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Electron Affinity of Organic Molecules.

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Pedersen therefore proposed that the term “gel filtration” should be replaced by the term “exclusion chromatography” for the high molecular-weight range.
Other authors disliked the name “filtration” fcr this
chromatographic process and suggested the terms
“restricted diRusion chromatography” [ 1581 or “gel
permeation chromatography” [30]. However, 7iselius
supports the hitherto usual term “gel filtration” [1801,
and points out the fundamental difference between this
and other chromatographic separations, namely, that
the concentration in the stationary phase can, at the
most, be only as great as that in the mobile phase.
[180] A. Tiselius, J . Porath, and P.-A. Albertsson, Science (Washington) 141, 13 (1963).
German version: Angew. Chem. 76, 635 (1964)
Translated by Express Translation Service, London
Some of the difficulties in nomenclature can be avoided by
agreeing to the division attempted in this summary (from
practical considerations, among others) into “gel filtration”
and “gel chrornatogrnphy”. Whatever name is given to the
process, its use guarantees clean and easy separation of substances with different molecular sizes, as shown by the profusion of the applications quoted.
I should like to thank Doz. Dr. B. G‘elotte, A. B. Pharmacia, Uppsala (Sweden),for information and stimulating
discussions.
Received, October 22nd, 1963 [A 384/177 IEI
Electron Affinity of Organic Molecules
BY PROF. DR. G. BRIEGLEB
INSTITUT FUR PHYSIKALISCHE CHEMIE DER UNIVERSITAT WURZBURG (GERMANY)
Dedicated to Dr. W . Foerst on the occasion of his 65th birthday
Theformation of organic radical ions by direct electron exchange between a donor molecule
D and an acceptor molecule A is becoming increasingly imporfant,particularly with regard to
reaction mechanisms involving intermediate radical ions. All redox processes also involve
electron exchange. Finally, the knowledge of electron affinities and ionization energies of
organic molecules is necessary for a quantitative theoretical interpretation of intermolecular
resonance and of the semiconductor properties of crystalline organic compounds and molecular compounds.
The electron-donating properties of organic molecules
can be estimated from their ionization energies which
for some organic molecules are directly measurable.
However, direct experimental determination of the
electron affinities of organic molecules, which govern
their electron-accepting properties, was not possible
until only a few years ago. This paper deals with the
methods and means OF determining the electron affinities
of organic molecules; the discussion will be restricted
to neutral molecules and will exclude consideration of
radicals as electron donors and acceptors. The electrondonor and electron-acceptor properties of aromatic
hydrocarbons will be treated separately in Chapter IT.
All the electron donors discussed this article are listed
in Table 1 .
Ej = -EA
=
EA-
- ‘A
+ eo-
The ionization energy I of an electron-donor molecule D
is equal to the energy Ei of the highest occupied orbital
of D in the ground state or to the difference in energy
(E,-ED+ + e o - ) between the cation D+ + en- and the
uncharged molecule D.
Angew. Chem. internat. Edit.1 Vol. 3 (1964) 1 No. 9
=
ED - E D’ +
eo-
The difference in energy between the lowest vibrational
level of the ground state of the neutral molecule and
the corresponding level of the ion is termed the a d i a b a t i c ionization energy, or the a d i a b a t i c electron
affinity, respectively. The energy difrerence between the
energy level of the ground state and that part of the potential curve of the ionized state to which - applying
the Franck-Condon principle - transition is most likely
to occur, is denoted as v e r t i c a l ionization energy and
v e r t i c a l electron affinity, respectively.
I. Organic Electron Acceptors
(excluding Aromatic Hydrocarbons)
The electron affinity EA of an electron-acceptor molecule A is the energy E j of the lowest unoccupied orbital
of A in the ground state and is equal to the difference
in energy (EA--EA + eo- ) between the anion A- and
the unch,irged molecule A + eo-.
(1)
Ei = -I
(2)
1. Electron Affinities Derived from the Electron
Absorption Coefficient
Lovelock [ 1 ] determined the electron absorption caefficient y,
(3)
~~
--
Ie-I
=
[e-10 ex&% [A])
--
[ l ] J. E. Lovelock and R . S . Lipsky, J. Amer. chern. SOC.82, 431
(1960); J. E. Lovelock, Nature (London) 187, 49 (1960); 188,401
(1960); 189, 729 (1961).
617
Table 1. Electron donors discussed In the present paper. The numbering adopted here is also used in the text, in Figures I , 2, 3, 7, 8, and 9. and in
Tables 3, 5 , 6 , 7, and 8. Numbers 66, 85, and 87 have keen omitted for technical reasons.
No.1
Name
Formula
Nainz
0
1 ]Benzene
Hexaphene
5
6 Benzo[c]pentaphene
7 Dibenzo[c,m]pentaphene
Naphtha[2,3-clpentaphene
8
Benzo[a]perylene
9
D~henzo[a,j]peryIene
I
0
10 Triphenylene
Formula
s
Name
0.1
Formula
&
I
9 Acenaplxhylene
8
8
s
n
1 Benzo(c]tetraphene
2 Toluene
3 m-Xylene
4 Mesitylene
OCH3
3c,
L>CH,
IH,C
<>CH,
H3C
CHI
I Dihenzo[a,n]perylene
D~henzo[h,pqr]perylene
@?
Trihenzo[h,k,pqr]perylene
88
2
-
3
4 Di henzo [h,klper ylene
'8 20-Methylcholanthrene
H3C@
I
U
15 trans-Stilhene
s
15 Benzo[tuv]bisanthene
OCH=CHC
,9 9,10-Dimethylhenz[a]anthracene
n
16 Benzo[j]terrylene
10 Aniline
I I N,N-Dimethylaniline
$2
18 Benzo[e]pyrene
I-Naphthylamine
Cfi,
QNHz
ON(CH,),
CP
$3 2-Naphthylaminc
19 Benzo[a]pyrene
19 Benzo[a]phenanthrene
20 Dibenz[a,j]anthracene
21 Picene
1
52 Dihenzo[e,l]pyrene
22 Pentaphene
23 Perylene
54 Dihenzo[a,i]pyrene
2 4 Benzo[ghi]perylene
55 Naphtho[2,3-e]pyrene
25 Coronene
68
26 Biphenyl
@
32 Benzo[c]fluorene
&H?
56 Naphtho[2,3-a]pyrene
1
57 Dinaphtho[2,3-a:2,3,-h]pyrene
%
93 9-Benzylidenefluorene
94 Benzo[ghi]Auoranthene
95 Benzo[h]fluoranthene
K
I
3 1 Trihenz[a,c,J]anthracene
33 Benzo[c]chrysene
34 Benzo[a]tetraphene
w
6 3 Dibenzia,j]anthanthrene
64 Dihenz[a,k]anthanthrene
Angew. Chem. internat. Edit. 1 Vol. 3 (1964) 1 No. 9
for the capture of thermal electrons e- by a few organic
molecules A. This represents the only possibility for
experimental determination of the relative order of
electron affinities. Under experimental conditions for
which [ec]
[A], the absolute electron absorption
becomes identical with the equilibrium
coefficient
constant K A 121:
>
x
gAgA'ge-
K - - - e
(4)
A-
EA/RT -~[A-I
- 1-41 [e-I
The electron affinity EA can be calculated from Equation (4) by setting gA-/gA = 2, and hence gA-/gA .ge- = l.
Lovelock reported several electron absorption coefficients x'
K'A referred to chlorobenzene as standard
(Table 6. Column 2). Wentworth and Becker [3] determined the absolute equilibrium constant for electron
capture by anthracene. Using this value and Equation
(4), they obtained a value of 0.42 eV for the electron
affinity of anthracene. For a given acceptor A, and for
Klanthracene
= 12 [za], we have:
5
RT In K,
(5)
=
EAA = RT In K ,
+ 0.42 - RT In 12
The electron affinities of a few organic acceptors,
calculated for
= KIA using Equation ( 5 ) are listed
in Table 5 , Column 3.
x'
2. Electron Affinities derived from the Electron-Transfer
Energy and from the Ionization Energy of the Electron
Donor
On the basis of Mulliken's theory [4] of intermolecular
interaction between electron-acceptor and electron-donor
molecules, it is possible to derive an expression for the
energy FjcT of electron transfer from D to A in terms
of the ionization energy I of the electron donor and the
electron affinity EA of the acceptor [5,6]:
(6)
FjcT
~
=
I - (EA-E)
+
E is predominantly t h e energy of t h e C o u l o mb attraction b e
tween D+ a n d A- after t h e transfer of o n e electron f r o m D t o
A . E also comprises t h e polarization energy [6, 71 a n d a n
energycontribution AH from t h e intermolecular b o n d D.... A
in t h e ground state. C2 can be calculated theoretically [61.
Theoretically, Equation (6) is based o n t h e adiabatic ionization energy. However, i t is not yet certain whether t h e
various methods for determining t h e ionization energy of
organic molecules (e.g., t h e electron impact a n d photoionization methods) yield the adiabatic or t h e vertical ionization
energy.
It has been established that E and C2 can be assumed
to be nearly constant, for electron-donor/electron-acceptor complexes (EDA complexes) of the same bond
type [IO], so that for complexes involving one and the
same acceptor, Equation (6) reduces to Equation (7).
C1 can b e determined experimentally [ I l l by plotting 7cr
against I for t h e EDA complexes o f a given acceptor with
various donors, when Equation (7) is normalized t o a standar d , e.g. naphthalene.
