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Force Constants and Bond Orders of Nitrogen Bonds.

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Force Constants and Bond Orders of Nitrogen Bonds
BY PROF. DR. J. GOUBEAU
LABORATORIUM FUR AWORGANISCHE CHEMIE
DER TECHNISCHEN HOCHSCHULE STUTTGART (GERMANY)
The force constants and the corresponding bond orders of nitrogen bonds have been calculated from the vibrational spectra (infrared and Raman spectra) of a great number of
nitrogen compounds. Plotting the maximum bond order of stable nitrogen bonds against
the sum of Pauling’s electronegativities of the bonding partners ( x x ) leads to one continuows
curve for the N - X bonds where X represents elements of the first and the second short
period of the periodic table. Furthermore, when the bonds formed between these elements
are arranged in a coordinate system in such a way that the position of each bond is determined by the diference between the electronegativities of the bonding partners ( A x along
the ordinate) and the sum of the electronegativities of the bonding partners ( r x along the
abscissa), the bonding partners capable of forming multiple bonds all lie within a closed
domain, where their position can be correlated with their polymerizability and other reactivities of the multiple bonds. Also discussed are the orders of bonds between nitrogen and
some transition elements. In an appendix, the present methods used to calculate force
constants and bond orders are surveyed.
A. Introduction
cumulation of high formal charges. The types of bonds
thaLnitrogen can form are summarized in Table 1.
The conditions for the formation of x-bonds have been
previously summarized as follows [ I ] :
1) electron deficiency on both bonding partners;
2) the sum of the Pauling electronegativities of the
partners must be at least 5;
3) the difference between the Pauling electronegativities
of the bonding partners must be small.
We have endeavored in the past few years to verify these
conditions on further examples and, in particular, to
define more accurately condition 3). Condition 2) has
recently been confirmed by Herber and Stockler [*I, who
investigated the Mossbauer effect of tin compounds and
showed that the isomeric shift is invariably proportional
to the sum of the electronegativities of all bonding
partners involved.
We have chosen nitrogen for our investigations, since
it forms a great number of compounds and has the high
electronegativity of 3,0, enabling it to form bonds with
numerous elements in such a way that the sum of the
electronegativities is at least 5 . Unlike the halogens, and
like oxygen, nitrogen can form x-bonds without the acTable I . Types of bond f o r m e 3 between nitrogon a n 3 a partnor X .
Imines, imides
1
e e
E-E-H
tt X = N H tt X E N H
In addition to the preferred bonds, in which the two atoms
involved have a low formal charge (cf. Table l), bonds in
which the nitrogen carries a formal positive charge are also
important. This formal positive charge is to some extent
compensated by the fact that nitrogen is at the same time the
negative end of the polar bonds which it forms in most cases.
B. Calculation of Bond Order
The following discussion is based on vibrational spectra
(infrared and Raman spectra) furnishing data for the
force constants, which in turn are used for the calculation of the bond orders. In these calculations, the geometry of the molecule and the assignment of the observed frequencies to the fundamental vibrations is also
accounted for. Fadini, Sawodny, and Ballein [31 have
used a new method of calculation, enabling them to obtain the force constants and all interaction constants
directly from the measured frequencies without a trialand-error procedure 141. The resulting valence bond
constants, the only ones discussed below, have the
advantage of being mutually comparable, since they
are obatined by the same method. Deviations of more
than 5-10 % correspond to actual changes in the nature
of the bonding.
Siebert’s rule r4.51 has been used for the calculation of
the force constants of single bonds. This rule gives good
results, suitable for a general comparison of all force
constants. Siebert’s values for single bonds correspond
.
Nitrides
[I] J . Goubeau, Angew. Chem. 69, 77 (1957).
[ 2 ] R . H. Herber and A. Stockler, Chem. Engng. News 42, No.
28, 66 (1964).
Angew. Chem. internat. Edit.
VoI. 5 (1966) 1 No. 6
.
[3] A. Fudini, Z . angew. Math. Mechan. 44, 506 (1964); W. Sawodny, A. Fudini, and K . Ballein, Spectrochim. Acta 21, 995
(1965).
[4] J. Goubeuu, Angew. Chem. 7 3 , 305 (1961).
[5] H. Siebert, Z. anorg. allg. Chem. 273, 170 (1953).
567
to sp3 hybridization. Since differences in the hybridization can lead to large (up to 50%) variations in the
force constants of single bonds, it has been found necessary to introduce an empirical correction factor 111.
Using these theoretical force constants for single bonds,
we can calculate the bond order, b. It is assumed that
higher force constants are essentially due to x-bonds.
Siebert has given two equations for the bond order.
One is based on the assumption that the bond is proportional to the force constant up to about b = 1.5,
the other being an empirical formula (giving good
results) for b > 1.5 (see Appendix 111). The calculations
are generally based on direct proportionality even in the
case of higher bond orders. It is further possible to correlate, on a theoretical basis, the bond orders of single
and double bonds in certain compounds with the force
constants, and then to find the bond orders corresponding to the other force constants, as has been done in
Tables 3 and 5 (column btr).
and, therefore, it can form multiple bonds (at the most a
three-electron bond in NF2). Distributed over two N-F
bonds, this leads to a maximum bond order of 1.25,
while in N F the bond order may reach 2.0.
fexp. ImdyneiAI
fcalc. (single bond) lmdyne/Al
4.38
4.26
4.85
4.05 [el
b
1.03
1.20
5.90 171
3.87 [81
4.05 IS1
1.45 P a l
C . Order of Bonds between Nitrogen and Elements
of the First Period
11. The N-B
T h e parameters of the N-0, N-N, a n d t h e N-C bonds a r e
well known a n d need no further discussion. In all three cases,
the b o n d order can b e between 1 a n d 3 according to the force
constants.
Bond
The most important results concerning force constants
and bond orders of N-B bonds in various compounds
are listed in Table 3 191, together with the possible
Table 3. Force constants f and bond orders b of N-B bonds (X
H, CH3, F, CI; Y
i
=
H. CH,)
Formula
X~BNYJ
I. The N-F
Canonical forms
-
tl
b1 I
-2.8
0.7
I .0
4.0
I .0
1.3
6.0
1.5
1.7
6.3
1.6
I .8
RzBNCO
6.6
1.7
1.9
XzBNYz
7.c
1.8
2.0
BN;‘]
7.0
1.8
2.0
BN
8.3
2. I
2.3
Bond
Although, on the basis of the sum of electronegativities,
the formation of multiple bonds should be possible in
nitrogen trifluoride, this compound possesses no electron deficiency (condition 1 in Section A), and can,
therefore, form no x-bonds. In the compounds NF2 161
and N F 171, which have recently been described together
with their vibrational spectra, the nitrogen does not
attain an electronic octet (i.e. it is electron deficient),
[6] F. A . Johnson and C . B. Colbuin, Inorg. Chem. I , 431 (1962);
M . D . Harmony and R . J. Myers, J . chem. Physics 37,636 (1962).
[7] M. E. Jacox and D . E. Milligan, J . chem. Physics 40, 2451
( 1964).
mesomeric limiting structures. The bond orders bI are
based on Siebert’s value for the single bond (f = 3.9
mdyne/&. In calculating thebond order bII, it was assumed that compounds of the type of X3BNY3 are
characterized by a single bond, while BN23- contains a
[8] In such a case, the value for the a-bond represents an intermediate value, since the number of free electron pairs changes
in the transition from single bond to double bond. For N F two
values are obtained (3.87 and 4.05).
[Sa] The bond orders were determined by a simple averaging
method, the extreme values being b = 1.39 and b = 1.52.
[9] Cf. J. Goubeau, Advances Chem. Ser. 42, 87 (1964).
H . Noth, G. Schmid, and Y . Chung (IUPAC Congress, Vienna,
1964) have found another example of a B-N bond with high
bond order in the adduct B(NR2)3,TiC14. However, since the
force constant has not been calculated, this example could not be
included in Table 3.
Angew. Chem. internat. Edit. 1 Vol. 5 (1966) No. 6
double bond. The intermediate bond orders were found
on the basis of an assumed proportionality between the
bond orders and the force constants. The difference between bI and b,, (about 0.2) is of little significance for
the subsequent discussion; the bI values are preferred
for purposes of comparison. The maximum bond order
has thus been found to be 1.8 for the N-B bond in compounds that are stable at room temperature.
