Gallup G.A. - Valence Bond Methods, Theory and applications (CUP, 2002)

24 2 H 2 and localized orbitals

where on the		φ the anterior superscript indicates the multiplicity and the posterior
subscript indicates the			m s value. The	η±1 /2	are the individual electron spin functions.
If we let		P i j	represent an operator that interchanges all of the coordinates of the
i th and	j th	particles in the function to which it is applied, we see that
			P 12	1 φ0	= −1 φ0 ,	(2.5)
			P 12	3φm s	= 3φm s ;	(2.6)

thus, the singlet spin function is antisymmetric and the triplet functions are symmetric with respect to interchange of the two sets of coordinates.

2.1.2 The spatial functions

The Pauli exclusion principle requires that the total wave function for electrons (fermions) have the property

	P	12	(1 , 2)		= (2	, 1) = − (1 , 2)	,	(2.7)
but absence of spin in the ESE requires
	(2 S +1) m	s (1 , 2) = (2 S +1) ψ (1 , 2) × (2 S +1) φm					s (1 , 2) ,	(2.8)
and it is not hard to see that the overall antisymmetry requires that the spatial
function	ψ have behavior opposite		to that	of		φ in all cases. We emphasize that		it
is not an oversight to attach no			m	s	label to	ψ in Eq. (2.8). An important principle

in quantum mechanics, known as the Wigner–Eckart theorem, requires the spatial
part of the wave function to be independent of		m s	for a given	S .
Thus the singlet spatial function is symmetric and the triplet one antisymmetric. If
we use the variation theorem to obtain an approximate solution to the ESE requiring
symmetry as a subsidiary condition, we are dealing with the singlet state for two
electrons. Alternatively, antisymmetry, as a subsidiary condition, yields the triplet
state.
We must now see how to obtain useful solutions to the ESE that satisfy these
conditions.
	2.2	The AO approximation
The only uncharged molecule with two electrons is H			2 , and	we will consider this
molecule for a while. The ESE allows us to do something that cannot be done in the
laboratory. It assumes the nuclei are stationary, so for the moment we consider a
very stretched out H	2 molecule. If the atoms are distant enough we expect each one

2.2

The AO approximation

25

to be a normal H atom, for which we know the exact ground state wave function.

1

The singlet wave function for this arrangement might be written

1 ψ (1, 2)

= N [1s a

(1)1s

b (2)

+ 1s b

(1)1s

a (2)],

(2.9)

where 1s

a

and 1s

b

are 1s

orbital functions centered at nuclei

a and

b , respectively,

and

N

is the normalization constant. This is just the spatial part of the wave function.

We may now work with it alone, the only influence left from the spin is the “+”

in Eq. (2.9) chosen because we are examining the singlet state. The

function

of

Eq. (2.9) is that given originally by Heitler and London[8].

Perhaps a small digression is in order on the use of the term “centered” in the last

paragraph. When we write the ESE and its solutions, we use a single coordinate

system, which, of course, has one origin. Then the position of each of the particles,

r i

for

electrons and

r α for

nuclei,

is given by a

vector

from

this

common

origin.

When determining the 1

s

state of H (with an infinitely massive proton), one obtains

the result (in au)

1s

(r )

1

exp(− r ),

(2.10)

=

√

π

where

r is

the radial

distance from

the origin

of

this H atom problem, which

is

where the proton is. If nucleus

α = a

is located at

r a

then 1s

a

(1) is a shorthand for

1s a

= 1s

(|r 1

− r a

|) =

1

exp(−| r 1

− r a

|),

(2.11)

√

π

and we say that 1s

a (1) is “centered at nucleus

a ”.

In actuality it will be useful later to generalize the function of Eq. (2.10) by

changing its size. We do this by introducing a scale factor in the exponent and write

1s (α, r ) =

π

exp(−α

r ).

(2.12)

α3

When we work out integrals for VB functions, we will normally do them in terms

of this version of the H-atom function. We may reclaim the real H-atom function

any time by setting

α = 1.

Let us now investigate the normalization constant in Eq. (2.9). Direct substitution

yields

1

= 1 ψ (1, 2)| 1 ψ (1, 2)

(2.13)

= |N

|2 ( 1s a

(1)|1s

a (1) 1s

b (2)|1s

b

(2)

+ 1s

b (1)|1s

b (1) 1s

a (2)|1s

a

(2)

+ 1s

a (1)|1s

b (1) 1s

a (2)|1s

b

(2)

+ 1s

b (1)|1s

a (1) 1s

b (2)|1s

a (2) ),

(2.14)

1The actual distance required here is quite large. Herring[26] has shown that there are subtle effects due to exchange that modify the wave functions at even quite large distances. In addition, we are ignoring dispersion forces.