T h e electron affinities listed in Column 4 of Table 6, were
calculated f r o m C1 according t o Equations ( 6 ) a n d (7) using
t h e value E = Ecou~= -4.3 eV for the x,r-complexes. It was
assumed here t h at t h e average intermolecular distance for t h e
z,x-complexes is nearly constant. This assumption is supported by t h e fact t h at t h e validity of Equations (6) a n d (7)
has been demonstrated f o r a large number of E D A complexes [I 11. T h e charge o n D+ a n d A- is shifted towards t h e
center of t h e complex molecule. Because of t h e approximations used, electron affinities derived f r o m C1 (Table 6,
Column 4) have a n uncertainty o f ab o u t *0.2 eV.
In Table 2, the electron affinities calculated by Person
[9] and by Jorfwer and S?kolov [I21 for the l~:.l~gens
Table 2. Electron affinities [eV] for halogens and iodine chloride.
CI, Eqs. (6) and
cz
I-(EA-E)
EAvertical [ * ]
I
EA from Eq. ( 2 5 )
[**I 112.161
I
~
[2] [e-] : Stationary electron concentration at equilibrium; [e-10:
electron concentration before approach of acceptor A to e!ectron
gas; g: number of realizable quantum states ("quantum weight").
For e-, g = 2 ( s = + 1 1 2 ) ; see,for example. 7 . L . H i l l : Introduction
to Statistical Thermodynamics, Addison-Wesley Publ. Co., Reading, Mass. (U.S.A.) 1960. Equation (3) can be written, to a good
approximation, as follows: [e-l= [e-10 - [A-l= leIo(1 - x[Al).
Then, for the case where [e-3
[A] or [A-] << [ e - ] ~we obtain:
>
[2a] If K c ~ BIS~ the electron-capture equilibrium constant for
chlorobenLene, then, for an acceptor A, we have:
KA/KCIBz
K'A and Kanthracene/KCIBz = l 2
and hence: (4') KA = ( K ' ~ i 1 2 )' Kanthracene,
The change in free energy for electron capture by A is by definition -EA = -RT InKA, and, hence, using Equation (4'), we obtain Equation (5).
[3] E. W. Wrntworrh and R . S . Becker, J. Amer. chem. SOC.84,
4263 (1962).
[4] R. S. Mnlliken, J. Amer. chem. SOC.72, 600 (1950); 74, 811
(1952); J. chem. Physics 19, 514 (1951); J. physic. Chem. 56, 801
( 1952).
[5] S. H. Hustings, J. L. Franklin, J . C. Schiller, and F. A. Matsen,
J . Amer. chem. SOC.75, 2900 (1953); G. Briegleb and J. Czekalla,
2. Elektrochem., Ber. Bunsenges. physik. Chem. 59, 184 (1955);
6 3 , 6 (1959); Angew. Chem. 72, 401 (1960); R . Foster, Nature
(London) 181, 337 (1958).
161 For further details, see G. Briegleb: Electronen-DonatorAcceptor Komplexe. Springer, Berlin-Gottingen-Heidelberg, 1961.
Angew. Chern. internat.
Edit. Vol. 3 (1964) No. 9
1.7 = 0.5
1.56
I
I
1 . 2 1 0.5
1.48
1
1
1.3
+ 0.4
1.3
I
I
1.7 1 0.6
1.43
[*I Calculated theoretically from the potential curves for Hall and
Hal-2 by applying the Morse function and making use of the interatomic equilibrium distance, dissociation energy, and fundamental
vibrational frequency.
[**I Relative t o the iodine atom with EA = 3.23 eV [I51
_
~ -~ _
[7] The valence bond energy Eval contributes to a slight extent
to the polarization energy Epol and vice versa. In general, Epol
should be equal to about 0.1 Ecoul..Only for iodine-amine complexes does the contribution of Eval apparently become larger
[8,8a], as a result of the relatively small intermolecular distance
(about 2.3 .&). According to Person [9], for iodine-amine complexes Eval w 0.25 ECoul,: In the case of n, a-complexes, for
example, those between iodine and amines, the constants C1 and
C Zin Equation (6) are markedly different compared t o those for
n,n- or n,a-complexes. H . Yada, J . Tanaka, and S . Nagnkura,
Bull. chem. SOC.Japan 33, 1660 (1960) proved this for iodineamine complexes. See also [8 a].
[8] See [ h ] ,pp. 7, 18. 76, 79, 81, 132, and 173.
[gal R. S. Mulliken, Annual Rev. physic. Chem. 13, 107 (1962)
[9] W . B. Person, J . chem. Physics 38, 109 (1963).
(101 x,o-; n,n-;n,a-; a,a-complexes; for further information see
161, pp. 5 et seq.
[l 11 C. Briegleb and J. Czekalla, [5]; see also [6], pp. 74 el seq.
[12] J. Jormer and U. Sokolov, Nature (London) 190, 1 (1961).
619
Table 3. Values of a and b [eVl from Equation (9) for electron-donorielectron-acceptorcomp!exes
2,6-Di.
Tetracyanoethylene
Chloranil
0.87 [*I
0.90
0.72 ['*I
Brornanil
chlorop-benzoquinone
Ch'orop-benzoquinone
0.85
0.85
___
1
0.93 1'1
0.97 [**I
p-Benzoquinone
I
'
g
3
5
Trinitrobenzene
0.85
0.97
0.85
~~~
4.86 [*] 3.48 [**I
[191
[*I
["I
I18al
4.5
1
4.94 ["I
5.27 [**I
4.36
4.19
1
3.89
4.5
_________
[18b,191
[18bl
[18b]
I181
[181
1181
I
3.9
i
2,4,6Trinitrotoluene
m-Dinitrobenzene
0.85
0.85
1,.
I
0.94
087
~ _ _ _ _ _ _ _
3 68
3.47
4.2
3.6
I
__________
l18cI I181
[181
1181
[IScI
[8aI
€or polycyclic aromatic hydrocarbons
Average values for substituted benzenes. To be precise, methoxybenzenes (a = 0.74; b
halogenobenzenes (a = 0.60; b = 2.34) [18al woold have to be distinguished.
and iodine chloride are given for comparison with
those of organic acceptors. The order of these affinities
seems reasonable.
X-ray studies [I31 of the complexes of iodine and the
other halogens with n-donors, particularly benzene,
have indicated that a configuration in which the axis
of the halogen molecule is vertical to the benzene plane
is more probable than the alternative configuration
with parallel axes. According to Equations (6) and (7)
E, and Ell differ by about 1 eV [14].
At sufficiently large ionization energies (I > 7.5 eV),
the third term in Equation (6) can be neglected:
Hence, for EDA complexes between one and the same
acceptor and various donors, the following approximation is often used 1171,
As comparison with Equation (6) shows, the term b in
Equation (9) contains the electron affinity of the acceptor and the Coulomb energy E ; thus, when E is nearly
constant, b can be taken as a measure of the electron
affinity and is therefore often referred to as the a p p a r e n t electron affinity. For EDA complexes of the
same bond type [lo] the term a can be assumed to be
constant, i. e., largely independent of the acceptor. This,
however, is not always true, particularly for interactions
with weaker donors [18] (cf. Table 3, where the values
of a and b obtained from Equation (9) are listed for a
few acceptors). The values of a and b for complexes of
the very strong electron-acceptors tetracyanoethylene
and bromanil with polycyclic aromatic hydrocarbons
[13] 0. Hassel, Molecular Physics 1 , 241 (1958); 0. Hassel and
K . 0. Stromnte, Acta chem. scand. 12, 1146 (1958).
I141 The value EA = 0 . S eV for 12 quoted by Briexleb and Czekalla 151 refers to the parallel model, which was originally adopted
because of symmetry considerations [4] arising from the quantum-mechanical group theory.
[IS] 0 . H . Pritchard, Chem. Reviews 5 2 , 529 (1953).
(161 It is improbable that the values of E for EDA complexes
of atomic iodine, molecular halogens, and TC1 are identical. However, the order of the electron affinities calculated by Jortner and
Sokolov [12] agrees -. within the limits of error - with the electron affinities calculated by Person [9] for the vertical model.
1171 H . McConnel, J. S . Ham, aad J . F. Platt, J . chem. Physics
21, 66 (1953).
[18] R . Faster, Tetrahedron 10, 96 (1960).
620
=
3.64), methylbenzenes (a = 0.68): b = 3.01), and
differ from those obtained for their complexes with
substituted benzenes [18a, 191.
Judging from the strengths of their EDA complexes
with aromatic hydrocarbons, the following order of
electron affinities can be derived from the b values
listed in Table 3:
bromanil > tetracyanoethylene > chloranil = 1,3,5trinitrobenzene > 2,6-dichloro - p - benzoquinone >
chloro-p-benzoquinone > p-benzoquinone > 2,4,6trinitrotoluene > m-dinitrobenzene.
This order is not ic7 agreement with the relative order
of electron affinities resultiog from Equations (9,(7),
(14), (23), and (25); see Table 6. Obviously, the constant
b in Equation (9) is less suitable to determine a relative
sequence of electron affinities.
From Equations (8) and (2) we obtain
As we shall see in the following, there are numerous
possibilities for expressing I or -Ei by means of other
quantities.