111. The N-Be
D. Order of Bonds of Nitrogen with Elements of
the Second Short Period
Bond
Beryllium compounds, like boron compounds, form
adducts with ammonia and amines [1*1. The reaction between beryllium alkyls and amines leads to substituted
beryllium amides 1191 which could not be prepared from
adducts of the type BeC12.2HN(CH3)3 by the abstraction of HC11201. Unlike the boron compounds, these
beryllium amides are known only in polymeric form.
The force constants have been calculated from the
spectra of berylliumdi-[bis-(trimethylsilyl)]amide [I31
and the adducts formed of ammonia, dimethylamine,
and trimethylamine with BeC12 and BeBr2[21]. In the
first case, the value f = 2.96 mdyne/A has been obtained
for the N-Be bond, while in the other cases, calculations on the basis of the antisymmetric vibrations of the
BeN2 group (600-785 cm-1) led to f = 1.2 to 1.4
mdyne/h;. Using these values and taking, with Siebert,
f = 3.1 mdyne/A for the single bond, we obtain b =
0.95 and b = 0.39 to 0.45, respectively. Since the beryllium amides are polymeric, we cannot expect them to
have appreciably higher force constants. In fact, there
is no indication at the present of higher bond orders in
N-Be compounds.
IV. The N-1.i
Bond
Only one very broad infrared band has been observed
for each of the simple compounds available, LiNH2,
LiNH, and Li3N. These bands are situated, respectively,
[lo]J . R. L. Anister and R . C . Taylor, Spectrochim. Acta 20, 1487
(1964);J . Goubeau and W . Sawodns, Z. physik. Chem. N.F. 4 4 ,
227 (1965).
[Ill J . Goubeau a r d H . K c l l u , Z . a n o r g . allg. Chem. 272, 303
(1953).
[I21 H.-J. Eecher, Z.anorg. allg. Chem. 287, 283 (1956).
[I31 B. L. Crawford and J . T . Edsall, J. chem. Physics 7 , 223
(1939);H . Watanabe, M . Norisada, T . Nukagawn, and M . Kubo,
Spectrochim. Acta 16, 78 (1960).
(141 J. Goubeair and H. Grabner, Che:n. Bir. 93, 1379 (1960).
[IS] J . Goubeaii and H.-J. Eecher, Z . anorg. allg. Chem. 268, I33
(1952).
[I61 J . Goiibeau and W. Anselment, Z. anorg. allg. Chem. 310,
248 (1961).
[I71 A . E. Douglas and G. Herzberg, Canad. J. Res., Sect. A 18,
I79 ( 1 940).
118Je.g.R.FrirkeandF.RObk~,Z.anorg.allg.Chem.170,25(1928).
[I91 See, e.g., G . E. Coates and F. Clocking, J. chem. SOC.(London) 1954, 22.
[20] N. Th. Rakintzis, Dissertation, Technische Hochschule
Stuttgart, 1957.
[211 Ch. Forker, Dissertation, Technische Hochschule Stuttgart,
1964; H . Burger, Ch. Forker, and J . Goitbeau, Mh. Chem. 96,
597 (1965).
Aiigew. Chem. intertiat. Edit.
at 403, 465, and 430cm-1 and can be attributed to
antisymmetric vibrations [221. The resulting force constants (0.44,0.61, and 0.55 mdyne/& respectively) are
probably too high. On the assumption that the single
bond has a force constant of 2.3 mdyne/A, the average
bond order for the N-Li bond is 0.23.
1 Vol. 5
(1966)
/ No. 6
I. The N-Cl
Bond
Owing to the presence of vacant d-orbitals in elements
of the second short period, electron deficiency exists here
even when the electronic octet is reached. This means
that double bonds can be formed, e.g. in NC13, unlike
the case of NF3 :
.. ..
,N-CI:
>NZ'CI:
H
The experimental data for the N-C1 bonds are listed
in Table 4.
Table 4. Force constants and bond orders of N- CI bonds in various
compounds I*].
HzNCl
HNClz
NCI,
N C1
NC10323
3.1
3.0
2.5-3.0
4.0
8.9
1 .&
1.2,
1.04-1.21
1.71
2.2
1231
[231
123,241
[71
[251
In nitrogen trichloride, the calculated force constant
varies with the C1-N-C1 angle; the lower value corresponds to a tetrahedral angle, the higher to an angle
somewhat smaller than 120 '. However, the geometry
of the molecule is not yet known. The bond orders in
chloramine and dichloramine confirm the prediction
based on the presence of vacant d-xbitals. These are
also partly responsible for the fact that chlorine is the
negative end in the N-C1 dipole 1261. The bond order in
NC1 is seen to be high; it is higher than in NF, probably
because of the smaller difference between the electronegaiivities of the bonding partners. The highes&force
constant (f = 8.9 mdyne /A) is exhibited by NC10:--, this
value being greater than that corresponding to a double
bond.
11. The N-S
Bond
Out of the reliable data for the force constants of N-S
bonds, some characteristic examples are listed in
Table 5 .
[22]V. Dorn, Dissertation,Technische Hochschule Stuttgart, 1964.
[ * ] According to Siebert, the sp3 single bond has a force constant
of 4.0 mdyne/>f.
[23] G. M . Moore and R. Badger, J . Amer. chem. Soc. 7 4 , 6076
(1952).
[24] E . Allenstein, Dissertation, Technische Hochschule Stuttgart, 1953.
[25] E. Kilcioglu, Diploma Thesis, Technische Hochschule Stuttgart, 1965;together with E. Jacob.
[26] E. Allenstein, Z.anorg. allg. Chem. 308, 3 (1961).
569
Table 5. Force constants and bond order of X-S
compounds.
bonds in various
Ref.
3.0
3.1
3.1
4.3
8. I
8.3
10.7
12.5
0.8
0.8
0.8
0.8
1.1
1.3
1.9
2.C
3.0
3.0
1.2
1.5
2.5
2.6
3.4
3.4
IV. The N-Si
Bond
Only a few examples have been selected from the chemistry of N-P compounds[331 to illustrate the range of
the force constants. Discounting the band spectrum of
PN in the gaseous state, which would probably give b 3, we find that the bond order for normal compounds
varies between 0.8 and 2.0.
The force constants in [C13P-N-PC13]+ [*I depend on
the P-N-P angle (which is again unknown):
Angle
f[mdyne/A]
I
I
120
I
62
I 6 9
140”
1
160
I 7 8
1
f [mdyne/A]
3.1
3.5
4.1
4.3
6.1
6.9
10.6
1 9 1
111
variois
b
0.8
I .3
I
1
1.7
2.0
3.4
.z
.z
O n t h e basis of calculations 1331, t h e most likely value for t h e
angle is 140OC; it is associated with the highest force con[27] F. Watari, Z . anorg. allg. Chem. 332, 322 (1964).
[28] H . Siebert, Z . anorg. allg. Chem. 292, 167 (1957); J . Nakagawa, S . Mizushirna, A . J . Saraceno, T . J . Lave, and J . V . Quaglinno, Spectrochim. Acta 12, 239 (1958) (calculated by present
author).
I291 W. Bubeck, Dissertation, Technische Hochschule Stuttgart,
1962.
[30] B. v. Cramon, Staatsexamen Thesis, Technische Hochschule
Stuttgart, 1958; F. Rauchle, Diploma Thesis, Technischc Hochschule Stuttgart, 1962.
[31] H . Richerr and 0 . Gleinser, Angew. Chem. 72, 585 (1960).
[32] H . Richert and 0 . Glemser, Z. anorg. allg. Chem. 307, 328
(1961).
[33] R . Boumgartner, CV. Saivoch7y, and J . Gorrbeau, 2. anorg. allg.
Chem. 340, 246 ( 1 965).
[ * ] The author thanks Frau Prof. Dr. M . Brcke-Goeliring [34]
(Heidelberg) for these data.
570
Table 7. Force corist2nls and b o n d orders of N-SI bonds In various
compounds.
I
f [mdyne/*,]
Single bond
HN[Si(CH,)312; 140
HN(SiCl3)?; 140’
WSiH3).1
N[Si(CH3)3]>’3:’;160’
3.3
3.5
4.0
4.1
4.2
>Us;
7.3
I
b
I
Ref.
