THE STRUCTURES OF THE FLUOROCHLOROlIETHAKES AND THE E F F E C T O F BOND TYPE ON CHEMICAL REXCTIT-ITY' L. 0 . BROCKWAY Gates and Crellin Laboratories of Chemistry, California Institute of T e c h m l o g y . Pasadena, Caltforniaz Receiiled Ocfober 16, 1935
The investigation of the structures of the nornial valence conipounds of the non-metallic elements was originally ptarted for the purpose of studying their itereocheniistry as shown by the bond angles observed in their compounds and the validity of a table of bond radii as indicated by observed bond distance.;. Results of the study of angles have shonn (2) that in the chlorine and methyl derivatives of the first row elements the bond angles are never more than about three degrees larger than the tetrahedral angle (109'28') observed in the symmetrical compounds of the 11Bb type. I n oxygen fluoride (OFJ as in water the bond angle is several degrees smaller than the tetrahedral value. The trifluorides and trichlorides of phosphorus and arsenic have angles from 101" to 104". Of greater importance in their relation to the chemical properties of substances are the bond distances. Several years ago Pauling and Huggins (8, 10) proposed a set of bond radii (table 1) whose sums represent the corresponding single-bond distances in conipounds in which the fourthgroup elements form four electron-pair bonds, the fifth group three, the sixth group two, and the seventh group one. The values for the fourth and seventh groups are half of the observed interatomic distances in the elements; the other valueq were obtained by interpolation with the aid of crystal structure data. The validity of these radii is supported by the observed distances for the methyl compounds (5). The values of the bond distances observed in eleven of these compounds are compared with the corresponding radius sums in table 2. I n 110 caw is the difference as great as the experimental error. These results show that a considerable degree of ionic character in Presented a t the Symposium on Molecular Structure, held a t Princeton 'Cniversity, Princeton, New Jersey, December 31, 1936 to January 2, 1937, under the auspices of the Division of Physical and Inorganic Chemistry of the American Chemical Society. Contribution S o . 573. 185
186
L. 0. BROCKWAE'
covalent bonds does not have an appreciable effect on the bond distances. Although the radii were obtained from bonds between like atoms they are apparently also applicable to bonds between atoms of quite different electronegatiyities, as illustrated here for several compounds in which appreciable electric moments are observed. TABLE 1 Sormal coisalent radii C
N
0.77
0 70
I
Si
I
1 17
0
0.66 ___
P
110
~
1
S
104
Ge
hS
Se
Sn
Sb
Te
1.40
1.41
1
1.37
F
1 i
0.64 c1
099
~
I
Br
I
1
1.33
TABLE 2 Bond distances and radius sums i n methyl compounds
c-c 1.55 f 0.02 1.54
si< 1.93 d~ 0.03 1.94 G e 4
1.98 f 0.03
2.18 & 0.03 2.17
I
1 1
1
N-c
1.47 f 0.02 1.47
'I
0 4
1.42 f 0.03 1.43 S-C
1.82 i: 0.03 1.81
I
F-C
1.42 f 0.02 1.41
c1-c 1.77 & 0.02 1.76 B
d
1.91 i 0.05
I I
Application of the radii to substances containing multiple bonds has shown that a new kind of conjugation frequently exists. Reference may be made to the chloroethylenes (4) for which the conventional bond structures are represented with a double bond b e t w e n the two carbon atoms and with single bonds connecting the chlorine and carbon atoms,
STRUCTURES O F THE F L V O R O C H L O R O M E T H l h ~ E ~
..
H
c1:
187
..