3. Electron-Transfer Energy as a Function of the
Ionization Energy of the Donor
For polycyclic aromatic hydrocarbons as donors, it
follows from Hiickel's simple A M 0 theory that :
(11)
E.=-l=a o +
f".P
In Equation (1 l), a0 is t h e Coulomb integral, the resonance
integral for t w o neighboring atoms (k, l), an d fH is t h e Hiickel
coefficient which expresses the contribution of p to t h e total
energy Ei. T h e electron overlap integral is assumed to b e zero,
an d t h e electron repulsion is neglected.
T h e decrease in electron repulsion on transition f r o m D t o D+
may be taken i n t o account (by a decrease in t h e Coulomb
attraction) in or0 [20]. In a simplification recently proposed by
-____
[ l sa] E. M . Voigt, Ph. D. Thesis, University of British Columbia,
1963.
[18b] M . Kinoshita, Bull. chem. SOC.Japan 35, 1609 (1962).
[18c] See [6]; pp. 79 et seq.
1191 H . Kurodn, M . KrrbaJ%zshi,M . Kinoshita, and S . TakenJoro,
J. chem. Physics 36, 457 (1962).
1201 G. W . JVhehelnnd and D . E. Mann, J. chem. Physics 17, 264
(1949) assume that the increase in a, expressed in units of 8, is
proportional to the positive charge (I-q)p on the carbon atom
(q is the relative negative charge on the carbon).
Angew. Chem. internat. Edit./ Vol. 3 (1964) No. 9
Ehrenson [22], OL can b e expressed in units of
a n d p a r e negative; la1 < la"():
p
a s follows
(cr,cro,
a
(12)
=
ao-(l-l/n)c@
where co is a n empirical proportionality factor a n d n is t h e
number of conjugated =-electrons in t h e aromatic hydrocarbon.
Substituting cc for CIO in Equation (11) and using
Equation (12), we obtain:
-Ei
(13)
=
I
=
- (f
-NO
H
-
n- 1
~~
n
w)p
= -a0
- xp
When cc = -9.878 eV, F = -2.11 eV, and w = 1.4 [21],
Equation (1 3) gives satisfactory values for the ionization
energy of aromatic hydrocarbons [22].
Substituting Equation (13) into Equation (lo), one
obtains :
(14)
GCT
=
(14a)
k
-EA-ao
=
(-EA-ao
+ E)-@
=
The charge-transfer band maxima GCT for EDA coniplexes between the acceptors given in Table 4 and the
donors listed in Column 5a of Table 6 and the corresponding x values are related by the linear function
represented by Equation (14). The values of p and k
calculated from this linear QCT vs. x relationship, and
the electron affinities derived from the k values obtained
from Equation (14a) by setting E = -4.3 eV and cco =
-9.88 eV are listed in Table 4 for the electron acceptors
tetracyaneethylene, chloranil, 2,4,7-trinitro-9-fluorenone, 1,3,5-trinitrobenzene, and iodine [23].
Table 4. Electron affinities of a few typical organic electron acceptors
and of iodine, calculated from the values of k obtained according to
Equations (14) and (14a) [*I.
1
1
k-xp
Tetracyanoethylene
1
I
I
-p [eVl 2.2
+E
This expression is based on the observation that the
values of EA and E for EDA complexes of analogous
donors (e.g. polycyclic aromatic hydrocarbons) with
one and the same acceptor are nearly constant.
According to Equation (14), CcT should be linearly
proportional to K. Figure 1 shows for the example of
the EDA complexes of tetracyanoethylene with aromatic hydrocarbons that this requirement is satisfactorily
fulfilled, considering the variability of E and CcT and
the uncertainty in I as obtained from Equation (13).
(The values of x and f, in Figure 1 are listed in
Columns 5 and 6 of Table 4.)
Chloranil
~
2'437Trinitro9-fluorenone
1 ;:;
I
;::
2.3
k [eVl
3.67
EA [eVl 1.9
1
1
Trinitrobenzene
I
2.6
1
1.9
&
;
['I For 12, E = -4.3 eV or E = - 3 . 5 eV, depending
parallel or vertical model is considered.
~
011
4.55
1.0 (11)
whether the
Within the limits of error (!c 0.2 eV), the electron
affinity values calculated using Equation (14) are in
good agreement with those calculated from C1 in
Equation (7); compare Columns 4 and 5 of Table 6.
The values of listed in Table 4 refer to the aromatic
hydrocarbons summarized in Column 5a of Table 6
and correspond to the values quoted by Streitwieser [21]
and Hoijtink [31].
9
4. Ionization Energy and Intermolecular Electron-
Transfer Energy as a Function of the Electron
Excitation Energy C , of Aromatic Hydrocarbons as
Donors
Setting for a given molecule, in general:
(15a)
Ei = cc+ f . p
(15b)
Ej = OL
and
+ T.8,
we obtain For the energy difference between the highest
occupied and the lowest unoccupied orbital :
1/
073
I
-100
EBZ
I
I
- 60
-80
x
I
I
-20
-LO
loZ--
Fig. 1. Linear dependence of the electron transfer energy CCT on the
value of K obtained from Equation (14) for tetracyanoethylene-hydrocarbon complexes (the numbers refer to the hydrocarbons listed in
Table 1).
[23] Instead of Equation (lo), Dewar and Lepley [241 use
C ~ =
T I-EA = E.-Ei
(10)
In other words, the energy term E is omitted. Moreover, Dewar
and Lep1e.v use Hiirkel's simple equation [Equation ( I l)] for I and
obtain, in place of Equation (14), a linear relationship between
CCT and f H :
(14')
~ C =
T k'-fH P
(14a')
k' = - E A - q
~
[21] A. Streitwieser and P . M . Nair, Tetrahedron 5 , 149 (1959);
A. Strritwieser, J. Amer. chem. SOC.82, 4123 (1960). - Hoijtink's
calculations [3 I], based on the one-electron half-wave reduction
potential of aromatic hydrocarbons (p. 626 and p. 629), in combination with Hiirkel's equation [Equation (15b)l and Whelnnd's
f'-values [20], gave an average value of p= -2.26 eV. See also
P. Balk, S. De Brui.vn, and G . J. Hoijtink, Recueil Trav. chim.
Pays-Bas 76, 860 (1957).
[22] S . E. Ehrenson, J. physic. Chem. 66, 706 (1961).
Angew. Chem. internat. Edit. 1 Vol. 3 (1964)
1 NO. 9
[*I
I
2.35
~
Iodine
However, the simple. uncorrected Huckel equation [Equation
(1 l)] gives less satisfactory values for the ionization energy than
Equation (13) 1211. In addition, the values of k' calculated by
Dewar and Lepley from Equation (14') differ by a factor of E
from those obtained using Equation (14). As a result, the EA
values derived from k' are 3.5-4.5 eV too high.
[24] M . J. S . Dewar and A. R. Lepley, J. Amer. chem. SOC.83,
4560 (1961); M . J . S . Dewar and H . Rodgers, ibid. 84, 39.7 (1962);
A. R. Lepley, ibid. 84, 3577 (1962).
62 1
1.95
(1)
From Equations (15a), (15b), and (16). we obtain:
E. = a
(17)
f
+ -.
f‘-f
Go
The data listed in Column 6 of Table 5, show that
f/(f’-f) is in rough approximation constant for several
aromatic hydrocarbons for which f and f‘ have been
calculated theoretically [26]. Hence, from Equation (12),
the ionization energy will be, in general:
b) by basing the calculation o n the 0,O-frequency Voo, provided
the latter can be identified [25].
Matsen [26] selected for CO the frequency of the longest
wavelength absorption maximum (Table 8, Column 6),
as recommended by Clar [27], and obtained
(19)
1
2
3
No.
I [261
Eq. (I5a)
-f‘ 1261
“I
Eq. (15b)
5
f
E‘-f
6
1.000
0.618
0.414
0.167
0.642
0.887
1.333
3.731
3.462
1.37
1.45
1.45
4
3.369
3.232
3.713
1.42
6
0.275
0.208
0.526
1.40
0.220
0.605
1.117
0.696
7
10
11
0.406
D.584
0.444
3.510
j.825
3.570
1.44
1.42
1.43
0.452
0.684
0.499
0.870
0.648
0.538
14
16
0.460
0.423
3.598
3.587
1.44
1.42
0.520
0.473
0.802
0.864
19
0.497
0.438
3.662
3.561
1.42
1.44
0.49 1
0.846
3.574
3.49
3.38
1.44
1.45
1.46
0.501
0.836
0.347
0.983
0.439
0.539
0.704
0.898
0.502
0.580
0.445
0.497
0.371
0.868
0.833
0.959
51
0.398
0.442
0.943
0.899
65
66
67
0.477
0.662
0.592
0.783
0.660
0.730
68
69
71
57
0.635
0.637
0.405
0.400
0.649
0.647
0.932
0.947
89
90
91
0.405
0.618
0.520
0.945
0.695
0.793
92
93
94
0.505
0.650
0.630
0.808
0.680
0.702
95
96
0.600
0.500
0.730
0.830
97
98
0.410
0.230
0.937
1.136
21
22
23
24
25
26
47
48
49
50
0.446
0.394
0.320
0.475
0.599
3.632
3.855
1.46
1.41
0.401
0.406
3.501
3.501
1.44
1.50
+ 0.857 CO [eV] (Fig. 2, 11)
I
=
5.156
+ 0.775 GO [eV] (Fig. 2, 111)
Considering the ionization energies (which have been
determined with some certainty for only relatively few
aromatic hydrocarbons), cf. Table 8, Column 4, and
using the 10frequencies (Table 8, Column 6), we obtain:
Eq. (1 3)
0.800
0.535
0.375
20
4.39
--K
fH
1
2
3
5
=
Becker and Wentworth [28] based their calculation on
the 0,O-frequency and obtained:
(20)
Table 5. Corrected and uncorrected Huckel coefficients and values
of -K for some aromatic hydrocarbons
I
(21‘
3
I
=
5.11
+ 0.701 Go [eV] (Fig. 2, 1)
35
Fig. 2. Experimentally determined ionization energies Iexp (Table 8,
Column 4) of some aromatic hydrocarbons as a function of their first
excitation energy GO (Table 8, Column 6); the numbers refer t o the
hydrocarbons listed in Table I . Differing ionization energy values were
obtained for Nos. 2, 3, 4, and 6 by various authors using different
methods. I: Equation (21); 11: Equation (19); 111: Equation (20).
