’ a a U

d

a Reference 17. Reference 18. Reference 16. In this thesis, available upon request, data from different sources have been discussed and ECCF have been computed. Reference 15. e Reference 21. Reference 28. g Reference 7.

All the hydrogen atoms which form hydrogen bonds are carriers of a relevant charge (>0.2 e) as already discussed in the previous section. Moreover, their charges show little deformability during the motion. In Table I11 the hydrogens have been divided into undeformable and deformable ones: those of the first class (large charges, small fluxes, F/(?I> 1) generally form hydrogen bonds of relatively large formation energy (AI3 > 2 kcal mol-’); those of the second class (small charges, large fluxes, p/pl < 1) do not. Let us stress that the capability of forming hydrogen bonds seems to be related both to the magnitude of the chafge and to its undeformability.’4y25Two examples support this statement. First, HBr has an undeformable hydrogen, but its charge is small; thus, it does not form hydrogen bonds with significant energy. Second, the hydrogen of N H 3 has a large charge, but its deformation during the motion is also large; accordingly, N H 3 usually is not an electron acceptor, with the exception of the case of the dimer. In the last part of Table I11 the values of 40, x,aq/dr, and w//31 are reported for some atoms carrying a negative charge. Generally (27) J. Morcillo, L. J. Zamorano, and J. M. V. Heredia, Spectrochirn. Acta, 22, 1969 (1966). (28) R. 0. Kagel, D. L. Powell, M. N. Ramos, A. B. M. S. Bassi, and R. E. Bruns, J . Chem. Phys., 77, 1099 (1982).

604

The Journal of Physical Chemistry, Vol. 88, No. 3, 1984

Gussoni et al.

TABLE IV: Atomic Equilibrium Charges q a o (e) and Principal Charge Fluxes aqa/arap (e K ’ )for H, 0, F, and CI in Different Molecular Surroundingsa “Atomic” Bonds

________

_-___-___-

C-H

CH,

C,H,

C,H,

C,H,

CH,CCH

CH,O

C,H,

C,H,

C,H,

C,H,

4H”

aqH/arCH

0.06 -0.19

0.05 -0.21

0.05 -0.21

0.08 -0.22

0.10 -0.21

0.10 -0.32

0.12 .-0.20

0.13 -0.22

0.15 -0.20

0.21 -0.02

C-H

CH,F

CH,C1

CH,Br

CH,I

CH,FZ

qHn

0.06 -0.21

0.09 -0.18

0.09 -0.17

0.08 -0.16

0.05 -0.17

C,H,F, 0.16 -0.11

0-H

H,O

CH,OH

qHn aqH/arOH

0.33 .-0.22

0.25 -0.11

aqH/arCH

-_____________

X-H

HF

qHo aqH/arXH

0.41 -0.11

HC 1 0.18 -0.00

HBr 0.12 -0.01

HI

-

____________

__

“3

0.05

HCN

0.21 0.22 -0.02 -0.01 ___-___

________--

______-__________

___I___-___

CH,CCH

----_-

0.27 -0.28

.-0.04


“Ionic” Bonds

x-0

H,CO

SO,

4on aqo/arxo

-0.31 -0.39

-0.23 -0.41

X-F

NaF

LiF

H,CF

H,CF,

H,C,F,

C,F,

NF,

PF,

qFO

-0.88 -0.02

-0.84 -0.04

-0.23 -0.47

-0.17 -0.55

-0.08 -0.45

-0.03 -0.67

-0.09

x-c NaCl LiCl ___ __-____

-0.06 -0.80 -

H,CCI

aqF/arxF

qC1°

aqcl/arxcl

-0.78 -0.07

-0.73 -0.13

-0.97

__-_____-___________---

-0.17 .-0.17

a For sources of data see Table 111; data of NF, and PF, were taken from ref 7. For the definition of “atomic” and “ionic” bonds see text.

TAGLE V: Atomic Equilibrium Charge qaO(e) and Charge Fluxes aq,/aR, (e rad-’), Rt = Angular Deformation, for Some Moleculesa

C H , C C H ~qH0 = 0.10

a q H / a u = 0.00,

.

__

aqH/aa = -0.02 aqa/ap = 0.02

a For source of the data see Table 111. a is the adjacent angle and 01’ the opposite one. i n order t o avoid problems connected with local redundancies,’ we define the charge fluxes with respect to the following group coordinates:

Y-

a~~~

-X

H C Methyl group coordinates are defined as

speaking, it seems that also for these atoms the capability of forming hydrogen bonds of relevant energy is related to the existence of a large qo and of a small flux; for instance,z9H 2 0 is a stronger electron donor than C H 3 0 H and even stronger than CHzO while SOz is not, usually, an electron donor. All the charge fluxes dq,/dR, (R, = any vibrational coordinate) contribute to the deformability of the charge on a; however, the most relevant among the charge flux parameters is always the

so-called principal charge flux dq,/dR,,, namely, the flux of the charge occurring on a when its bond is stretched. When q2 and dq,/aR,, have opposite signs, atom a tends to recover the charge that it has released upon formation of the bond. This situation suggests a tendency of the bond to dissociate into neutral atoms. When quoand dq,/dR,, have the same sign, the bond a@ tends to dissociate into ions. On the basis of our considerations on the values of dq,/dR,, so far available, it seems that H is always involved in “atomic” bonds30(Table IV). The same atoms which form “atomic” bonds with hydrogen may form “ionic” bonds with other atoms (Table IV). Of course, this definition of ionic bond suggests only a trend and does not imply an effective dissociation into ions; in order that an effective dissociation into ions occurs, the qa0 must be very large. For instance, we expect N a F to dissociate into ions since the charges on the atoms are close to e+ and e- even at the equilibrium; we do not expect the same to happen in SOz. Other interesting features of charge distribution can be inferred from a comparison of bending charge fluxes dq,/dR, (R, = bending coordinate; see Table V). A discussion of these parameters is somewhat tentative since a much larger set of data is required. Nevertheless, we want to stress some important points: (i) The bending charge fluxes are always much smaller than the principal charge fluxes (compare Tables V and IV). (ii) The bending charge fluxes are larger in systems with IT electrons (compare dqH/dU = 0.002 of CzH6with dqH/d(YHCH = 0.05 of C2H4and of CHzO). The bending charge fluxes are even larger whenever atoms with lone pairs are involved in the bending (compare the above values with dqH/damH = 0.07 of NH3). (iii) It is pleasing to notice that d q H / d a H C H increases with CYHCH,in agreement with the increase of qHowhen going from C(sp3) (CX’HCH = 109.47) to C(spz) ( a 0 H c H = 120) (Figure 4). Acknowledgment. This work was partly supported by Consiglio Nazionale delle Ricerche, Progetto Finalizzato Chimica Fine e Secondaria. (29) G. G. Pimentel and A. L. McClellan, “The Hydrogen Bond”, Freeman, New York, 1960. (30) G. Herzberg, “Molecular Spectra and Molecular Structure (Diatomic Molecules)”, Van Nostrand, Princeton, NJ, 1955.

Physical meaning of electrooptical parameters derived from infrared

Recommend Documents