H
c1:
The observed carbon-chlorine distances in these six compounds are from 0.07 to 0.03 A.U. less than 1.76 X.U., the radius sum and the value observed in the chloromethanes, according to the number of chlorine atoms in the molecule. This shortening, which has also been observed in phosgene and thiophosgene, may be explained on the basis of contributions to the normal states of the respective molecules of structures in which the double bond is between carbon and chlorine,
/
H
\
H
/
H
\*. c1:
=In empirical quantitative relation between bond character and bond distance (4) has been worked out for all molecules in which the normal state is best represented as a resonance among several individual electronic structures such that a particular bond is single in one and double in another. Additional examples of substances of this type are benzene, graphite, cyanogen, carboxylic acids, carbonates, nitro compounds, nitrates, and many others. The normal chlorides and fluorides of a number of the non-metals have also been measured (6, l), and the observed distances are compared n i t h the corresponding radius sums in table 3. The chlorides of silicon, germanium, tin, phosphorus, and arsenic and the fluorides of carbon, silicon, phosphorus, and arsenic show distances which are smaller than the respective radius sums by amounts which in most cases are several fold greater than the experimental error. The shortenings in the chlorides have been explained on the basis of the contributions of structures containing double bonds. Such structures are formed if a chlorine atom shares one of its extra electron pairs with the central atom. In the chlorides of the first row elements the existence of only four bond orbitals on the central atom prohibits this phenomenon; it will be noted that the shortening does not occur in carbon tetrachloride or in chlorine monoxide. The amounts of the shortenings observed in the other chlorides correspond to reasonable degrees of double-bond character. The fluorides, on the other hand, show striking anomalies. I n silicon tetrafluoride and phosphorus trifluoride the observed shortenings of 15
188
L. 0. BROCKW'AT
and 13 per cent below the respective single-bond distances are too great to be due to the contribution of double-bond structures because the maximuni difference obserTed between single- and double-bond distances is 10 per cent. Carbon tetrafluoride, moreover, .hon-s a 4 per cent shortening, TABLE 3 Bond distances and radius sums in chlorades and fluorzdes
I
C-Cl I
1 755 1 76 0 00
=z!
0 005
I
2 00 i 0 02
p - c r ,
GeC1
2 08 =k 0 03
2 16 f: 0 03
Sn-C1
2.30 i: 0.03 2.39 -0.09
I
-0 04
I
1
I
C-F
~
i
Te-Cl
2.36 =k 0.03 2.36 0.00 0-F
~
~
~
1.41
1
I
-0.05 Si-F
1.54 i 0.02 81 _1_ -0,27
1 1
1.52 i: 0.04 1.74 -0.22
I
l -
000
1.41 i 0.05 1.30 +O.ll
I41
1 1
2.315 =k 0.005 2.32 0.00
I
I
i I
P-F ~
1 983 i: 0 005 1 98
i 1
~
c1-c1
I
I
~
1
1
As41
~
2 21 __ -0 13
I
1
-0 16
'
~
2 00 i 0.02 2 09 1-09
2 16 __
-
1
1
si-c1
I
1 1 6 8 i 0 0 3 1 65 $ 0 3
I
~
0-c1
1
~
I
As-F
1.72 f: 0.02 1 85
, -0.13
I I
1
although the existence of only four bond orbitals on the carbon atom prohibits its holding even one fluorine atom by a double bond and the other three by single bonds. That this decrease in distance is not due to the electrostatic attraction of the negative fluorine for the positive carbon
GTRUCTERES OF T H E FLUOROCHLOROMETHAKES
189
atom is s h o ~ r nby the result given below for methyl fluoride in which the carbon-fluorine bond has the distance 1.42 =1.U., only 0.01 A.L. larger than the normal single-bond distance. The appreciable difference in the carbon-fluorine bond distances in the two compounds, amounting to 0.06 A . U . , indicates that the carbon-fluorine bond in carbon tetrafluoride iq se1 era1 thousand calories per mole stronger than in methyl fluoride, a fact which i4 related to the differences in the cheniical properties of the substances. I undertook the investigation of the fliiorochloroniethanes in order to discover the conditions governing the behavior of the carbon-fluorine bond and its effect on other bonds in the same niolecule and also to find an explanation in terms of bond structure which would not only account for the chemical properties of the fluoroniethanes but which might be applicable to the fluorine compounds of other eleinenb. The conipounds investigated include methyl fluoride, difluoromethane, and all of the fluorochlorometlianes except trifluorochloromethane. The structures of these niolecules were determined from electron-diffraction patterns obtained from the gases. The experimental procedure has already been described (2) and the details of the interpretation of the photographs \\-ill be published separately. The results are collected in table 4 together with those previously