Bond Energy and Bond OrderGeorgio Nebbia Citation: The Journal of Chemical Physics 18, 1116 (1950); doi: 10.1063/1.1747878 View online: http://dx.doi.org/10.1063/1.1747878 View Table of Contents: http://scitation.aip.org/content/aip/journal/jcp/18/8?ver=pdfcov Published by the AIP Publishing Articles you may be interested in A rotating bond order formulation of the atom diatom potential energy surface J. Chem. Phys. 95, 2216 (1991); 10.1063/1.460973 The Bond-Energy Bond-Order (BEBO) Model of Chemisorption J. Vac. Sci. Technol. 10, 89 (1973); 10.1116/1.1318049 Comparison of Relations between Covalent Bond Order, Energy, and Interatomic Distance for CarbonCarbon Bonds J. Chem. Phys. 17, 738 (1949); 10.1063/1.1747378 Dependence of Bond Order and of Bond Energy Upon Bond Length J. Chem. Phys. 15, 305 (1947); 10.1063/1.1746501 Relations between CarbonCarbon Bond Energy, Order, and Interatomic Distance J. Chem. Phys. 15, 77 (1947); 10.1063/1.1746296

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1116 LETTERS TO THE EDITOR

Bond Energy and Bond Order GEORGIO NEBBIA

University of Bologna, Bologna, Italy May 1. 1950

RECENTLY Kavanaul has compared the relations between bond order, bond energy and interatomic distance for

carbon-carbon bonds. Concerning his paper I should observe that bond energy and interatomic distance may be related in a rather different way from those reviewed by Kavanau, considering that all the C-C bonds do vary with continuity either in bond energy and in interatomic distance, the "pure" double bond of ethylene being but a particular case of the series of the bonds of organic chemistry, each with a certain grade of hybridization of the four binding electrons, starting from ethane and ending with acetylene.

keal mole

J: .... ~ !oJ oJ

200

~ 100 t-al

0 1.1

I I I 1.2 1.3 1.4 I.~ A

BOND L.ENGTH

FIG. 1. C -C bond.

TABLE 1.

Dissociation Bond length Bond energy energy"Q.

in A kcal/mole kcal/mole

Diamond n C -C 1.54 Q. obs. following Bichowski and

Rossini and Pauling =125 kcal/mole Ethylene 1 C -C 1.35

4 C-H 1.09 Q. obs. =444 kcal/mole

Acetylene 1 C -C 1.20 2 C-H 1.09

Q. obs. =300 kcal/mole Allene 2 C -C 1.33

4 C-H 1.09 Q. obs. =537 kcal/mole

Ethane 1 C -C 1.55 6 C-H 1.09

Q. obs. =583 kcal/mole Carbon monoxide 1 C -0 1.13

Q. obs. =211 kcal/mole Carbon dioxide 2 C -0 1.16

Q. obs. =338 kcal/mole Formaldehyde 1 C -0 1.21

2 C-H 1.09 Q. obs. =317 kcal/mole

Acetaldehyde 1 C-O 4 C-H 1 C-C

Q. obs. =561 kcal/mole Glyoxal 2 C-O

2 C-H 1 C-C

Q. obs. =547 kcal/mole Dimethyl ether 2 C-O

6 C-H Q. ob •. =668 kcal/mole

Hydrogen cyanide 1 C-H 1 C-N

Q. obs. =259 kcal/mole Cyanogen 1 C - C

2 C-N Q. obs. =404 kcal/mole

1.22 1.09 1.50

1.20 1.09 1.47

1.42 1.09

1.09 1.15

1.37 1.16

62.5

90 88

118 88

92 88

62 88

211

169

145 88

145 88 62

153 88 64

78 88

88 168

82 160

125 125

90 352 442 118 176 294 184 352 536 62

528 590 211 211 338 338 145 176 321 145 352

62 559 306 17(>

64 546 140 528 668

88 168 256 82

320 402

TABLE L-(Continued)

Dissociation Bond length Bond energy energy Q.