The ionization energies calculated from Equation (21)
are listed in Column 5 of Table 8.
Substitution of Equation (18) into Equation (10) gives
(since I = -Ei):
JCT = YGO + C-EA + E
(22)
Considering EDA complexes of the same bond type [lo]
between an acceptor and analogous donors (e.g.,
angularly condensed, unsubstituted aromatics), so that
E = -4.3 eV [29] may be considered as nearly constant,
we obtain :
JCT = YO” + C’
(23)
where C’ = C-EA
[*I The numbers refer to the hydrocarbons listed in Table
I
Equation (18) can be fitted t o experimental data only with
limited accuracy, because ionization energies are known to
within a n accuracy of 0.1 -0.2 eV for only a few aromatic
hydrocarbons. Moreover, often th e first excitation energy of
aromatic hydrocarbons cannot be determined with great
accuracy. This energy is determined a) by taking t h e maxim u m of the longest wavelength absorption b an d as Qo, an d
622
15
1
+E
-~
[25] The 0,O-frequency is the excitation energy required to raise
an electron from its zero vibrational energy level in the ground
state to the zero vibrational level in the excited state.
[26! See, among others, F. A . Muneiz, J. chem. Physics 24, 602
(1956); G. W. Wheland, J. Amer. chem. SOC.63,2025 (1941).
[27] E. Clurr Aromatische Kohlenwasserstoffe, Springer, Berlin
1941.
[28] R. S . Becker and W. E. Wentworth, J. Amer. chem. SOC.85,
4263 (1963).
Angew. Chem. internat. Edit.
Vol. 3 (1964) 1 No. 9
Substitution of Equations (18) and (21) into (23) reduces
the latter to:
Gcr
(23’)
=
0.701 ’70
+ 0.81-EA
[eV]
~ 9 1
The calculation of electron affinities from Equation
(23’) is limited by the inaccuracy in the values of I
obtained from Equation (18) or (21). The uncertainties
in the values of Iexp and GO or q00 are strongly reflected
in those of the electron affinities (variations of 0.2 to
0.3 eV). The electron affinities of acceptors in EDA
complexes with aromatic hydrocarbons of known excitation frequency GO can be calculated using Equation
(23’) provided the values of FCT for the EDA complexes
are known. The electron affinities of a given acceptor in
EDA complexes with various donors (Table 6, Column
6a) were calculated using Equation (23’) and the average
values are listed in Column 6 of Table 6.
Figure 3 shows a plot of the electron-transfer energies
FcT I.”. the first excitation energies GO of a series of
aromatic hydrocarbons for EDA complexes of 1,3,5trinitrobenzene, chloranil, and tetracyanoethylene. The
values given in parentheses in Equations (24a)-(24c)
correspond to the experimental straight lines shown in
Figure 3.
The deviations of the values given in parentheses from
those expected from Equation (23’) fall within a range
of inaccuracy of about 0.2 eV.
The averaged electron affinities calculated by different
methods, viz. according to Equations (9,
(7), (14), (23’),
and (25), are collected together in Table 6, Column 8.
For EDA complexes of a given acceptor (EA = const.)
with analogous donors (E = const.), it follows from
Equation (22) that:
GcT-SO
(23”)
(23”a)
C”
=
C-EA
=
(y-1) Go
+ C”
+ E = 0.81-EA
and from (21) that:
(23”’)
G0-GcT
=
0.299 90-0.81
+ EA
[eV]
Thus, although we have complexes of one and the
same acceptor, i.e. EA = constant, (+JCT) is not constant, as is to be anticipated if Hiickel’s simple relationships [Equations (15a) and (15b); f = f’] were valid;
instead, (Go-Gcr)
decreases as the wavelength of the
first excitation energy of the donor increases. The
empirical approximate constancy of F0-FcT [6, p. 741 is
valid only for aromatic hydrocarbons of nearly equal 5,)
within the limit of accuracy (about
0.2 eV) of Equation (23”’).
5 . Electron Affinities Relative to Chloranil
If we assume that for EDA complexes of the same bond
type [lo] and configuration [6,13], e . g . n,Tt-complexes
E may be considered to be nearly constant - a n assumption that is approximately correct then from Equation
(8) we obtain the following expression for two acceptors, relative to one and the same donor:
I
-
(25)
( ~ C T ) I - ( ~ C T ) k=
EAk-EA;
Hence, after normalization to a given standard acceptor,
Equation (25) can be used to calculate the electron
affinities of other acceptors of the same type [30]. In
Column 7 of Table 6, a list is given of electron affinities
relative to chloranil (EA = 1.37 eV) calculated using
Equation (25) .
25
rn
35
3
L
45
$[eVl-
Fig. 3. Relationship between the electron transfer energy ’ ~ C T
and the
for a few EDA complexes of the accepfirst donor excitation energy
tors tetracyanoethylene (I), chloranil (II), and trinitrobenzene (111).
The numbers refer to the hydrocarbons (donors) listed in Table 1 .
(24a)
Tetracyanoethylene:
(EA = 1.8 eV) qcT = 0.701 GO-0.99 (0.72) [eVj
(24b)
Chloranil:
(EA = 1.37 eV) JcT
=
0.701 Fo-0.56 (0.37) [eV]
I , 3,S-Tr initro benzene :
(EA = 0.7 eV) G,,
= 0.701 GO-0.10 (0.3) [eV]
(24c)
-
-~~
[29] For halogens and XY-type interhalogen compounds, we
must set E = -3.5 eV, in accordance with the vertical model
(p. 619). Equation (23’) then becomes
?CT = 0.701 50 + 1.61-EA
Angew. Chem. internat. Edit. 1 Vol. 3 (1964) No. 9
The agreement between electron affinities calculated
from Equation (25) and those obtained from Equations
(5),(7),(14), and (23’) in Columiis 3,4,5, and 6 of Table 6,
respectively, is satisfactory. The accuracy of the absolute
electron affinity values is certainly not better than about
0.2 eV. Considering this error, their relative order should
be correct.
[30] M . Butley and L. E. Lyons, Nature (London) 196,573 (1962)
use iodine a s acceptor standard. With iodine (EA = 1.8 eV) and
Equation (25), they calculate the values of EA for x-acceptors
(quinones, trinitrobenzene, tetracyanoethylene, etc). However,
since it is known that EDA complexes of iodine (which is a aacceptor) have an entirely different configuration than the corresponding complexes of x-acceptors (trinitrobenzene, chloranil,
etc.) [6], it is certain that the values of E [Equation (8)] for the
iodine complexes are different from those for the x-acceptors (cf.
the values of E given o n p. 621 for the parallel and vertical model
of iodine complexes with aromatic hydrocarbons). The values of
E for EDA complexes of iodine with aromatic hydrocarbons are
aboutO.S-1.0 eVIess than the values for n-acceptors. Correspondingly, the electron affinities (relative to iodine) of rr-acceptors
calculated by Butley and Lyons should be too high by the same
amount.
623
-- - 0
P ?
0
'9
m o
22
? P ? ?
- 0 0 0
L
lli
0
0 -
2nc
m
y1
b
n
8
c
g
-v
W
U
.-
c
P
L
0
rd
u
-0
m
L
lli
x
0
n
F
z
W
w
W
5
c-"
624
u
- 4
0
Z
m o - N
W
b
b
b
Angew. Chem. internat. Edit. / Vol. 3 (1964) / No. 9
Dibromopyromellitic dianhydride
5-Hydroxy-1,4-naphthoyuinone
Pyromellitic dianhydride
Diacetyl
Ethyl pyruvate
Diethyl maleate
Diethyl fumarate
Dimethyl oxalate
Tetrachlorophthalic anhydride
Maleic anhydride
Acctamide
Acetylacetone
Ethyl acetoacetate
Ethyl acrylate
Phthalic anhydride
Carbon tetrachloride
Hexachlorobenzene
Chloroform
57
59
60
61
62
63
64
65
66
67
68
69
70
71
72
73
74
0.59
0.40
0.39
0.54
3.16
0.00
0.78
2.57
3.00
0.95
0.42
0.34
0.32
0.25
0.65
0.58
0.57
1.08
-0.15
-1.00
3.84
3.00
2.90
0.08
0.60
0.60
0.60
0.59
0.59
3.30
3.23
3.23
3.18
3.10
- --
0.40
0.58
0.54
2.59
[**I According to P. R. Hammotrd, J. chern. SOC.(London) 1964, 471,
the following electron affinities are obtained for monosubstituted pbenzoquinones from the FCTvalues of the EDA-complexes with
hexamethylbenzene using Equation (25): NOrbenzoquinone: 1.45;
CN-benzoquinone: 1.04; F-benzoquinone: 0.91 ; CI-benzoquinone:
0.97; I-benzoquinone: 0.95; CH,-benzoquinone: 0.75. These values
correspond t o those given in Table 6.