1.1
I .2
1.2
1.3
2.2
Discounting molecular SiN which exist only at high
temperatures, we find that the maximum bond order is
1.3 for the N-Si bond in compounds stable at ordinary
temperatures. The maximum value of f = 4.8 mdyne/A
reported for the compound S=C=N-SiH3 [451 is not included in Table 7, since, on recalculation, we have obtained a considerably lower value. However, even for
the published value the bond order would be only 1.45.
In the case of silicon, compounds of the type >Si=NR,
which are monomeric in the case of sulfur and phosphorus, exist here only in polymeric form. Although we
have taken b = 1.3 as the maximum value, the bond
order may be somewhat higher in certain compounds.
V. The N-A1
180
Table 6 Force constants and bond orders of N - P bonds
compounds [**I
Bond
The reported values of the force constants vary,
probably because the bond angle exerts a strong influence on the value off.
With Siebert’s value of f = 3.7 mdyne/A for the single
bond, we obtain bond orders between 0.8 and 3.4,
taking into account the hybridization and assuming a
direct proportionality between f and b. However, since
the bond order cannot exceed 3, it is better to ascribe
the values b = 3 and b = 2 to NSF3 and HNSO, respectively, an3 to recalculate the other values on this basis
(brl). The corrected Siebert formula gives b = 2.70 and
b = 2.2 in these cases. The properties of these compounds
confirm the presence of the double and triple bonds
deduced from the bond orders.
111. The N-P
stant known t o us (f = 6.9). I t should b e added t h a t compounds such a s R N P C 1 3 , R N P F 3 , NPC12, a n d NPF2, for
which high b o n d orders are expected, exist only in polymeric
form a t r o o m temperature.
Bond
In this case and in the following ones, spectroscopic data
that may be used for the calculations are relatively
scarce, except perhaps for aluminum nitride [*21 and
the ammonia and amine adducts of aluminum com[**I Sieberr’s value for the sp3 single bond is f = 3.5.
[34] M. Becke-Goehring and W. Lehr, Z . anorg. allg. Chem. 327,
128 (1964).
(351 E. Steger, Z . anorg. allg. Chem. 309, 304 (1961).
[36] N . Lonhof, Dissertation, Technische Hochschule Stuttgart,
1964.
[37] J . v. Iriborne and D . G . de Kowalewski, J. chem. Physics 20,
346 (1952).
[38] Calculated by the author from data communicated by R .
Appel.
[39] J. Curry, L. Herzbcrg, and G. Herzberg, Z . Physik 86, 348
(1933).
[40] U. a. H. Kriegsnzanri, Z . Elektrochem., Ber. Bunsenges.
physik. Chem. 61, 1088 (1957).
[41] G. Gndmundsson, Dissertation, Technische Hochschule
Stuttgart, 1964.
[42] E. A . V. Ebsworth, J . R. Holl, M . J . Mackilon, D . C.
McKean, N. Sheppnrd, and L . A . Woodward, Spectrochim. Acta
13, 202 (1958); H . Kriegsmann and W. Fiirsfer, Z . anorg. allg.
Chcm. 2Y8, 212 (1959).
[43] 13. Biirger, W . Sriwodny, and I/. Wannagat, unpublished
work.
[44] F. A . Jenkins and H. rle Laszlo, Proc. Roy. SOC.(London),
Ser. A 122, 103 ( I 929).
[45] K. Sothianandon and J . C . Margrave, J . molecular Spectroscopy 10, 442 (1963).
Angew. Chem. internat. Edit. 1 Vol. 5 (1966)
/ No. 6
pounds [ l o , 461. The calculated force constants are between 1.9 and 2.5 mdyne/& to which correspond bond
orders between 0.6 and 0.8 (Siebert’s force constant of
the single bond is f = 3.0).
VI. The N-Mg
and the N-Na
Bond
Infrared spectra have been recorded for Mg3N2,
Mg(NH2)2, and NaNH2 “221, and the force constants for
the MgN (1.33 and 2.00) and NaN bonds (0.53) calculated only from antisymmetric vibrations on the basis
of the two-body model now appear to represent upper
limits. The resulting bond orders, b = 0.47 and 0.70
for MgN and b = 0.21 for NaN, are compatible with
the high polarity of these bonds.
Basically the same curve is obtained by plotting b
against Ex for the nitrogen bonds formed with elements
of the first short period and elements of the second short
period (Fig. 1 ; x x denotes the sum of the electronegativities of the bonding partners, and the b values
represent the maximum bond orders). Whenever S x >
5, double bonds can be formed (except in the case of F).
Recalling condition 2 which states that, for the formation of double bonds, the sum of the electronegativities
of the bonding partners must be at least 5, we note that,
when this condition is not fulfilled (i.e. for the nitrogen
bonds to Li, Be, B and Na, Mg, Al, and Si) the bond
orders do in fact lie below b - 2 (b = 1 at x x m 4.5).
Bond orders below 1 are due to the fact that the
polarity of the bonds increases with decreasing value of
>:x [31. The bond orders of 0.42 and 0.72 for the “single
bond” adducts of Be and R, respectively, fit in well with
this series. The coincidence of the curves for the two
short periods may be regarded as a confirmationiof the
second condition stated in Section A.
dimer (probably a four-membered ring) is stable only at
temperatures below -10 “C 1471, further polymerization occurring at higher temperatures. The N=Se bond
is, therefore, less stable than the analogous N=S band.
Approximate calculations based on the assumption that
the dimer forms a four-membered ring have given f =
3.4 and b = 1.3 for the N-Se bond.
In addition to insufficient experimental data on the
nitrogen bonds of the transition elements, calculations
are made difficult by the fact that the electronegativities
are not accurately known and the electrons can occupy
various orbitals. The force constants and the bond
orders of the nitrogen bonds of the transition elements
are listed in Table 8.
None of the elements has been subjected to exhaustive
investigations concerning their bond parameters. Instead,
ranges of single values are given in Table 8 for the force
constants of azides, since the N3 group was treated as a
point mass in the simplified approximation. The lower
limit corresponds to mass 14, the upper to mass cu. 30.
T a b l e 8. Electronegativities (x). stretching frequencies (v), force constants (f), a n d bond
orders (b) of the nitrogen bonds of transition elements
Ref.
Y
f acc. t o
Siehert
v [cm-’1
I .6
406
565
3.2
I .8
555
3.7
560
615
Ill0
1.9
475
470
530
715
815
3.6
1073
2.0
552
2.1
495
565
780
3.8
1045
Na Mo
Ai
P
SI
S
CI
2.1
2.2
0
LI
Be
B
C
N
O
F
E. Order of Bonds between Nitrogen a n d Elements
of Higher Periods Including Transition Elements
The bond order for NBr is 1.7[71, and the bond is,
therefore, similar to that in NCl. The compound
RN=Se=O does not exist as a monomer and even its
[46] H . Roszitiski, R . Dautel, and W. Zeil, 2. physik. Chem. N.F.
36, 26 (1961).
Aiigew. Chem. Iiiternaf. Edit.
/ Vol. 5 (1966) / No. 6
4.1
3.8
1025
1023
1084
3.9
fex,.
b
1.1
2.0-3.1
2.0-3.1
2.0-3.1
0.3
0.6- I .O
0.5-0.8
2.5
0.7
2.2
8.0
I .8
2.1
0.5-0.8
0.5
0.6
4.8
1.3
8.8
2.2-3.1
7.9
1.6-2.5
2.1
4.9
8.0
8.0
9.0
2.4
0.6- I .O
2.2
0.4-0.6
0.5
1.3
2.1
The higher values, which are more probable, lead to
bond orders of 0.6 and 1.0 in accordance with expectations. When the values of b,,
for Ti, V, Mo, W, Re
and 0 s are plotted against c x (Fig. 2), the points for
Ti, Re, and 0 s lie on the curve, while Mo, and particularly W and V (the last two in their chloronitrides with
b > 2) lie above the curve. The 0 - V bond in oxygencontaining compounds has a similar anomalous position.
[47] J . Goubenu and U . Weser, Z . anorg. allg. Chem. 319, 276
(1963).
[48] K . Dehnicke, Habilitation Thesis, Technische Mochschule
Stuttgart, 1964.
[49] J . Strahle, Dissertation, Technische Hochschule Stuttgart,
1964.