inA kcal/mole kcal/mole

Methyl acetylene C-C 1.21 113 113 C-C 1.46 70 70

1 C-H 1.057 90 90 3 C-H 1.09 88 261

Q. obs. =536 kcal/mole 537 Cycloexane 12 C-H 1.09 88 1056

6 C-C 1.53 57 342 Q. obs. = 1397 kcal/mole i398

Methyl cyanide 1 C-C 1.49 60 60 I C-N 1.16 148 148 3 C-H 1.09 88 264

Q. obs. =470 kcal/mole 472 Benzene 6 C-C 1.39 85 510

6 C-H 1.09 88 528 Q. obs. = 1039 kcal/mole i038

Pyridine 5 C-H 1.09 88 440 4 C-C 1.39 85 340 2 C-N 1.37 79 158

Q. obs. =938 kcal/mole 939 Pyrrole 1 N-H 84 84

4 C-H 1.09 88 352 2 C-N 1.42 65 130 2 C-C 1.35 87 174 I C-C 1.44 70 70

Q. abs. =811 kcal/male 810 Pyrazine 2 C-C 1.39 79 15S

4 C-N 1.35 81 324 4 C-H 1.09 88 352

Q. abs. =833 kcal/mole 834 Acetamide I C-C 1.51 66 66

1 C-N 1.38 89 89 1 C-O 1.28 120 120 3 C-H 1.09 88 264 2 N-H 84 168

Q. abs. =708 kcal/mole 707 Urea I C-O 1.25 118 118

2 C-N 1.37 83 166 4 N-H 84 336

Q. abs. =620 kcal/mole 620 Furan 1 C-C 1.46 70 70

2 C-C 1.35 86 172 2 C-O 1.40 88 176 4 C-H 1.09 88 352

Q. abs. = 770 kcal/mole 770

To a C-C bond is therefore associated an energy and an inter­atomic distance related between them by the curve shown in Fig. 1 and drawn with the data of Table I, from which has been drawn also the curves of Figs. 2 and 3 for the C-O and the C-N bonds, respectively.*

The dissociation energy of a molecule is given directly by the sum of the various bond energies present in the molecule itself and, in any way, avoiding the concept of single, double and triple bond energy, the resonance energy is "lost." Furthermore the following picture, apparently more physically correct, could be used instead of that of bond order. The interatomic distance could be related to the sum of the ".-electron charges upon the bonded

kgJ mole

~ II: !oJ Z !oJ

C Z g

1.1 1.2 1.3 1.4 15 A

BOND L.ENGTH

FIG. 2. C -0 danb.

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LETTERS TO THE EDITOR 1117

keel mole

200

~~ .... . ; 100 • m •

o _I I IJ 1.2 1.3 1.4 1.5 A

SOND LENGTH

FIG. 3. C-N bond.

atoms, obtainable with the molecular orbital treatment.2 There exist many uncertainties for a relation of this type since we know but little about the true charges in ",-electrons near the linked atoms and about the C-C distance in many conjugated systems. We know, e.g., that the 2 and 3 atoms of butadiene have low charge and that their distance is high (1.46A), and that the mesocphenan­threne-type atoms have high charge and short bond (1.35A in pyrene, 1.38A in dibenzanthracene and coronene, following Robertson).3

On the other hand the calculations give (1.00+ 1.(0) ",-electrons both for the ethylene carbon atoms and for two carbon atoms linked together in benzene, the interatomic distances being 1.35 and 1.39A, respectively. This may be due to our too rough calcula­tions of charge densities. The limiting cases should be acetylene with 4 ",-electrons upon its two carbon atoms (1.20A) and ethane with zero ",-electrons upon his two carbon atoms and with C-C distance 1.54A.

* In such curves the heat of vaporization of carbon has been assumed. for an easier comparison with the well-known data of Pauling (L. Pauling. The No,ture of the Chemical Bond (Cornell University Press. Ithaca. New York. 1939)). as 125 kcal/mole. following the old data of Bichowski and Rossini (F. R. Bichowski and D. R. Rossini. The Thermochemistry of the Chemical Substances (Reinhold Publishing Corporation. New York. 1936). In any way the data may be transformed. assuming a new heat of vaporiza­tion (see reference 1 and also K. S. Pitzer. J. Am. Chern. Soc. 20. 2140 (1948)). by a simple shifting of the curves. All the heats of dissociation have been recalculated from the heats of combustion. corrected to the gaseous state of the molecules. referred to hy Wheland (G. W. Wheland. The Theory of Resonance (John Wiley and Sons. Inc .• New York. 1944). From the same author have been obtained the interatomic distances.

I J. L. Kavanau. J. Chern. Phys. 17. 738 (1949). 'C. A. Coulson and H. C. Longuet-Higgins. Proc. Roy. Soc. (London)

A191. 39 (1947). • J. M. Robertson. Acta Crystall. 1. 101 (1948).