['I For key numbers t o donors, see Table 1.
58
Nitrobenzene
1,2-Dinitrobenzene
2.4.6-Trinitrotoluene
2,4,6-Trinitro-m-xylene
2,4-Dinitrotoluene
Dinitrophenol
Chlorobenzene
Broniobenzene
Iodohenzene
Hexachloro benzene
Azobenzene
1,4-Dicyanobenzene
1,3,5-Tricyanobenzene
1,2,4,5-Tetracyanobenrene
43
44
45
46
47
48
49
50
51
52
53
54
55
56
0.6
50, 51, 65-69,71.
3.71
3.65
0. I
0.84
0.86
1.09
[i]
10, 11. 14, 16, 21, 23, 25, 47, 49-51,
No. 2, 3, 6, 7, 10, 14, 16, 21. 23, 25, 47-49, 71.
[h] No. 2, 3, 5-7,
71.
Is] No. 2, 3, 6, 7, 10, 11, 14, 16, 21, 23, 25, 47, 49, 51, 71.
[r] No. 2, 3, 6, 26, 47.
[q] No. 1-4. 6, 7, 10, 14, 16, 21, 23, 25, 47, 49, 72-75, 77, 79--54, 99.
6, 7, 10, 14, 16, 21, 23, 25, 47, 49, 68.
[p] No. 2-4,
No. 2, 3, 6, 23, 47, 49.
101 No. 1-3, 15, 23, 47, 49, 72-74, 77, 80-82, 84, 99.
[€I No. 2-4, 6, 7, 10, 14, 16, 21, 23, 25, 47, 49.
[g]
[n] No. 1-4, 6, 7, 10, 14, 16, 21, 23, 25, 26, 47, 49, 72-86, 99.
[el No. 2, 3, 6, 10, 25, 47.
[k] No. 2, 3, 6, 7, 10, 14, 16, 21, 23, 25, 26, 47, 49, 72-77.
[I] No. 73-75, 77.
f1.31
t8.60
f0.8J
t0.32
t0.52
t0.55
[m] No. 2, 3, 6, 10, 25, 26. 47, 85.
65.68, 69.71, 89-98.
0.15
0.57
0.41
0.3,
0.3
0.25
0.58
0.6
0.6
0.6
0.6
0.6
0.85
1.16
0.9
0.1
0.4
t O
0.5
.'- 0
0.6
0.4
0.1
0.6
0.4
0.4
0.57
0.6
0.4
[dl No. 2, 3, 6, 7, 10, 14, 16, 21, 23, 25, 47, 49.
21, 23, 25,47-51,
41.77
[rl
47
47.77
86
47.77
86
86
86
47
86
86
86
86
[cl No. 1-3, 6. 7, 10, 11, 14, 16, 20, 21, 23, 47, 49, 51, 69.
[bl No. 2, 3, 5-7,10,11,14,16,
[a] No.l,2,6,7,10,11,14,16, 20, 21,23,25,26,47-49,
0
0.56
0.57
0.6
0.9
0.8s
1.16
0.1
0.4
+ 0
0.57
0.6
0.4
0. I
-0
-
6. Electron Affinities from Polarographic Half-Wave
Reduction Potentials
The first step of reversible polarographic reduction of
an electron acceptor or of reversible oxidation of an
electron donor proceeds as follows:
= -3.66 eV, obtained as an average value from the
~ ~6, Column
~ ~ lo), we obtain:
values of A F (Table
EA
(28)
=
-ERed
‘/2
+ 1.41
Figure 4 shows a plot of EA against E;Zd [34]; the
agreement with Equation (28) is satisfactory. The
straight line in Figure 4 corresponds to Equation (28‘).
(28’)
EA
=
-1.04 EtZd
+ 1.39 [eV]
MuccuZ[31] was the first to show tinat, for angularly
condensed aromatic hydrocarbons, there is a relationship between the polarographic one-electron half-wave
reduction potential, ECZd, and the energy of the
highest unoccupied orbital, Ei.
The reduction potential, measured against a saturated
calomel electrode, is given by the expression :
Fig. 4. Linear relationship between the electron affinity
and the
typical
half-wave reduction potential
according to Equation (28’); for
acceptors in zcetonitrile and dimethylformamide as
the numerical values and numbering of the acceptors, see Table 6.
solvent. The half-wave potentials are largely independent of the solvent. Usine EA’ and EA (Table 6,
T h e average value of AFsoiv (-3.66 eV) for t h e acceptors in
Column 6) in Equation (27), it is possible to calculate
Table 6 is of t h e correct magnitude, a s indicated by comparithe difference in solvation energy of the anion and of
son with the solvation energies t h e anions o f aromatic hydrothe neutral acceptor molecule A F (Table
~ ~6, Column
~ ~
carbons (Table 8, Columns 8 an d 9). T h e values of AFsolv for
lo). SinceA-FsoIv AFsolV,the major part of A F ~ t h e~acceptors
~ ~(Table 6) are relatively constant; this is probably d u e t o t h e fact that t h e negative charge of the A- ion
is represented by the solvation energy of the anion.
AFsolv m A-FsolV
is nearly constant [35]. For h-Fsolv is located in certain parts of t h e molecule (e.g. -C=O, -C=N,
Peuvdr [34] measured ECZd for a number of
>
.-
[31] A. Mnccoll, Nature (London) 163, 178 (1949); L . E. Lyons,
ibid. 166, 193 (1950). In addition: A. Pullman, E . Pullman, and G.
Berthier, Bull. SOC.chim. France 591 (1950); G . J. Hoijtink, Recueil Trav. chim. Pays-Bas 74, 1525 (1955); G. J . Hoijrink, E. de
Boer, P . H . van der Meij, and W . P . Weijlnnd, ibid. 75,487 (1956);
P . Balk, S . de Bruyn, and G. J. Hoijtink, ibid. 76, 860 (1957); P.
H . Given, Nature (London) 181, 1001 (1958); P . H . Given and
M . E. Peover: Advances in Polarography. Pergamon Press, London 1960, p. 948; M . Eatley and L. E. Lyons, Nature (London)
196,573 (1962); M . E. Peover, J . chem. SOC.(London) 1962, 4540.
[32] E . A . Matsen, J. chem. Physics 24, 602 (1956); R . M . Hedges
and F. A . Matsen, ibid. 28, 950 (1958); G. J . Hoijtink and J. van
Seliootan, Recueil Trav. chim. Pays-Bas 71, 1089 (1952); A. Lainlinen and S. Warzonek, J. Amer. chem. SOC.64, 1765 (1942).
[33] The electrochemical process:
Asolv
+ C l ~ o +~ vHqiq
+
AFolv
7
(HgCl)solv:(E~~d)
-NO2 groups), whereas in the anions of polycyclic hydrocarbons, the negative charge is dispersed over larger portions
of t h e molecule. Therefore, in these cases A-Fsolv is not
constant.
Since Equation (28) is fulfilled satisfactorily, Equation
(28’) can be used to estimate the electron affinities of
acceptors for which only Et:d has been measured.
Electron affinities calculated in this manner from
Equation (28’) are listed for a few acceptors in Column
1 1 of Table 6.
According to Equations (27) and (8), there must be
some relationship between CCT and EZ
:d
[35] when
CcT is based on EDA complexes of different acceptors
with one and the same donor:
can be separated into the following steps
Agas + e-
+
~ s o l v+ Agas
Hqiq
+
+ Cl&lV
(29)
: (-EA)
A&
gas + ~
+ (HgCI),,IV
o
:~( ~ ~
l vA Fs o ~ v - A ~ s o ~ v )
v s o=
+ e- (in Hg)
: ( A F H ~ ~ c I=
,
e- (in Hg) + Hgliq
+ e-
E
H
~
~= c0.53
~ ~eV)
: (4.54 eV)
According to Equation (26a), negative values of ECZd should be
used for strong acceptors, and positive values for weak acceptors
(the sign is sometimes given in the opposite sense). AFsolv is
negative, EA is positive by convention. It is assumed that electron
addition a t a dropping mercury electrode according to Equation
(26a) is sufficiently fast and reversible so that overvoltage phenomena and secondary reactions can be neglected. Moreover, no
proton addition should take place (solvents devoid of proton
activity, such as acetonitrile, dimethylformamide, or dioxane,
must be used).
1341 M . E. Peover, Nature (London) 191,702 (1961); Trans. Faraday SOC.58, 656, 2370 (1962).
[35] M . E. Peover, Trans. Faraday SOC.58, 1656 (1962).