[50] M . Schober, Dissertation, Technische Hochschule Stuttgart,
1964.
[51J J . Clioffand G. A . Rowe, J. chern. SOC.(London) 1962,4019.
[52] G. Wilkiiison, Inorganic Syntheses VI, 169.
[53] J . Lewis and G . Wilkinson, J. inorg. nuclear Chem. 6, 12
(1958); L . A . Woodwardand J. A . Creighron, Trans. Faraday SOC.
56, 1266 (1960); U. Hiibler, Staatsexamen Thesis, Technische
Hochschule Stuttgart, 1962.
57 1
2.1
2.4
Most element pairs lying within the boundaries drawn
for the first short period (shaded area) can similarly reach
a bond order of 2, the exceptions being the chlorine bonds
of P, S, and C [owing the polar bonding (halogen effect)] and the S-S bond. The latter however, is clearly a
limiting case, since S=S bonds do exist in S2 mdecules
at 800°C[541. These anomalies indicate that, as expected, Puuling's electronegativities represent only a
rough approximation.
0L 5" " " "47' "
L
19
-
53
51
zx
m
55
20
Fig. 2. Maximum bond order hmax of the nitrogen bonds OF transition
and other elements as a function of the suni of Pading's electronegativities of the bonding partners (Xx).
15
It cannot yet be decided whether the electronegatitivies [*I
and/or force constants belonging to the single bonds are incorrect, or whether the transition metal bonds are in fact
different. A more accurate knowledge of the hybridization
may throw some light on this problem.
110
Q
05
F. Comparison with Bonds Involving n o Nitrogen
0
In order to compare the polarities of all possible bonds
that may be formed within the first short period, we have
constructed Fig. 3 by plotting the sum of the electronegativities against the difference of electronegativities
of all such combinations. The three apexes of the triangle represent the three types of bond in pure form.
rn
15
50
55
E5
70
Fig. 4. Diagram f o r element pairs from the first and second short period,
obtained by plotting S x against Ax. The shaded area comprises bonds
with b 2 2 in compounds stable at ordinary temperatures (curve I :
b = 1; curve 11: h % 1 . 5 ; curve 111: b R+ 3).
G . Conclusions
Although it may seem like playing with numbers, the
above treatment does lead to significant conclusions
concerning the stability and the reactivity of multiple
bonds. Thus, we can quote again the following examples
(cf. Fig. 1):
I
x
Q
3 :NrN:
10
3 HCEN:
t
\ '
N6
C X =6.0
ti3C3N3
cu= 5.5
H3B3N3H3
Z X =5.0
-0-
3 HB-NH
0
-20
30
50
40
60
ZX
.65183,
70
80
Fig. 3. Diagram for element pairs from the first short period, obtained
by plotting the sum of the electronegativities (Ex) against the difference
of electronegativities (Ax) of the bonding partners. The shaded area
2 in compounds stable at ordinary tempercomprises bonds with b
a tures.
>
The shaded area contains the element-pairs that can
form multiple bonds (b
2.0) in compounds stable at
ordinary temperatures. The boundary to the top right
is determined by the electron deficiency (condition l),
that to the top left by the difference, and to the left by
the sum of the electronegativities (conditions 3 and 2,
respectively) .
Fig. 4 has been obtained in the same way for element
pairs formed in the first and second short periods, the
range being now set by 2lx = 4.5 and 7.0 and by Ax =
0 and 2.0 comprising the more interesting combinations.
[*I The literature values of tlizse electronegativities differ by 0.1
to 0.2 eV. We have taken the mean values, but the highest ones
would bring V, W, and M o nearer to the curve.
572
->
In the case of nitrogen with >:x = 6.0, the triple bond is
so stable that no polymer is formed. In the case of HCN
with r x = 5.5, both monomeric and polymeric forms
exist, while only the polymer is known for HBNH
characterized by x x = 5.0. In terms of the diagram in
Fig. 4, the lower right element pairs form very stable
multiple bonds with a small tendency to polymerize. The
stability of the multiple bonds decreases, and the polymerizability increases, as one proceedes towards the top
left, coming to bonds formed by carbon, phosphorus,
and boron frequently encountered in polymer chemistry. The element pairs N - 0 , C-0, and B-0 form
similar rows, as do C1-0, S - 0 , and P-0.
The N - 0 system affords a further example of the polymerizability increasing with decreasing Cx: having 10
valence electrons, NO@ shows no tendency to polymerize. Nitric oxide has 11 electrons, and its y x value
must be lower; accordingly, it dimerizes in the liquid
._
[54] R . Mueder, Helv. phys~caActa 2/, 41 1 (1948)
~
Angew. Chem. infernut. Edit.
/ Vol. 5 (1966) 1 No. 6
state a t a low temperature. The species NO 3, having 12
electrons and being formed in t h e reaction of NO with
alkali metals in liquid ammonia, exists only as a dinier[551. This shows t h a t the position of t h e element
pairs in Fig. 4 is not fixed, but varies according to t h e
type of bond and t h e o t h e r bonding partners. The
polymerizability of C=S has recently been confirmed in
t h e case o f FzCS which polymerizes in t h e presence o f
Insertion of this result into Eq. (4) leads to a set of linear
algebraic equations:
(F-AB)q == 0;
= wz,
(6)
which have non-vanishing solutions for q only at certain ).
values (eigenvalues). These eigenvalues are connected with
the frequencies of the fundamental oscillations by the relationship A = 4x*cW, and they satisfy the secular equation:
det(F-AB)
=
0.
(7a)
fluorides 1561.
Among the interesting bonds lying outside the shaded area in
Fig. 4, whose boundary is not yet well defined in the low-Ax
region, we may mention the P-C and the P-S bonds, which
have been studied with respect to their reactivities but
which require further spectroscopic investigation.
H. Appendix
Calculation of Force Constants and B o n d Orders
(by Dr. W. Sawodny)
I. Derivation of the Oscillation E q u a t i o n
157-591
The theoretical treatment of molecular vibrations starts with
the assumption that the atoms can be considered as points of
mass and the bonds between them as weightless springs, i.e.
that we are faced with a mechanical system of coupled oscillators. It is further assumed that the very small motions of the
mass points about their rest positions are harmonic vibrations and the potential curve can, therefore, be represented
approximately by a parabola. The kinetic energy T and the
potential energy V of the oscillating system are given by
2V
=
C fijqiqj = q’Fq
(fij
=
( 1 b)
fji)
IJ
where the qi,j are displacements of the mass points from
their equilibrium positions. The potential energy coefficients
f i j are described as force constants, and are a measure of the
strength of the springs, i.c. of the forces acting between the
atoms. For sufficiently small and slow vibrations, the coefficients bij and fij can be considered independent of qi.
By substituting expressions (la) and (lb) into the energy
equation
T-: V
=:
const.
(2)
The substitution G
leads to
=
B-1, carried out for convenience sake,
det(GF-~Eh)
=: 0
(7 b)
where E is the unit matrix.
Besides the eigenvalues A (fundamental vibrations) obtained
from the vibrational spectra, the determinant contains the
matrix G, which is made up of the atomic masses, the bond
angles, and the bond lengths of the molecule, and matrix F,
which is constructed of the desired force constants[*].
The derivation of the secular equation (7b) is independent of
the choice of coordinates for the oscillator displacement,
provided that they satisfy conditions (la) and (lb). On the
other hand, the coefficients bij and fij are determined by the
coordinates used. Instead of Cartesian coordinates, linear
transformations of these are generally used.
Most convenient are the internal valence coordinates, which
embrace only the internal vibrations of the molecule without
its translational and rotational motions (these have zero
frequency). The movement of the atoms is thus described in
terms of changes in either the bond angle or bond length. The
force constants thus assume the following physical meaning:
the diagonal eiements of the F matrix represent the valence
and the deformation constants, while the non-diagonal
members denote the coupling constants between the valence
and deformation modes. This is known as the General
Valence Force Field (GVFF) model.
The mathematical effort rapidly increases as the order of the
determinant (7b) increases, though it may be simplified by
virtue of molecular symmetry. Within a symmetry class, the
fundamental molecular vibrations are subdivided into various
symmetry species. One can find linear combinations of the
inner coordinates that correspond to the symmetry properties of the fundamental vibrations. These are called ”symmetry coordinates”. The overall secular determinant of the molecule can then be factorized into several smaller determinants
whose orders are given by the number of fundamental vibrations of the corresponding symmetry species. These smaller
determinants can then be solved separately.