obtained for the cliloroniethanes (11) and for carbon tetrafluoride (6, 1). The estimated probable errors indicate the relative certainty with n hich the various interatomic distances have been determined. In some of the compounds containing both chlorine and fluorine the fluorine-chlorine and chlorine-chlorine separation3 are given n i t h smaller errors than those for the shorter distances. This is due to the fact that the theoretical diffraction patterns are less sensitive to changes in the distances between the atoms whose scattering powers are smaller. I n some of the cases, as indicated below, awmiptioiis about the bond angles have been made which aid in fixing probable valueq for the smaller distances, but in every such case the ranges through which these distances may vary without causing pronounced disagreement between the theoretical and observed diffraction patterns are included n-ithin the limits of the assigned errors. The FCF bond angle in difluoroniethaiie is only slightly larger than the tetrahedral angle and is appreciably smaller than the corresponding ClCCl angle in dichloromethane. The high electronegativity of fluorine relative to hydrogen appears not to be an important factor in affecting the bond angles. Measurements on trifluoromethane would furnish an additional test of this point. The results on fluorotrichloroniethane shox that the fluorine-chlorine repulsion is less than the chlorine-chlorine repulsion; although the number of separations of each kind is the same, the bond angle for the first is 4 ' smaller than that for the second. In difluorodichloro-
190
L. 0. BROCKWAY
0 3
0
01
0
c.l
0
01
-
0 0
3
01
3
22
m01
c.l
9 0
M
9
0
-5
8
$I
si
c.l M
01 8
c.l
m m m
909
0 0 0
m m
00 0 0
zz $I+
1
STRUCTURES O F THE FLUOROCHLOROYETHANES
191
methane the chlorine-chlorine bond angle is larger than either the fluorinechlorine or fluorine-fluorine angle. These results together with comparisons with the chloroniethanes support the conclusion that the fluorinefluorine and fluorine-chlorine bond angles on carbon generally are about equal while the chlorine-chlorine angle is larger. This was assumed in the analysis of the measurements on difluorochloro- and fluorodichloromethane. It is evident that the repulsions which fix the halogen bond angles on carbon are determined chiefly by the relative sizes of the atoms and not by their relative electronegativities. The carbon-fluorine bond distances listed in the first column of the table show that the difference originally observed between methyl fluoride and tetrafluoromethane occurs between other members of this series. -1value close to 1.36 A.U. is observed for difluoro-, tetrafluoro-, and difluorodichloromethane. DifluorocJdoromethane also belongs to this group, for although the determination is less accurate the carbon-fluorine distance in this compound cannot be more than 0.02 or 0.03 ATJ. greater than 1.36 A.V. A value of about 1.41 A.U., on the other hand, is observed for methyl fluoride, fluorochloromethane and fluorodichloromethane. Only in the case of fluorotrichloromethane is the uncertainty in the determination as large as the difference between the two values for the carbon-fluorine distance; here the existing evidence supports the larger value and this compound, too, probably belongs to the second group. The striking feature in this classification is that the methane derivatives containing two or more fluorine atoms fall into the first and those containing only one fall into the second group. The value 1.41 A.U. is the sum of the single covalent radii for carbon and fluorine. Since it is observed in methyl fluoride, and the methyl compounds in general ( 5 ) have the normal values for single covalent bonds, it is probable that the carbon-fluorine bonds in the second group of compounds represent the normal type of such bonds. The bonds in the first group of compounds are then about 4 per cent shorter than the normal value. This shortening interpreted in terms of bond energies indicates that a carbon-fluorine bond in the first group is about 8000 or 10,000 calories per mole stronger than one in the second group. The carbon-chlorine bond distances in the second column of table 4 show that in most of these compounds the bond has the normal distance, 1.76 A.U., observed in the chloromethanes. In particular the chlorine bonds in fluorochloroinethane and fluorotrichloroniethane are not shortened. I t is to be noted, however, that there is a possible effect in the two compounds having two fluorine atonis on the carbon. The shortening is small and uncertain in any event, and we may conclude that if there is any strengthening of the carbon-chlorine bond it is much less pronounced than the strengthening observed in some cases in the carbon-fluorine bond.