626
ECZd = GCT-(I
+ E-AFsoIv-5.07)
If the ionization energies of pyrene, hexamethylbenzene,
and tetramethyl-p-phenylenediamineare taken as 7.8,
8.0, and 6.7 eV, respectively, and E = -4.3 eV is used
for the Coulomb energy of the EDA complexes of the
acceptors listed in Table 6 (for complexes with tetramethyl-p-phenylenediamine E = -3.9 eV), the following
linear relationships result for the different acceptors,
when A F in ~Equation
~ ~ (29)
~ is assumed to be -3.66
eV :
(30)
Pyrene: Eczd
=
GcT-2.09 (2.13) [eV]
Hexamethylbenzene: EcZd = C,,-2.29
Tetramethyl-p-phenylenediamine:Er)Zd
(2.40) [eV]
=
CcT-1.35
(1.365) [eVl
Angew. Chem. internat. Edit. Vol. 3 (1964)
I No. 9
The numerical values refer to straight lines plotted in
Figure 5 which best fit the experimental Ef)zd and ~ C T .
values. Because of the uncertainty in E and the
variability of A F ~ deviations
~ ~ ~ , are not of great import an ce .
tron-transfer energy GCT and the electron affinity EA
of the acceptor. Applying the ionization energies and
values of E given above for pyrene, hexamethylbenzene,
and tetramethyl-p-phenylenediamine, Equation (8) gives
the following linear relationships :
(33)
7'
t
Pyrene: GCT
Hexamethylbenzene: qCT
=
-EA
=
-EA
+ 3.51 (3.42) [eV]
+ 3.7 (3.75) [eV]
Tetramethyl-p-phenylenediamine :
JCT = -EA
+ 2.8 (2.78) [eV]
3
--7=25
The straight lines in Figure 6 refer to the numerical
/=la
values given in parentheses. In view of the relatively
large inaccuracy in the EA values, the agreement
between Equations (33) and (8) is satisfactory.
Fig. 5. Correlation between the half-wave reduction potential of the
acceptor Et:d
[34,35] and the electron transfer encrgy ?CT f o r donor
acczptor complexes, according t o Equations (29) and (30). Donors:
tetramethyl-p-phenylenediamine (I), hexamethylbenzene (II), and
pyrene (111). The numbers on the experimental points refer t o acceptor
numbers given in Table 6.
7. Electron Affinities from Polarographic HaIf-Wave
Oxidation Potentials
According to Equation (26bj, polarographic oxidation
of aromatic unsubstituted and substituted hydrocarbons
may be described in terms of a relationship corresponding to Equation (28) in which the first polarographic half-wave oxidation potential is a function of
the ionization energy of the hydrocarbon. Pysh and
Yung [36] found empirically:
(31)
1 = 1.473 EG:
+ 5.82
From Equations (31) and (8), using E
obtain:
(32)
EA
=
-GcT
=
--4.3 eV, we
+ 1.473 E: + 1.52
The values of :E
:
[36] and VCT for E D A complexes between the donors listed in Column 12a and the acceptors
listed in Column 1 of Table 6 have been determined experimentally. Thus, for the donors in Column 12a, the electron
affinities can be caiculated and averaged using Equation (32)
[see Column 121. However, the electron affinities obtained
from Equation (32) are somewhat low compared to those
obtained and averaged using Equations (23') and (25) [see
Table 6 , Column 81.
8. Correlation between Electron Affinity and ElectronTransfer Energy
It follows from Equation (8) that, for complexes of one
and the same donor with various acceptors, an approximately linear relationship must exist between the elec[ 3 6 ] E . S . Pysk and N . C. Yung, J. Amer. chem. SOC.8.5, 2124
(1963).
Angew. Chem. inrernat. Edit. 1 Vol. 3 (1964) 1 No. 9
Fig. 6. Electron-transfer energy SCT of E D A complexes as a function
of the electron affinities F A of various acceptors, according t o Equation (33). Donors: tetramethyl-p-phenylenediamine(I), pyrene (II), and
hexamethylbenzene (111). For the
values and acceptor numbering,
see Table 6.
9. Electron Affinity and Chemical Constitution
The relatively great inaccuracy in the electron-affinity
values precludes discussion of the more refined effects
induced by the structure of the compounds involved.
The strongest acceptors are 2,3-dichloro-5,6-dicyano-pbenzoquinone, tetracyanoethylene, 2,6-dinitrobenzoquinone, tetracyano-p-benzoquinone,
2,3-dicyanobenzoquinone, and bis(dicyanomethy1ene)quinones. The electron affinities of unsubstituted quinones lie in the following order: 9,lO-anthraquinone < I ,4-naphthoquinone *
9,lO-phenanthrenequinoneM 1,2-naphthoquinone w pbenzoquinone < o-benzoquinone < 1 &diphenoq u i none.
Alkyl substituents tend to reduce the electron affinity,
the effect becoming more pronounced as their number
increzses. Electrophilic (halogeno, CN, NOz) groups
627
tend to augment the electron affinity of a compound,
this effect increasing with the number of such substituents. C N and NO2 tend to increase the electron
affinity to a larger extent than do halogens. C1, Br, and 1
do not differ appreciably in this respect. The nitroaromatics are relatively good acceptors : 2,4,7-trinitrofluorenone > 1,3,5-trinitrobenzene KY 1,4-dinitrobenzene > nitrobenzene > 1,3-dinitrobenzene > 1,2dinitrobenzene. With the exception of 1,2-dinitrobenzene, the electron affinity increases with increasing
number of nitro groups. Unexpectedly, the electron
affinity of 1,3-dinitrobenzene is less than that of nitrobenzene. (Since the electron affinity of these two nitro
compounds has been determined only once, and by
two different methods, this difference may be due to
the difference in the methods used.) The fact that the
electron affinity of 1,2-dinitrobenzene (in contrast to
that of I ,3- and 1,4-dinitrobenzene) is practically zero
is consistent with steric inhibition of the mesomeric
stabilization of the 1,2-dinitrobenzene anion, i. e., both
nitro groups cannot become coplanar with the benzene
ring. Substitution by methyl generally decreases the
electron affinity: 1,3,5-trinitrobenzene > 1,3,5-trinitrotoluene > 1,3,5-trinitro-rr;-xylene;
also, 1,3-dinitrobenzene > 2,4-dinitrotoluene. The polycyanobenzenes are
definitely poorer acceptors than the corresponding
nitrobenzenes. The electron affinity of polycyanobenzenes increases with increasing number of cyano groups.
Simply substituted halogenobenzenes are relatively good
acceptors. An increase in the number of halogen s u b
stituents on the benzene ring does not seem to exert as
strong an influence as does substitution by N02. This is
exemplified by hexachlorobenzene.
Molecules with two adjacent electrophilic carbonyl
groups, e.g. diacetyl, ethyl pyruvate, and dimethyl
oxaloacetate, are relatively good electron acceptors
(comparable to trinitrobenzene). The electron affinity
of these compounds and that of the electrophilic
-CO-C-C-COgroup, e.g., in diethyl maleate and
diethyl fumarate, are nearly equal.
Compounds with F-dicarbonyl groups have lower electron affinities (acetylacetone, ethyl acetoacetate).
Phthalic anhydride is practically devoid of electronacceptor properties. It acquires such properties by
electrophilic substitution (tetrachlorophthalic anhydride).
The relatively high electron affinity of C C 4 is noteworthy. This is consistent with the fact that, in the
presence of CC14, donors such as hexamethylbenzene
show a characteristic absorption band indicative of
electron-donor/electron-acceptorinteraction [36a]. Carbon tetrachloride and chloroform are o-acceptors
which, despite their high electron affinity, have an EDA
interaction energy that seems to be lower than that
of Ti-acceptors with comparable electron affinities
[36d].
~
[36a] D . P. Stevenson and G . M . Coppinger, J. Amer. chem. SOC.
84, 149 (1962); M . Templetfon, J. chem: Physics 37, i61 (1962);
F. Dorr and G . Buttgereit, Ber. Bunsenges. physik. Chem. 67, 867
(1963); R. Anderson and J . M . Pramnitr, J . chem. Physics 39,
1225 (1963).
628
11. Aromatic Hydrocarbons as Electron Acceptors
1. Quantum-Mechanical Calculation of
Electron Affinity
According to Pople and Brickstock [37], the energy of
the highest occupied and that of the lowest unoccupied
orbital of an aromatic hydrocarbon are:
F [38] can be calculated by the semiempirical ASMOSC
[42] method. The electron affinities calculated by Pople
and Brickstock [37] and by Hedges and Matsen [39]
from Equation (34b) are listed in Columns 3 and 4 of
Table 7. sc' was normalized by Hedges and Matsen to
the ionization energy of naphthalene (8.12 eV) to x' =
-3.87 eV.
These calculations [37,39] were carried out by Hoyland
and Goodman [40] with even fewer assumptions; they
also took into account the chanses in electron density
following removal or capture of an electron, as well as
the effect of theo-electron energy. The electron affinities
calculated by these workers [40] are listed in Column 5
of Table 6.
The w-approximation in the simplified treatment
proposed by Ehrenson [22] gives an expression for the
electron affinity corresponding to Equation (13):
(34')
-E.J
=
EA
=
-a
o-(f,-
'
7
Q) p = -mo- %p
+
As already mentioned, the o-approximation [Equation
(13)] gives satisfactory values for the ionization energies
of aromatic hydrocarbons when sc = -9.878 eV, 9 =
-2.1 1 eV, and w = 1.4. On the other hand, the electron
affinities obtained from Equation (34) with w = 1.4 are
much too high. To obtain good agreement with the
electron affinities calculated from Equation (34') by
Hedges and Matsen [39], Poples [37], and Hoyland and
Goodmart [40], as well as with the experimental values
of Wentworth and Becker [3], the very high value of
[37] J . A . Pople, Trans. Faraday SOC. 49, 375 (1953); A. Brickstock and J . A . Pople, ibid. 50, 901 (1954); J. A. Pople, J . physic.