Nevertheless, electronic computxs generally must be employed in the treatment of the vibration problem, particularly for
molecules of low symmetry or containing many atoms 1601.
we obtain
‘ / z (i’B6
+ q’Fq)
=
const.
(3)
Differentation with respect to time gives the equation of
motion
Bq+Fq=O,
(4)
a possible solution of which is
q
~
=
a sin wt;
q = ~ -ad
sin c d = - d q .
(5)
.~
[55] J . Goubeau and K . Laitenbergtr, Z. anorg. allg. Chem. 320,
78 (I963); K . Laitenberger, Dissertation, Technische Hochschule
Stuttgart, 1964.
[56] C. Walter, US-Pat. 3032537 (1962).
[57] R . Zurmiihl: Matrizen. Springer, Berlin-Gottingen-Heidelberg 1961.
[58] E. B. Wilson Jr., J . C . Decius, and P . C. Cross: Molecular
Vibrations. McGraw-Hill, New York 1955.
[59] J. M . Mills in M . Davies: Infrared Spectroscopy and Molecular Structure. Elsevier, London 1963.
Angew. Chem. internat. Edit.
1
Vol. 5 (1966)
/ No. 6
11. Solution of the Secular E q u a t i o n
Equation (7b) is termed the characteristic equation of the
eigenvalue problem, and can be solved by various methods
for the eigenvalues A [ 5 @ . This type of calculation presupposes
a knowledge of the force constants and can be used to confirm the assignment of the fundamental vibrations for a given
[ * ] If there is only one fundamental vibration, the determinant
reduces to one single equation gf - A = 0 which, for a diatomic
molecule, leads to 4 ~ i v2z = f
I:(
3
-t -
.
[60] D . E. Mann, T . Sliimanouchi, J. H . Meal, and L . Fano, J .
chem. Physics 27, 43 (1957); J . Overend and J . R . Scherer, ibid.
32, 1289 (1960); T . Shimanouchi and I . Suzuki, ibid. 42, 296
(1965); J . Aldous and 1. M . Mills, Spectrochim. Acta 18, 1073
(1962); D . A . Long, R . B. Gravenor, and M . Woodger, ibid. 19,
937 (1963); E. W . Schmid, Z . Elektrochem., Ber. Bunsenges.
physik. Chem. 64, 533 (1960); D. Papouselc and J . Pliva, Collect.
czechoslov. chem. Commun. 28, 755 (1963).
573
molecule by using the constants of compounds having a
similar structure 1611. However, the problem is generally the
reverse: we know the vibrational assignment and seek the
coefficients of the F-matrix. The derivation of Eq. (7b) in the
manner specified above involves the basic difficulty of having
to calculate for each symmetry species n(n+ 1)/2 GVFF-force
constants from n eigenvalues (in the case of a symmetric
F-matrix); as a result, the system of equation is insufficiently
defined for n > 1 .
We can now attempt to find further defining equations by the
use of additional experimental data, or to reduce the number
of required force constants to n by use of simplifying assumptions concerning the force field, or else to employ mathematical approximation methods.
1. A d d i t i o n a l E x p e r i m e n t a l D a t a
T o obtain additional information, fundamental vibrations
of isotope-substituted molecules can be used. The force field is
considered to be independent of the masses of the system, i.e.
the isotope substitution does not change the F-matrix, but
the fundamental vibrations are changed and a new G-matrix
results. Thus, additional equations for the force constants are
obtained. However, the mass of the isotope-substituted molecule must differ sufficiently from that of the reference molecule, since otherwise the additional equations roughly duplicate the original ones, and the indeterminacy of the F-matrix
is retained 1621. Furthermore, as the frequency shifts decrease,
their measurement involves an increasing error. Therefore,
H-D substitution is the most convenient for the calculation
of force constants, although owing to tne small masses, the
anharmonicity cannot be neglected introducing a further uncertainty. In addition, there are some isotope rules, such as
the product rule of Teller and Redlich [631, which represent
relationships independent of the force constants, and decrease the number of independent equations available for the
determination of the F-matrix.
The vibration-rotation interaction constants can also be
utilized in the determination of the force constants. Both the
centrifugal stretching constants DJ, DJK, and DK (641 and
the Coriolis coupling constants <i [651 depend on the F-matrix,
Although their relationship is a more complicated one, these
constants can still be used successfully.
The force constants are also related to the mean vibration
amplitudes obtained from electron diffraction studies 1661 as
well as t o the intensities of Raman and infrared bands[671,
although all these measurements are not yet sufficiently accurate for the application of the data in the calculation of the
force constants.
Unfortunately, these additional data are at present available
only for simple molecules. It may also happen that even the
use of such information leads to no unique solution of the
G V F F constants, and resort must then be taken to the application of simplified potential fields.
2. S i m p l i f i e d P o t e n t i a l F u n c t i o n s
To reduce the number of forcc constants, we can make
certain assumptions concerning the magnitudes of some
[61] See, e . g . , H.-J. Becher, W. Sawodny, H. Nfith, and W. Meister, Z. anorg. allg. Chem. 314, 226 (1962); H.-J. Becher, ibid.
271, 243 (1953).
I I ,molecdar Spectrosc3py 13, 338
1621 See, e.g., J . L. D U ~ I C UJ.
(1964); L. Becknlarnt, L . Gutjalir, and R . Mecke, Spectrochim.
Acta 21, 141 (1965).
[63] 0.Redlich, Z . physik. Chern. B 28, 371 (1935).
[64] D . Kivelson and E. B. Wilson J r . , J. chern. Physics 21, 1229
(1953).
[65] J . H . Meal and S. R . Polo, J. chern. Physics 24, 1126 (1956).
[66j S . J . Cyvin, Acta chern. scand. 14, 959 (1960).
1671 B. L . Crawford, J. chern. Physics 20, 977 (1952); G. W .
Clmrirry and L . A . Woodward, Trans. Fardday SOC.56, I l l 0
( I 960).
574
constants or the relationships between them. One possibility
is to borrow some of the constants from other molecules.
Particularly suited for this are constants obtained with the
aid of the Urey-Bradley field (681. However, this process
becomes questionable when the changes in the bonds caused
by the modified surroundings are to be investigated.
Simplification of the potential functions is best illustrated by
means of an example: a symmetric non-linear three-body
system ZXz of symmetry Czv executes three fundamental
vibrations, two of the type A1 and one of the type B1. The
potential energy is given, in accordance with Eq. ( I ) , by
hi)+ 2fr, A r l Ar2 + fa Az2 I 2fr,
2~ = f, ( h +
(Aria
Ar2x)
Therefore, the G V F F contains four constants to be determined.
The following two simplifying assumptions can be made. In
one case, we assume that all interaction constants are negligibly small, so that they can be put equal to zero. The potential function then contains only two constants:
This model is known as the Simple Valence Force Field
(SVFF). In the second case, we assume that the potential
energy is a pure square function of all the internuclear
distances, i.e. that the force acting on a n atom is the sum of
the attractive and the repulsive forces due to all the other
atoms, these forces coinciding with the lines between the
nuclei and depending on internuclear distances alone. This
is the ionic model of the Simple Central Force Field (SCFF),
in which the potential energy is given by:
2V
=
fr
car:+
Ar?$
+ fRhR2
Both the deformation constants and the coupling terms have
thus been expressed in terms of the interaction between the
two X atoms not bonded to each other. This postulate is
inadequate in the case of linear and planar molecules, since
it leads to a zero frequency for the out-of-plane vibrations.
With certain molecular configurations, such as tetrahedral
ZX4 molecules, not only the squared, but also the linear
potential energy terms must be taken into account. The rest
position of the atoms is then described as a state of equilibrium between the attractive and repulsive forces.
Both the SVFF and SCFF system are characterized by a
redundancy of parameters. Thus, in the given example, three
fundamental vibrations are available for the determination
of two force constants. The third equation can then be used
to verify the results. The results reveal that both postulates
involve oversimplifications [@I. It appears, therefore, that we
cannot neglect all the coupling constants. The SVFF treatment gives the better result when the central atom is heavy in
comparison to the ligands (weak coupling), whereas it is the
SCFF treatment that leads to the better approximation when
the central atom is light and the ligands are heavy (strong
coupling).