Electronic Kinetic Energy in Gases at High Pressures

T. L. COTTRELL Research Department. 1. C. 1. Ltd .. Nobel Division.

S/evenston. Ayrshire. Scotland May 31.1950

APPLICATION of the virial theorem to a substance of volume V and pressure p gives

tiT=3ti(PV)-tiU, (1)

where T is the average kinetic energy of electrons and nuclei and U is the total energy.l Michels et al.2 have used this result to calcu­late changes in average kinetic energy brought about by com­pression of gases; ti(p V) was directly measured and tiU was calculated from the observed P-V-T relationships. tiT was found to increase with pressure, and it was pointed out that this must be due to an increase in electronic kinetic energy. It was shown that a change of the correct order of magnitude in electronic kinetic energy could be demonstrated by solving the wave equa­tion for the hydrogen atom with the boundary condition that >It=0 at r=ro.2.8 It is of interest to try to interpret these results

on a molecular basis, particularly because such an interpretation may enable one to discuss the effects of pressures greater than those used by Michels (3X103 atmos.).

Slater4 has pointed out that, for a diatomic molecule,

T= -E-R(dE/dR), (2)

where E is the electronic energy, as a function of the internuclear distance R. This shows that as the nuclei are moved closer together than the eqUilibrium distance, there is a sharp rise in kinetic energy, and a relatively fiat rise in total energy. Thus, if the com· pression of a gas composed of diatomic molecules could be re­garded as pushing the nuclei a little closer together, the qualita­tive explanation of the effect is clear. The large effect of compres­sion on the kinetic energy is very striking, but it is notable that a compression which increases the kinetic energy by an amount as great as the dissociation energy of the molecule leads to only a small increase in total energy. Using the E(R) function and con­stants given by Linnett,· we may use (2) to calculate how E and T vary with R for nitrogen:

R(A) E(kcal./mole) 1.094 0.00 1.092 0.01 1.050 3.45

AT(kcal./mole) 0.00 6.9

172.

The force required to bring about a given change in Rand hence in T can thus be obtained from spectroscopic data; more­over, the corresponding pressure required can be obtained from Michels' experimental work. From this force and pressure the "effective area" of the molecule can be calculated. This "effective area" should be about the same as that of the van der Waals sphere, but close agreement is not to be expected: Michels' re­sults on the hydrogen atom suggest that the area might be greater. The results for nitrogen are of the correct order, and their approximate constancy with increasing pressure suggests that this viewpoint is justifiable. They are as follows:

AT(kcal.) 1 2 3 4 5 6 Area (A') 143 151 152 149 145 142.

The "effective area" seems to be decreasing at high pressures; thus tiT calculated from it for still higher pressures is likely to be an upper limit. In nitrogen at 105 atmospheres, for example, 180 kcal. is an upper limit to the change in kinetic energy, and 4 kcal. to the change in energy. Dissociation should take place at about 106 atmospheres.

1 Schottky. Physik. Zeit •. 21,232 (1920). 2 Michels et ai .• Physica 4, 981 (1937). 8 Michels and de Groot. ibid. 16. 183 (1950). • Slater. J. Chern. Phys. 1.687 (1933). • Linnett, Trans. Faraday Soc. 36. 1123 (1940) •

Configuration of the Tartaric Ion CL!!MENT DUVAL

Laboratoire de Chimie AND

JEAN LECOMTE

Laboratoire des Recherches Physiques, Sorbonne. Paris. France June 14, 1950

I N a Letter to the Editor, E. E. Turner and K. Lonsdale1 have recently called attention to the two possible configurations for

d-tartaric acid in the crystalline state, H-atoms and OH-groups being arranged approximately in a plane and the carboxyl groups staying either both in front of behind the plane of the paper, or one in front and the other behind. We think the second configura­tion is the only correct one.

Some years ago, we recorded and studied2 the infra-red absorp­tion spectra, between the wave-lengths of approximately 6-15 1', for a series of metallic tartrates (11 dextro, 7 laevo, 8 meso and 5 racemic) in the powder state. d- (and 1-) tartrates have a spectrum formed of numerous bands, whose number is greafly reduced in the case of mesotartrates and racemates; the individual positions of bands differ in the three cases of d (or 1), meso, and racemic.

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