Chem. 61, 6 (1957); C. J. Koothaan, Rev. mod. Physics 23, 61
(1951).
N
ywy=Coulomb interaction between one electron a t atom and
one electron at atom V. Pwv= resonance integral for two adjacent
atoms I* and v. N = number of occupied orbitals in the aromatic
hydrocarbon. X i w and Xi, = Huckel coefficients of the ri-th and
,i-th atomic orbitals in the i-th molecular orbital. The %-values
can be calculated with good approximation using the simple
MO theory.
[39] R. M. Hedges and F. A. Matsen, J. chem. Physics 28, 950
(1958). See also: A. Streitwieser and P. hf. Nuir, Tetrahedron 5,
149 (1959); A . Streitwieser, J. Amer. chem. SOC.82, 4123 (1960);
W.F. Wolf, Abstracts J. Amer. chem. SOC.Meeting, New York
1960.
[40) J . R. Hoyland and L. Goodman, J. chem. Physics 36, 12, 21
( 1962).
Angew. Chem. internat. Edit. / Vol. 3 (1964) No. 9
= 3.77 must be used in Equation (34’) [43,44];
similarly, a value of o = 3.73 must be used for phenanthrene, benzo[c]phenanthrene, and triphenylene [44]. No
theoretical explanation is available for the fact that the
calculations of I using Equation (13) and of EA using
Equation (34’) require different values of w.
In Column 6 of Table 7, a list is given of the electron
affinities calculated by Ehrenson [22] according to the
w-method, using a w = 3.5. Column 2 contains the
electron affinities obtained for a few aromatic hydrocarbons by Wentworth and Becker [3] using the electron
absorption coefficients determined experimentally by
Lovelock [l].
2. Semiempirical Methods for Determining
Electron Affinities
w
Table 7. Electron affinities [eV] derived for aromatic hydrocarbons by
various methods.
a) Electron Affinity and the Longest- Wavelength
Absorption Band of the Donor
It follows from (34) that
(EA
(35)
+ 1)/2
=
CL’
On the basis of the electron affinities determined experimentally for a few aromatic hydrocarbons (Table 7,
Column 2) and the experimental values for the first
ionization energies Iexp (Table 8, Column 4), one
obtains in agreement with Becker and Wentworth [28]
an average value of 8.14 eV for the sum (I + EA). From
Equations (21) and ( 3 9 , we obtain for CY’= 4.07:
6
(36)
Theor. [411
2
1
3
I
0.42
-=
1-1.63
1-1.40
1-1.40
I
I
I
I
1
I
0.49
-0.14
0.64
-0.20
0.61
-1.590
I
I
-0.246
0.147
4
1.035
5
-6
0.20
7
I 0.46
10
0.14
-0.06
1 1
I
0.62
-0.28
--
1
0.25
1
I
0.014
1
11
17
=
3.03-0.701
<”
The electron affinities calculated from ijo (Table 8,
Column 6) using Equation (36) are given in Column 2
of Table 8.
The differences between the electron affinities calculated
by the ASMOSC method (Table 6, Columns 3, 4, 5,
and 6), those obtained from 30 (Table 8, Column 2)
using Equation (36), and the few experimental EA
values
(Table 7, Column 2) may be considerable. The
EA values obtained by averaging all EA values are
given in Column 3 of Table 8.
0.446
1 1
0.75
-~
19
-0.033
EA
0.33
201
-0.14
1iG
1
I
I
I
1I
1
I
I
0.313
t
e?
0.463
0.442
0.727
23
24
0.573
0.50
0.385
-0.78
-0.290
-0.37
27
-
28
I
-0.41
29
-0.35
-- 47
0.39
0.68
-~
48
--
0.55
0.410
I
49
50
’
-~
I-
51
65
0.417
0.92
I
0.676
0.683
0.590
0.052
[41] S. E. Ehrenson, J. physic. Chem. 66, 706, 712 (1961).
[42] Antisymmetrized Molecular Orbital Self-consistent.
[43] S. E. Ehrenson [41] uses w = 3.5.
[44] D . R . Scott and R . S . Becker, J. physic. Chem. 66, 2713
(1962).
Angew. Chem. internat. Edit. 1 Vol. 3 11964) 1 No. 9
Fig. 7. Averaged electron affinities of aromatic hydrocarbons (calculated by different methods) as a function of the first excitation energy
Go. The straight line corresponds to Equation (37); for numerical data
and key numbers to hydrocarbons, see Tables 8 and 7, respectively.
Their absolute accuracy is never better than f 0.2 eV.
Figure 7 shows a plot of these electron affinities against
$0 for aromatic hydrocarbons. The linear relationship
(37)
-
EA = 2.991-0.700 <o
obtained in this manner is in good agreement with
Equation (36), which was derived directly from experimental data, and with the expression of Becker and
Wenrworth [28] :
(38)
EA = 2.92-0.762 Goo
629
It is noteworthy that the linear functions (21) and (36),
which agree best with the experimental values for
ionization energies and electron affinities, have the
s a m e gradient.
I
!
which is in good agreement with Equation (42). The
A F values
~ ~derived
~ from
~
Equation (44) are given in
Table 8, Column 9.
The decrease in solvation energy with increasing electron
affinity of aromatic hydrocarbons is possibly associated with
the distribution of the negative charge over the entire molecule. The solvation energy decreases, and the increase in
electron repulsion energy o n transition from the neutral
molecule to the anion also becomes less (Le., the electron
affinity of the aromatics increases) as the range of dispersion
of the negative charge over the aromatic anion increases. Also
the interaction energy with the solvent molecules decreases
with increasing radius of the aromatic anion.
Fig. 8. Half-wave reduction potentials ERed [47Jof aromatic hydro-
%
carbons as a function of avera_eed electron affinities
calculated by
different methods. The solid straight line corresponds to Equation (39).
For numerical data and key numbers to hydrocarbons, see Tables 8
and 1, respectively.
b) Electron Affinity, Hal'
Solvation Energy
Wave Reduction Potential and
Figure 8 shows a plot of ECId (Table 8, Colunm 7 [45])
obtained from Equation (27) vs. EA (Table 8, Column
3). The EtId and EA values fit Equation (39):
(39)
E t r d = 2.416-0.692
It follows from Equations (39) and (27) that there is a
functional relationship between AFsolvand EA. If we set
(40)
AFsolv = -Tsolv
+
6.EA
then from Equations (27) and (40) we obtain:
(41)
-ERed
'12
=
EA (1-6)
+ qsOlv-5.07
[eV]
and, by comparison with Equation (39):
6 = 0.308 and qsOlv= 2.654
Hence:
(42)
AFsolv = 0.308 EA-2.654
Equation (42) can be checked by using Equation (27)
to calculate A F ~ (~Rt;,' ~ from Column 7,
from
Column 3 of Table 8).
(43)
AFsolv = EA-5.07
+ ECrd = EA-EA'
The values of AFsolv calculated using Equation (43)
are listed in Column 8 of Table 8, and give rise to the
following linear relationship :
(44)
AFsolv = 0.2714 m-2.624 [eV]
[45] I. Bergman, Trans. Faraday SOC.50, 829 (1954); G. J.
Hoijtink, Recueil Trav. chim. Pays-Bas 73, 355 (1954); 74, 1525
(19551, where references to previous publications are given. The
Etzd-values of Eergman must be increased by 0.5 eV in order to
obtain the EtId-values measured in 7 5 % dioxane against a
saturated Hg2CIz-electrode. The E:zd-values measured in 75 %
and 96 % dioxane show practically no differences.
630
Fig. 9. Solvation energy AFsolv of aromatic hydrocarbons as a function
of the electron affinity
according to Equation (44); for numerical
data and the key numbers of the aromatic hydrocarbons, see Tables 8
and 1, respectively.
Extrapolation of Equations (21) and (37) to GO + O is of
questionable validity because these equations have been
shown to be valid only within the rather narrow range from
90 = 2.5 to 4 eV. Extrapolation to GO + 0 would assume I
and EA to be linear functions of CO [Equations (21) and (37)]
even for very low values of 90;this, however, has never been
proved. Though the difference (ILEA) = (1.4 90 2.12 [eV])
[Equations (21) and (37)] decreases with decrease in the
excitation energy of aromatic hydrocarbons, Equations (21)
and (37) give different values, viz. Iij,=o = 5.1 1 eV and EAj,=o
= 2.99 eV. However, according to Huckel's theory, VO should
tend to zero and I should tend to EA for infinitely large, conjugated aromatic ring systems, e.g. in graphite. The work
function of graphite is 4.39 eV [46].
For hydrocarbons, we should expect IG,+O # EA since the
electron repulsion energy of the hydrocarbon anion is greater
than that of the hydrocarbon cation. Accordingly, I;,+o >
EAij,,o.
Finally, it should be borne in mind that Equations
(21) and (37) refer to two-dimensional aromatic ring systems
in the gaseous state.
+
3. Electron Affinity and Chemical Constitution [ "1
a) The electron affinity increases with the number of rings
fused together in a line, whereas the ionization energy decreases correspondingly (see Table 1, Nos. 1-5, 6-9, 10- 13,
and 14-18).
[46] A . Eraun and G. Eusch, Helv. physica Acta 20, 33 (1947).