The results can be improved by the introduction of additional
interaction terms, so as to have the same number of force con~
~~
-
[68] See, e . g . , J. Overend and .I. R . Scherer, J. chcm. Physics 32,
1296, 1720 (1960); 33, 446 (1960); 34, 547 (1961).
[69] G. Herzberz: Molecular Spectra and Molecular Structure.
Vol. 11, Infrared and Iiarnin Spectra of Polyatomic Molecules.
Van Nostrand, Toronto 1951.
Angew. Chem. internut. Edit. / Vol. 5 (1966)
/ No. 6
stants to be calculated as the number of the available fundamental vibrations. These modified models have been used in
most force-constant dettrminations. The above mentioned
equalization may be achieved by the introduction of the
required number of G V F F coupling terms into the SVFF
model (Modified Valence Force Field, MVFF). Although
the choice of the interaction terms is arbitrary, we can often
rely o n analogy and past experience. In many cases all interactions between identical internal coordinates are used, since
this advantageously leads to a diagonal F-matrix in the derivation of the oscillator equation with symmetry coordinates.
The potential energy of the non-linear Z X 2 molecule is then
given by
2v - fr
+ Ari) + frr Arl Ar2 -1- r,l&
A model which combines the SVFF and the S C F F concepts
is the Urey-Bradley Force Field model (UBFF) [701. It contains the valence force constants and the deformation constants of the SVFF model and describes all couplings in
terms of forces acting between the non-bonded atoms:
2V
=
fr (Art
L
Ari) + fa Aa2+ fRAR2
This treatment generally affords valence force and deformation constants that are appreciably smaller than those obtained by the valence force fields, but gives relatively large
van der W a d s terms. As mentioned above, the force constants can often be transferred from one molecule to another
similar one, but in some cases the UBFF model must be
further modified to obtain meaningful results [711. For easy
comparison with force constants obtained by other methods,
the UBFF model is frequently employed to calculate the
potential energy with symmetry coordinates, but the resulting
constants are transformed into those of the valence force
field.
Attempts have also been made to derive simplified potential
fields by considering the coupling from the point of view of
bond theory. Thus, Heath and Linnett 1731 have used the
change in hybridization that occurs during the vibration, and
developed the Orbital Valence Force Field (OVFF). When
an atom is displaced from its equilibrium position by a change
of the bond angle, the overlap of the orbitals decreases. In
certain deformation vibrations, the orbitals of the central
atom can follow this motion either by a partial rotation of
the system or a rehybridization with change of the s and p
contributions. We thus arrive at a new definition of the deformation constants. The OVFF model resembles the UBFF
model in describing the coupling forces by the interactions
between the non-bonded atoms. Rehybridization of the
orbitals of the central atom can give rise to yet another definition of the coupling between valence and deformation modes :
a n increase in s-character shortens the bond, while an increase
in the p-component lengthens it. These considerations have
been incorporated in a model by Mi//s[731 called the Hybrid
Orbital Force Field (HOFF) model. This treatment has led
to good results in force constant calculations[74], but is at
present applicable only when d orbitals d o not participate in
the bond. A similar, though more complicated, model has
been developed by KingC751 for the NH3 molecule; it is known
as the Hybrid Bond Force Field (HBFF) model.
We must also mention a n attempt at using the interactions of
a x-electron system to describe relations between some of the
G V F F force constants. This is called the Tc-bond Interaction
Valence Force Field (x-IVFF) concept, and its use has made
it possible to reduce the number of constants to be determined
in the case of Ni(C0)4 from 17 to 10[761.
Table 9a. Force constants of tetrahedral molecules Z X j calculatcd \\it11 diNcrcnt potential field..
Com pd.
Method
OVFF
UBFF ( V )
MVFF
Fudini
GVFF (Rot. const.,
Isotop.)
OVFF
UBFF
UBFF (V)
MVFF (I)
MVFF (11)
Fudini
GVFF (Tsotop.)
GVFF (Rot. const.)
UBFF
U B F F (V)
Fudini
GVFF (El. diffract.)
GVFF (Raman int.)
MVFF
U B F F (V)
Fudini
GVFF (Isotop.)
OVFF
UBFF
U B F F (V)
MVFF
Fudini
GVFF (Rot. const.)
GVFF (Isotop.)
MVFF
Fudini
GVFF (Raman int.)
1701 H . C. Urey and C . H . Brodley, Physic. Rev. 38, 1969 (1931).
1711 T . Shimanouchi, Pure appl. Chem. 7, 131 (1963).
[72] D . F. Heath and J . W . Linne/t, Trans. Faraday Sot. 44, 873
( 1948); J . w.Linnetr and P. J . Wheotley, Trans. Faraday soc.
45, 33 (1949).
Angew. Chem. internat. Edit. / Vol. 5 (1966)
i Nu. 6
fI
4.95
4.944
5.37
5.391
5.050
4.3s
3.75 ( I < )
6.15
9.15
6.72
6.876
6.25
6.97
1.76 (I<)
2.98
3.268
3.59
3. I57
2,77
2.743
2.764
2.769
5.9
5.4 (K)
5.95
5.94
6.400
6.16
6.57
2.75
3.077
3.124
frr
FIE
~
--
fra
~~
0.036
0.051
-
0.008
0.142
0.151
0.036
0.6 (F)
1.35 (F)
1 .oo
~~
0.81
0.788
0.99
0.75
C.656 (F)
0.45
0.3 70
0.24
0.415
-
0.431
0.436
0.420
0.460
0.430
0.844
-0.3 (F')
0.50
1.13
0.567
0.90
0.59
-0.097 ( F
0.24
0.458
0.36
0.406
-
0.034
0.028
0.027
0.15 (F)
0.43 (F)
0.39
0.4 I
0.254
0.33
0.21
0.441
0.436
0.420
0.486
0.469
0.02 I
0
0.016
0.5 (z)
I .33 I
0.71
0.15 ( H )
0.62
0.71
0.71
0.710
0.70
C.71
0.08 ( H )
0.33
0.317
0.32
0.34
0.207
0.190
0, I88
0.188
1.014
1.12
1.01
-_
0.26 (%)
0.46
0.417
0.38
0.420
0.207
0.225
0.231
0.231
0.56
-0.31 (F')
0.26
0. I68
0.03
0.19
0.33
-
0.232
0.21 3
0.100
0.1 10
0.02 ( H )
0.26
0.268
0.3 ( x )
0.50
3.268
0.454
0.47
0.27
3.26
3.158
3.157
3.082
0.468
0.44
0.25X
0.236
3.235
[73]I . M . Mills, Spectrochim. Acta 19, 1585 (1963).
[741 J. Aldous and 1. M. Mills, Spectrochim. Acta 19, 1567
J ' L. D'"cn'l,
ibir'.
Igo7
1751 14'. T. Kitis, J. chem. Physics 36, 165 ( I 962).
[76]L . H . Junes, J. molecular Spectroscopy 5, 133 (1960).
7''
575
3. M a t h e m a t i c a l A p p r o a c h
Mathematical approximation methods may be used to eliminate the indeterminacy of the secular equation (7b) caused
by the fact that the number of force constants to be calculated
exceeds the number of experimental data available. Thus, in
one approach, the matrix product A = G F is reduced to the
Jacobian canonical form, i.e. the determinant (A] now
consists only of the product of the diagonal elements. The
eigenvaiue Ak involving the fundamental vibration vk is
considered as characteristic of the coordinate qk, but corrected for the interaction with the rest of the molecule. The
fundamental vibrations are first arranged in a series on the
basis of their magnitudes, and the individual terms are then
successively removed, so that the force constants fij (i > j,
will depend only on eigenvalues between Ai and hn[771.
A recent method [31 based on the Cayley-Hamilton theorem
gives n(n+ 1)/2 equations for n(n+ 1)/2 force constants, but
these equations are not mutually independent. However, this
mutual dependence can be eliminated by systematic slight
alterations of the equations and the use of an approximate
solution. In practice, we start with fully uncoupled vibrations
as a first approximation. For this case, a fully defined diagonal
F-matrix from the diagonal G-matrix can be obtained. Using
this F-matrix as a starting point, a complete F-matrix containing all the coupling terms can be calculated by the iterative and stepwise build-up of the known, complete G-matrix
from the diagonal matrix.