192
L. 0. BROCKWAP
The chemical properties of the fluorochloroniethanes and fluorochloroethanes have been reported by Nidgley and Henne. They have studied the influence of fluorine atoms in the molecule on the reactions in which chlorine (or bromine) is substituted by fluorine. The following quotation from their recent survey (7) s h o m the very close relation between the structural anomalies observed above and the chemical behavior of these compounds: ‘‘ . . . The presence of a lone fluorine atom in an aliphatic molecule produces an unstable compound, which shows a great tendency to liberate hydrogen fluoride. In contradistinction, compounds having a halogen atom on the carbon bearing the fluorine are more stable, The presence of two fluorine atoms on the same carbon is attended with (1) a very great increase of the strength of the carbon-fluorine bonds; (2) a strengthening of the linking of any third halogen atom that may be present on this same carbon; and (3) a great increase ofJhe stability of any halogen atoms present on adjacent carbon atoms.” The structural data are in accord with these conclusions. The shortening and strengthening of the carhon-fluorine bond are found to occur in the methane derivatives having two or more fluorine atoms on the carbon atom. Accepting the chemical evidence for the effect on chlorine as a third halogen atom, n-e know from the observed bond distances that the strengthening of the carbon-chlorine bond is considerably less than that of the fluorine bond. Among fluoromethanes and fluorochloroniethanes the very great increase in stability occurs only in these having two or more fluorine atoms, the differences between methyl fluoride and the monofluorochloromethanes being much less. An explanation of the structural and chemical properties of the fluorine derivatives of aliphatic compounds can be developed from the assumption that on account of its great electronegativity fluorine has a tendency to take the shared electron pair from the carbon atom and become a negative ion. When this ionic character of the bond is high the fluorine atom is reactive; a decrease in the ionic character will make the fluorine atom less reactive. The saturated monofluorides with the exception of the first four in the series of normal alkyls lose hydrogen fluoride spontaneously. According to the foregoing postulate the ionic character of the fluorine bond in these compounds is so pronounced that the separation into a negative fluoride ion and a positive alkyl ion actually occurs; the latter then loses a proton and is transformed into an olefin. All of the saturated monofluorides are hydrolyzed quite easily by concentrated acids and bases. I n basic solutions the high concentration of hydroxyl ions favors the formation of the carbon-oxygen bond, which has less ionic character than the fluorine bond, and the alcohol is formed, The splitting off of hydrogen fluoride observed in acid solutions may be attributed t o an increase in the tendency of the
STRVCTURES OF T H E FL~OROCHLOROMETHAKES
193
fluoride ion toward a complete heparation from the alkyl group due to the high concentration of hydrogen ions and the alkyl ion again becomes an olefin molecule. The ionic character of the carbon-fluorine bond is decreased when another group is present in the niolecule n hich tends to reduce the relative electronegativity of the carbon and fluorine atoms. This will occur in general when one of the hydrogens is substituted by a more negative group, and under these circumstances the fluorine compound should be more stable. This increase of stability is observed in the fluorine compounds containing other halogens. The ionic character and reactivity of the carbon-fluorine bond may be decreased by the presence of a negative atom on the adjacent carbon as shown by the stability of monofluoroethanol and monofluoroacetic acid. It will be noted that in the fluorochloro compounds both halogen a t o m are less reactive than in the respective mono derivativeq. This mutual effect is probably due to the increase in the negativity of the carbon atom occurring on the substitution of either halogen into the molecule containing an atom of the other so that both carbon-halogen bonds have less ionic character and reactivity than in their respective nionohalogen parents. The very marked increase in stability when the second substituent is fluorine we attribute to an additional effect. If the normal state of difluoromethane is represented in terms of the individual electronic structures, whose resonance together determines the properties of the substance, the three structures I, 11, and 111 would be expected by analogy to methyl fluoride to contribute to the normal state of the molecule, I1 and I11 giving ionic character to the carbon-fluorine bonds.
..
:F:
H-C-F: I
..-
:F: ..
..
H-Cf-F
..
I
..
:F:
.. .. :
1
H-C+
H I1
H I
..-
:F: *.
H I11
I n addition, however, the two structures IV and 1-have low energy values and can make a significant contribution.
..-
:F: .. H-C=F H IV
**
:F I
..
H-C
,
+
II
H 1-
:Fr *.