[*I The fact that aromatic hydrocarbons are electron donors as
well as electron acceptors enables them to undergo intermolecular
EDA-interaction. Wentworth and Chen proved this using the
spectral shift of the absorption spectra or a few polycyclic hydrocarbons on addition of hexamethylbenzene in heptane solution:
J. physic. Chem. 67, 2201 (1963).
Angew. Chem. internat. Edit.
Vol. 3 (1964) 1 No. 9
Table 8. Energy data for aromatic hydrocarbons.
1
2
3
Hydrocarbon no.
[*I
EA
Eq. (36)
EA
Iexp
-1.5
9.24 [47]
-0.3
8.26 [481
8.12 I471
15
6
7
4.35
2.48
2.89
2.70
3.31
1.96
2.61
2.49
1.635
2.43
2.35
1.3612.51
1 2.30
I8
1 9
-
I
1
2
4
-0.01
I
7.3 [471
3
0.71
0.5
4
1.18
1.0
5
1.52
1.2
7.43
7.00 1491
6.96 1501
6
0.05
0.1
8.06
8.03
8.10
8.03
7
0.61
0.56
7.6 I491
7.57 [47]
8
1.11
1.1
9
1.45
1.45
10
-0.02
I1
0.54
12
1.06
-0.05
0.5
1.41
1.4
14
0.31
0.3
16
0.55
0.56
-
17
1.04
0.9
18
1.36
1.36
19
0.27
0.15
-
20
0.55
0.6
2.63
6.62
2.15
7.6
4.24
7.53
3.45
I 7.03
I 6.69
8.0 [491
8.1 [SO]
I
7.55 [47]
I 7.6
2.74
-
2.03
1
1.69
2.17
4.36
2.47
3.55
-
2.04
I 6.72
2.48
2.28
2.25
8.17
I 7.08
1.1
-
13
I 2.65
I
2.53
2.81
1.705
2.30
2.30
1.43
2.23
3.88
2.47
7.95 [49]
15
21
0.38
0.4
22
0.61
0.6
23
1.03
0.85
24
0.78
0.6
25
[28]
1481
1491
[SO]
6.95
0.48
0.3
llrs 3.53
2.83
I 7.09 2.37
6.77
2.46
1.75
I 2.43
1.35
2.36
I
T 3.93 2?45/2.67
3.53
2.07
I 2.42
I 7.58 -
7.8 [49]
7.15 [49]
7.07 [501
( 7.6 I491
17.63 [50]
+I
1
7.11
3.77
2.29
3.45
2.025
2.85
1.75
I
11
1
1
II
2.47
2.32
2.23
2.64
2.49
2.34
2.24
2.55
2.41
2.38
2.26
2.58
I 2.47
2.37
2.51
2.42
2.46
2.47
I
2.39
7.35
3.20
1.985
2.47
2.46
7.65
3.63
2.14
2.59
2.53
I 2.45
I 2.52
I 2.49
1 2.47
-0.5
-
26
I
1
30
0.43
0.4
3.71
2.15
31
0.52
0.5
3.59
2.07
32
0.56
0.56
-
3.53
2.09
I 2.07
I 2.42
33
0.31
0.3
3.87
2.23
2.53
34
0.80
0.8
-
35
1.05
1.os
1.72
1 2.30
I 2.34
36
0.66
0.66
2.41
2.45
37
0.66
0.66
-
38
0.78
0.8
39
1.29
1.3
40
1.43
1.4
41
1.29
1.3
b) A benzene ring fused on at the 1,2-positions i.e. onto the
[ a ]side of a linear system, e.g. in Table 1, on passing, from No.
2 to No. 6 and from No. 3 to No. 7, from No. 4 to No. 8, and
from No. 5 to No. 9, affects the electron affinity and ionization energy only slightly. A benzene ring fused o n at the
3,4-positions, i.e. the [c] side, (e.g. on going from No. 7 to
No. 11, from 8 to 12, or from 9 to 13 in Table 1) [*I, at the
5,6-positions (from No. 6 to 14, from 7 to 16, from 8 to 17,
Angew. Chem. internat. Edii. 1 Yoi.3 (1964)
No. 9
1 7.82
-
7.33
-lI
I1 T
7.35
6.84
6.84
1.965
I 2.32
1
1I
1.50
I 2.28
1 2.27
3.17
1109 2.82
3.37
2.37
1.995
3.37
3.20
2.39
2.54
2.41
2.45
2.41
2.47
2.27
2.47
[*I With the exception of triphenylene (No. 10, Table l), which
has an extraordinarily small EA.
1471 H . Watanabe,J. chem. Physics 22, 1565 (1954); 26,542 (1957).
1481 M. E. Wacks and V . H . Dibeler, J. chem. Physics 31, 1557
(1959).
[49] According to Equation (7); see G . Briegleb and J. Czekalla,
Z . Elektrochem. Ber. Bunsenges. physik. Chem. 63, 6 (1959);
Angew. Chem. 72,401 (1960).
63 1
Table 8 (continued).
1 1=
4
9
Iexp
-Wolv
Eq. (441
I
Hydrocarbon no.
[*I
42
43
Et(36)
I
0.73
I
0.7
10.8710.9
7.27
___--
44
7.11
45
6.73
46
47
48
49
I-I0.5
0.42
1
0.40
-1-1
1
1 1 ::::1 1
7.72 [481
17.55 [49]
7.53 [SO]
3.08
2.86
2.31
1.415
2.39
2.27
2.35
-__
2.24
2.24
6.46
1.93
1.175
2.22
---
2.17
7.72
3.72
2.11
2.47
2.49
0.4
0.7
--
7.37
0.77
50
7.27
51
7.41
~
1.945
52
7.75
2.185
1I t: I I
2.42
2.52
7.04
2.75
7.30
3.12
1.815
1.65
22;
---
53
54
~
7.60
3.55
~
2.005
7.03
7.38
59
1 I I
I ::::I1 II
2.74
1.65
2.31
2.48
c
_
_
-
61
-1.08
1.1
1.03
1.02
1.0
~-
63
-1.15
65
1.15
1.55
64
I
7.10
1.0
62
1.55
I
2.33
2.40
1 2.32
58
~
7.06
7.11
11.78
2.86
::::5
2.36
1.65
2.34
1.69
2.36
r / 2 6 8 / F 2.36
l
6.58
I 2.10 I 1.30 I 2.22
-
2.52
I 2.70 I 1.65 1 2.29
2.15
1.28
2.27
-6.62
57
2.43
2.48
7.00
56
1
I
2.52
55
60
2.43
2.35
I 2.21
I 2.32
1 2.42
I 2.34
1 2.33
I1
2.35
2.31
2.20
0.5
['I For key t o hydrocarbon numbers, see Table 1.
and from 9 to 12), at the 7,s-positions (from No. 6 to 19,
from 7 to 20, and from 11 to 31) has essentially no effect on
the electron affinity and ionization energy. Accordingly, the
EA and I values for tribenz[a,c,g]anthracene (No. 31) and dibenzo[a,g]chrysene (No. 32) compared to dibenz[a,g]anthracene (No. 20) and benzo[a]chrysene (No. 30), respectively,
are practically identical.
c) Beginning with pyrene (Table 1, No. 47), a n additional
benzene or naphthalene ring fused o n at the 1,2- and/or 6,7positions (i.e. the [u] or [f]sides) has no particular effect on
the electron affinity and ionization energy. This is indicated
by comparison of No. 47 with Nos. 48, 52, and 55. The electron affinity of pyrene increases, and its ionization energy decreases, as a result of fusion of benzene or naphthalene rings
onto sides other than the [a],e.g. at the 3,4-, 8,9-, and 9,lOpositions([c], [h],and [i] sides). This can be seen by comparing
the electron affinity of pyrene with those of Nos. 49, 53, 54,
56, 57, and 58. Fusion onto the [c, h] sides has a greater effect
on the electron affinity than [c, i] fusion (Table 1 ; increase in
[50] From E,:
632
Equation (31), [36].
EA in going from Nos. 53 and 57 to Nos. 54 and 58, respectively).
d) Fusion of one or two benzene rings onto the [a] or [n, g]
sides (Table I , Nos. 39 and 40, respectively) of perylene
(Table 1 , No. 23) increases the electron affinity and decreases
the ionization energy to the same degree as does double fusion
onto the [b, h ] or [ u , j ] sides (No. 44 or No. 41). On passing
from perylene (No. 23) to benzo[l]perylene (No. 24) and to
coronene (No. 25), the electron energy decreases and the
ionization increases correspondingly.
e) The following compounds have the highest electron affinities, and, correspondingly, the lowest ionization energies:
benzo[g]terylene (No. 46): EA = 1.4 eV; I = 6.46 eV; dibenz[b, Alanthanthrene (No. 64): €A = 1.55; I = 6.58 eV;
dibenzo[a,g]perylene (No. 40): EA = 1.4; I = 6.7 eV;
benzo[a]pentacene (No. 9): EA = 1.45; I = 6.69 eV; dibenzo[u, clpentacene (No. 13): €A = 1.4, I = 6.72 eV, and dibenzo[a, dlpentacene (No. 18); €A = 1.4 eV, I = 6.77 eV.
Received, December 16tb, 1963 [A 365/17L IEI
German version: Angew. Chem. 76, 326 (1964)
Translated by scripta technica, New York
Angew. Chem. internat. Edit. I Vol. 3 (1964) 1 No. 9
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