Table 9 b. Force constants of non-linear symmetric Z X ? molecules calculated uith c!'i?'xent potential fields.
Compd.
Method
H2
MVFF
UBFF
Tcrkiiigron
Fadini
GVFF (Isotop.)
UBFF
MVFF (1)
MVFF (11)
Torkingrou
Fudini
GVFF (Isotop.
and Rot. const.)
UBFF
MVFF
Torkingron
Fadini
GVFF (Isotop.)
GVFF (Rot. const.)
MVFF
Torkinglon
Fadini
GVFF (Rot. const.)
UBFF
IMVFF
0
ClOl
SO?
0 3
NO?
Torkington
FzO
frr
f,
Fadini
GVFF (Isotop.)
MVFF
Torkingron
Fadini
GVFF
(Rot. const.)
8.250 (K)
8.437
8.330
8.422
8.454
7.150(K)
7.16
6.78
7.144
6.8 I5
7.018
10.09 (K)
9.86
10.172
9.887
10.02
10.006
4.64
6.338
5.624
5.701
8.46 (K)
10.022
10.818
10.1 1 1
10.406
3.08
5.320
3.964
3.950
0.083 (F:
-0.118
-0.225
-0. 132
-0.100
-0.441 (F:
__
-0.22
0.724
-0.164
-0. I70
0.21 (F)
-0.12
0.290
-0. I08
0.03
0.024
0.703
2.160
1.440
1.523
3.03 (F)
1.652
2.210
1.741
2.024
-
-0.252
0.862
0.806
frcx
0.407 (F'
-
0.104
0
0.234
0.014 (F'
0.18
-
0.328
0.022
0.006
0.31 (F')
-
0.452
0.044
0.20
0.189
-
0.964
0.297
0.332
-
0.851
0.068
0.535
-0.14
0.680
0. I63
0.137
fi
0.984(H)
0.751
0.772
0.i54
0.761
0.823 (H)
0.63
0.635
0.778
0.618
0.651
0.993 ( l i )
0.8 I
0.796
0.808
0.793
0.793
1.696
1.131
1.308
1.285
0.395 (H)
I .036
0.870
1.021
1.097
1.13
0.788
0.715
0.724
The UBFF force constants have the following meaning:
K
F
stretching force constant
=
H
x
repulsive force constant between non bonded atoms
repulsive force constant, linear term
F'
bending force constant
~
internal tension
UBFF(V) denotes constants of valence force field calculated from
UBFF.
~
[77] P . Torkington, J. chem. Physics 1 7 , 1026 (1949).
[77a] T . Shimanourhi, J . Nakagawa, J . Hireishi, and M . isliii,
J . molecular Spectroscopy 19, 78 (1966).
576
The results of calculations carried out for some simple molecules using various force fields are listed in Tables 9a and 9b,
but their discussion must be preceded by brief mention of the
concept of coupling. If, in the triangular ZXz molecule used
earlier, the mass of the central atom m, is large in comparison
to that of the ligands m,, then the central atom will remain
virtually static during the vibrations:
In this case, the vibrations are practically pure valence or
deformation modes, owing to the lack of any coupling via thc
central atom. When, however, the mass relationship is the
reverse (i.e. m, < m,), it is predominantly the central atom
that oscillates; because of its similar motion in both vibrations, strong coupling between the two modes results.
Although the vibrations are no longer pure stretching and
deformation vibrations, they still retain some of their original
character, namely, the lower frequency is predominantly the
deformation mode and the higher frequency, the valence
vibration. On the other hand, in the case of strong coupling
between two valence vibrations complete mixing can take
place. When, for example, a tetrahedral molecule YZX3 is
characterized by the mass relationship of m, < m, = my,
we can no longer describe the symmetrical valence vibrations
as vzy and vzx3 sym., because these motions actually represent an in-phase and out-of-phase vibration vyzxj.
In addition to this coupling of masses, the coupling of vibrational energies also plays a part, as for example, in the NO2
molecule. In this case, the relationship m, < m, does hold,
but the valence vibration of the N-0 bond of bond order 1.5
and the deformation vibration differ in their energy so much
that they can interact only to a small extent.
Weak coupling does not mean that the coupling constant will
be very small. Calculations carried out on H 2 0 , C H 4 , and
NO2 by the GVFF method have given by no means negligible
values for the interaction constant frW However, the coupling
constant can vary within wide limits without notable deviations in the valence and the deformation constants. As coupling increases, the system becomes more and more sensitive
to changes in the interaction constants. It is, therefore, not
surprising that the MVFF method gives poor results in the
case of strong coupling. Quite often, no real solutions are obtained at all, as is the case with FzO, C 1 2 0 , C F 4 , and CC14;
nor can the situation be remedied by neglecting frr,since its
value cannot be neglected in most cases (e.g. CF4).
By contrast, the UBFF method gives a better approximation
for strongly coupled systems, but it leads to erroneous answers in the case of weak coupling. Whereas the constants
obtained for C F 4 and CC14 are in good agreement (after
transformation of the valence coordinates), the UBFF method
Angew. Chem. internut. Edit.
1 Vol. 5 (1966) No. 6
must be modified to result in reasonable values for CH4.
Similarly, additional interaction constants must be introduced in the calculations o n silicon compounds and molecules with resonance structures (e.g. NOz) [711. The resulting
modified model (MUBFF) gives good results, but generally
contains more constants than the available number of vibration frequencies “J7al.
The force constants obtained by the O V F F method are very
similar to those of the UBFF method. The H O F F method o n
the other hand, represents a considerable improvement. In
calculations on YZX3 molecules[741, it was shown to be
superior to the MVFF and the U B F F methods, although for
tetrahedral ZX4 molecules it again involves more constants
than there are vibration frequencies available.
The mathematical method put forward by Torkington [771
generally leads to excessively high coupling constants. By
contrast, the results obtained by the Cayley-Hamilton theorem[31 for intermediately coupled systems (CF4, SiF4, and
SiC14) are in good agreement with those obtained by the
G V F F method. For weak coupling, the resulting coupling
constant is practically zero, but as mentioned earlier, this has
only a small effect on the valence and deformation constants.
The method is not applicable to strongly coupled valence
vibrations, since a certain vibration character is a prerequisite
for the validity of the first approximation (uncoupled vibrat ions).
Summing up, we may say that unambiguous answers concerning the coupling constants are available at the present only in
very few cases, even with an exact knowledge of the vibrational frequencies and possible additional data as well as all
the molecular dimensions (for the influence of the latter,
see [781). Thus, even the values obtained by the G V F F method,
using various additional data, show considerable deviations,
as for example in the case of SiF4. As regards the valence and
the deformation constants, on the other hand, the appropriate
choice of the potential function ensures a n accuracy of about
10%. Higher accuracy presumably is obtained by the application of a unified method.
Deformation force constants sometimes cannot been given in
an explicit form, owing to the presence of “redundant coordinates” (e.g. in the case of ZX4 tetrahedra, cf. Table 9a). For
the time being, it is reasonable, therefore, to confine the discussion to valence force constants, which, incidentally, are
also the most important parameters as regards the properties
of bonds.
Table 10. Comparison of valence force constants calculated from e m pirical rules and from experimental data.
Force constants according. t o :
Bond
CH
NH
OH
BC
BN
BO
BF
Budsrr
G,,l.dJ
Sirbert
4.33
6.33
7.67
2.04
2.55
4.45
5.81
6.94
2.93
3.5.7
4.05
5.40
6.30
7.20
3.3P
3.94
4.50
2.92
3.14
2.93
3.77
4.31
4.78
cc
CN
co
CF
CCI
SiH
P IH
SiF
Sic1
PCI
AlCl
SIC
PC
GeH
GeCl
2.78
7.68
3.44
2.33
I ..57
1.99
1.16
1.53
1.90
2.82
I .25
4.49
5.06
3.76
4.54
5.25
5.85
3.54
2.29
2.74
3.28
2.13
2.50
1.82
2.20
2.56
4.05
4 73
5.40
6.08
3.40
3.74
4.00
2.17
2.10
4.20
2.35
2.52
2.18
2.80
3.00
3.60
2.27
5.05
6.00
6.45
3.64
-2.8
4.16
4.87
4.45
4.49
5.22
6.97
3.59
2.77
3.16
6.57
3.12
3.52
1.97
2.93
3.35
2.61
2.66
1
(CHn)
(NH?)