194
L. 0. BROCKWAY
In these structures one fluorine atom is able to form a double covalent bond by sharing one of its extra pairs with carbon, on n-hich a bond orbital has been released by the other fluorine atom n-hich has assumed an ionic form. These structures making equal contributions to the normal state would strengthen and shorten the carbon-fluorine bonds. With the aid of the relation between bond distance and double-bond character (9) the observed shortening can be accounted for by a 10 per cent contribution from each of IV and T. This formulation is applicable only to the molecules having two or more fluorines on one carbon, a distinction which agrees with the experimental evidence. Double-bond character of the bonds between halogen atoms and nonfirst-row atoms has already been well established; for these heavier atoms, which have a larger number of orbitals in the valence shell, there is no such restriction on double-bond formation as there is for carbon. The shortening effect of double-bond character is not found in the monofluorochloromethanes presumably because chlorine has not the great electronegativity which would allow the assumption of an extreme ionic form in these compounds. ' I n fluoroethylene, also, the greater stability may be explained by the contribution of some double-bond character to the carbon-fluorine bond. Here as in the chloroethylenes (4) one of the important structures has a single bond between the carbon atoms and double bond to the halogen. The representation for fluoroethylene is the following :
/
\
H I
I1
H
I11
Measurements should show that the carbon-fluorine bond in this molecule is shortened without the presence of a second fluorine atom. When other halogen atoms are present the same effect n-ould occur, although more structures would contribute to the normal state. We can predict that there is not the great difference in the stability of the mono- and di-fluoroethylenes that is observed in the corresponding saturated compounds. The decreased reactivity of chlorine in the fluorochloromethanes is apparently due chiefly to the effect of the fluorine in decreasing the ionic character of the carbon-chlorine bond. I n the difluorochloromethanes there is some evidence of double-bond character in the chlorine bond, but to a much lesser extent than in the fluorine bonds. The stabilizing effect of two fluorine< on chlorine attached to the adjacent carbon atom is ac-
STRUCTURES OF THE FLUOROCHLOROMETHANES
195
counted for by the increased negativity of the carbon and the leqser ionic character of the chlorine bond. The foregoing discussion has assumed that clieniical reactivity is dependent upon (1) the degree of ionic character in bonds which are largely covalent and ( 2 ) the contribution of double-bond character to single bonds. From the chemical data it appears that the second effect has the more pronounced influence on reactivity. From the structural data it appears that the second effect has much the greater influence on the bond distances. Indeed the change in distance with ionic character must be very small. This conclusion is supported by the observed carbon-nitrogen bond distances in cyanides and isocyanides (3) which have the same value in spite of the difference in the ionic character of the bonds. That great differences in chemical reactivity may follow small changes in structure has been discussed theoretically by Sutton and Pauling (12). The explanation offered for the rnethanes does not solve the question of the abnormal shortening in the phosphorus and silicon fluorides. The 3 to 5 per cent shortenings observed below the double-bond values for these compounds have not yet been explained. The contribution of triple-bond structures is not very plausible, in view of the sniall tendency which second row elements have toward triple-bond formation. It is evident from the foregoing results, however, that the mixed halides of phosphorus and silicon should be investigated for variations in the fluorine bond distances. The kindness of Dr. llidgley and Dr. H e m e of the Jlidgley Foundation and of l I r , R. J. Thompson of Kinetic Chemicals, Inc., in supplying the materials is gratefully acknowledged. To Professor Linus Pauling I am indebted for the benefit of many consultations and suggestions. REFERENCES (1) BROCKWAY, L. 0 . : J . Am. Chem. Soc. 67, 958 (1935). (2) BROCKWAP, L. 0.: Rev. Modern Phys. 8, 260 (1936). (3) BROCKWAY, L. 0.: J. Am. Chem. Sor. 68, December (1936). (4) B ~ o c ~ w . a YL., BEACH, J. I*., ASD PACLISG, L.:J. .4m. Chem. soc. 67,2693, 2705 (1935). ( 5 ) BROCKWAY, L. O., A N D JE~YKINS, H. 0.: J . Am. Chem. Soc. 68,2036 (1936). (6) BROCKWAY, L. O., A N D JVALL,F. T.: J. Am. Chem. SOC.66, 2373 (1934). (7) HESVTE, .4.L., A N D ~ I I D G L E T., YJR.: , J . Am. Chem. Soc. 68,882 (1936). (8) PAULIXG, L.: Proc. Natl. Acad. Sei. 18, 293 (1932). (9) PAULISG, L., BROCKWAY, L. O., A N D BEACH,J. T.: J . .4m. Chem. SOC.67, 2705 (1935). (10) PAULISG, I,., ASD HUGGISS, XI. Id.: Z. Icrist. 87,205 (1934). (11) SUTTOS, L. E., .ISD BROCKWAY, L. 0.: J. h i . Chem. SOC.67,473 (1935). (12) SUTTOS,L. E., ASD P ~ U L I S G L .,: Trans. Faraday SOC.171,939 (1935).
o.,