(HzO)
B(CH3)3
BX,--NR>
B(0H)y
B F ~
CH,-CHI
N(CH,)?
O(CH3),
CF,
CClJ
SiH,
PHP
SiF4
SIC14
Pclg
AICij
Si(CH,).,
P(CH,$
GeH,
GeCI,
where kij is a constant determined by the position of the
bonding atoms in the periodic table, k‘ is a correction factor
for ions, and ZA and ZB denote the number of valence electrons of the two atoms.
Another relationship involving the number of valence electrons z i has been put forward by Guggenheimer[811 in the
form:
f = R(zAzB)l12/rt
where R and t are constants whose values depend on the
polarity of the bond. Gordy’s relationship [821 involving the
electronegativities
Fah(XAXB
;?-) 3’4 c
is a special case of Guggenheimer’s formula[831. The constants
a and c here generally have the values of 1.67 and 0.30 (other
values have to be assumed when one of atoms involved has
only one valence electron) ; b denotes the bond order.
Siebert
151
has devised the particularly simple formula:
111. Empirical Rules for Valence Force Constants
There have been attempts to correlate the valence force constants with other bond properties such as internuclear distance, dissociation energy, and the number of electrons in the
bonding partners. The literature contains many such empirical
correlations, though they are generally derived and verified
only for diatomic molecules (Table 10).
The most important empirical rules are as follows:
Badger’s ruler791 correlates the force constant f with the internuclear distance r :
where dij and Cij are constants whose values are determined
by the position of the bonding atoms in the periodic table.
A similar correlation is given by Clark’s formula [801:
where the constant P invariably has the value of 7.20, Zi is the
nuclear charge and ni is the principal quantum number of the
valence electrons of atom i. This formula enables us to calculate the single bond values fi without knowing any experimental data such as bond lengths. However, this formula is
strictly valid only for bonds of sp3 hybridization. It can be
shown that the formation of free electron pairs at the bonding
partners (i.e. the setting-up of overlapping p-orbitals)
weakens the bond. This can be taken into account by the
introduction of an empirical correction factor in Siebert’s
formula“]. Furthermore, the single bond value of the force
constants increases with increasing s-character of the
bond 1841.
[81] K. M . Guggenheimer, Proc. physic. SOC. (London) 58, 456
(1946); Discuss. Faraday SOC.9 , 221 (1950).
[82] W. Gordy, J. chern. Physics 14, 305 (1946).
1831 R. L . Williams, J. physic. Chem. 60, 1016 (1956)
[84] J . Goubeau, unpublished work.
[78] R . R . Hart, Developments appl. Spectroscopy 4 , 171 (1965).
[79] R. M . Badger, J. chem. Physics 2, 128 (1935).
[80] C . H . D . Clark. Philos. Mag. 18, 459 (1934); H . S. Allen u.
A. K . Longair, Philos. Mag. 19, 1032 (1935).
Angew. Chem. internat. Edit. / VoI. 5 (1966) / No. 6
1851 L. Solem, J. chem. Physics 38, 1227 (1963); L S a l e m , J. chem.
Physics 38, 1227 (1963); P . Phillipson, ibid. 39, 3010 (1963); 44,
633, (1966).
[86] S. Bratoz u. M . Allavena, J. chem. Physics 37, 2135 (1962);
J. Chirn. physique 60, 1199 (1963).
577
Siebert also gives a formula for the determination of the bond
order b from the single-bond value f l :
b = rbfhiri
ri
Moreover, we can
tibe
up to b = 1.5 whenever the distances rl and Tb are not known.
However, for bond orders over 1.5, we must use the formula
-
0.29
7 0,71
Acknowledgments
an even simpler relationship
b = fb/fl
b
cules (H20, NH3) force constants have been calculated from
energy values obtained from an approximative wave function [86J.
fbbifl
containing empirical correction factors; fb always is the experimental force constant.
Finally, it should be mentioned, that also attempts where
made to find relations between quantum mechanical data, as
the electron density, and force constants [851. For small mole-
The author wishes to thank his colleagues for their assistance in the syntheses and the spectroscopic investigations, and especially Dr. W. Sawodny for the calculation
of numerousforce constants. He also wishes to express his
gratitude to the Deutsche Forschungsgemeinschaft, the
Ministry of Scientific Research, and the Fonds der Chemischen Industrie for their financial support.
Received: November 26, 1964; revised March 14th, 1966 [A 518 IE]
German version: Angew. Chem. 78, 565 (1966)
Translated by Express Translation Service, London
C 0M M U N ICAT I0N S
Amination of Hydrazines to Yield Triazanes
By Prof. Dr. Ernst Schmitz, Dip].-Chem. S. Schramm, and
Dipl.-Chem. Heide Simon
Institut fur Organische Chemie der Deutschen Akademie der
Wissenschaften, Berlin-Adlershof (Germany)
Experiments with l5N have shown that hydrazine can be
aminated by hydroxylamine-0-sulfonic acid, the reaction
proceeding by way of triazane [I].
We have found 2-acyloxaziridines [21 to be extremely effective
aminating agents. The compound ( I b) aminates cyclohexylamine at room temperature within one minute, giving the
hydrazine derivative (2) (m.p. and mixed m.p. [31 141 "C).
This reaction led to attempts to aminate hydrazines to
triazanes.
C&11-NH-NH-CO-NH-C&i
H,
(2). 61%
-, y - C O - m - R
in which fission of the N-N bond is coupled with dealkylation. Dilute acid liberates ammonia (0.7 mol) and formaldehyde (0.5 mol) from compound (3) within a second [reaction
(b)], and because of this competing reaction the yield of the
iodide reaction amounts to only 70-75 %. Neutral aqueous
solutions of compounds (3) and (4) decompose five minutes
at room temperature with formation of tetramethyltetrazene
( 5 ) and urea or phenylurea, respectively [reaction (c)].
Compound (5) was identified by gas chromatography,
phenylurea by melting point and mixed melting point, and
urea by means of xanthydrol. The same products are formed
when the crystalline compounds ( 3 ) and (4) are kept at
room temperature for one hour. During a melting-point
determination, compound (4) reacts at 80-85 O C and then
shows the melting point of phenylurea.
Compounds (3) and ( 4 ) are formulated as triazanium
betaines, in analogy to the 2,2-dialkyltriazanium salts obtained by R. COsl[41 from N,N-dialkylhydrazines and hydroxylamine-0-sulfonic acid. Final proof of structure was
provided for the triazanes prepared analogously from monoalkylhydrazines. A crystalline compound (6) (m.p. 63-65 "C)
separates in 30 % yield from a solution of ( l a ) and methylhydrazine in tetrahydrofuran; (6), too, oxidizes iodide to
iodine. Its thermal or acid-catalysed decomposition leads to
ammonia, formaldehyde, and semicarbazide [reaction (d)] ;
the semicarbazide contains the N-N bond obtained on
formation of the triazane.
2-Acyloxaziridines, ( l a ) and ( I b ) , suspended in benzene,
dissolved immediately on addition of N,N-dimethylhydrazine,
and crystalline compounds separated in 70-80% yield in
1-2 minutes at 5-10°C.
Elemental analyses showed that products containing four
N-atoms were formed with loss of benzaldehyde. These
products liberated iodine from acidic iodide solution in the
cold [reaction (a)]. As with peroxides, oxaziridines, and
diaziridines, acid initiates an intramolecular redox reaction,
I z + NH3
5 3 3
H2N-%-NG-CO-NH-R
n
578
-
P
----)
\
CH3
(3) ,R = H
I , ,
z
(a)
CHlOf NH3 (b)
I
hydrazine
derivatives
with unknown
f
On reaction with cyclohexylhydrazine, compound ( I b ) affords the triazane (7) (yield 70 %). The position of the alkyl
group, not rigidly proved for (6), was made certain for this
product (7) : the colorless crystals (7), when brought into
contact with acid, pass within seconds into yellow crystals
(m.p. 89 "C), the azo compound (8) being formed with loss
of ammonia. In (8) the cyclohexyl group has the position
required by the triazane structure (7).
structure
u
Angew. Chem. internat. Edit.
Vol. 5 (1966) J No. 6
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