26

2 H

2

and localized orbitals

where we

have

written out all of

the

terms.

The

1

s

function

in Eq.

(2.10) is

normalized, so

1 s a (1) |1 s a (1) = 1 s b

(2)

|1 s b

(2)

,

(2.15)

= 1 s b

(1)

|1 s b

(1)

,

(2.16)

= 1 s a

(2)

|1 s a

(2)

,

(2.17)

= 1 .

(2.18)

The other four integrals are also equal to one another, and this is a function of the

distance,

R , between the two atoms called the overlap integral,

S ( R

). The overlap

integral is an elementary integral in the appropriate coordinate system, confocal

ellipsoidal–hyperboloidal coordinates[27]. In terms of the function of Eq. (2.12) it

has the form

S (w) = (1

+ w + w2

/3) exp(

−w),

(2.19)

w = α R ,

1 ψ (1 , 2) is

(2.20)

and we see that the normalization constant for

=

+

N

[2(1

S

2 )]−1 /2 .

(2.21)

We may now substitute

1 ψ (1 , 2) into the Rayleigh quotient and obtain an estimate

of the total energy,

E ( R

) = 1 ψ (1 , 2) |H

|1 ψ (1 , 2) ≥ E

0 ( R ),

(2.22)

where

E 0

is the true ground state electronic energy for H

2 . This expression involves

four new integrals that also can be evaluated in confocal ellipsoidal–hyperboloidal

coordinates. In this case all are not so elementary, involving, as they do, expan-

α = 1 in all of the

sions in Legendre functions. The final energy expression is (

integrals)

E

2 h

2

j1 ( R

)

+ S ( R

)k 1 ( R )

j2 + k 2

1

,

(2.23)

+ 1

= −

1

+

S ( R )2

+

S ( R )2 + R

where

h

= α2 /2

− α,

(2.24)

j1

= −[1 − (1 + w)e −2 w ]/ R

,

(2.25)

k 1

= −α(1

+ w)e −w ,

(2.26)

j2

= [1 − (1 + 11 w/8 + 3w2 /4 + w3/6)e −2 w ]/ R

,

(2.27)

k 2 = {6[S (w)2 (C + ln w) − S (−w)2 E 1 (4 w) + 2 S (w) S (−w)E 1 (2 w)]

+ (25 w/8 − 23 w2/4

− 3w3

− w4 /3)e −2 w }/(5 R ),

(2.28)

l2

= [(5/16 + w/8 + w2 )e −w − (5/16 + w/8)e −3w ]/ R

.

(2.29)

	2.4	Extensions to the simple Heitler–London				treatment	27
Of these only the exchange integral of Eq. (2.28) is really troublesome to evaluate.							−w,
It is written in	terms of	the overlap integral		S (w ),		the same function of
S (−w ) = (1	− w + w2 /3)e w , the Euler constant,				C	= 0.577 215 664 901 532 86,
and the exponential integral		E 1 ( x ) = x	∞ e −y	dy
					,		(2.30)
				y

which is discussed by Abramowitz and Stegun[28].

In our discussion we have merely given the expressions for the five integrals that appear in the energy. Those interested in the problem of evaluation are referred to Slater[27]. In practice, these expressions are neither very important nor useful. They

are essentially restricted to the discussion of this simplest case of the H 2 molecule and a few other diatomic systems. The use of AOs written as sums of Gaussian

functions has become universal except for single-atom calculations. We, too, will use the Gaussian scheme for most of this book. The present discussion, included for historical reasons, is an exception.

	2.3 Accuracy of the Heitler–London function
We are now in a position to compare our results with experiment. A graph of		E ( R )
given by Eq. (2.23) is shown as curve (e) in Fig. 2.1. As we see, it is qualitatively
correct, showing the expected behavior of having a minimum with the energy rising
to infinity at shorter distances and reaching a finite asymptote for large		R values.
Nevertheless, it misses 34% of the binding energy (comparing with curve (a) of
Fig. 2.1), a significant fraction, and its minimum is clearly at too large		a bond
distance.	2

	2.4 Extensions to the simple Heitler–London treatment
In	the last section, our calculation used	only	the function of	Eq. (2.9),		what	is
now called the “covalent” bonding function. According to our discussion of linear
variation functions, we should see an improvement in the energy if we perform a
two-state calculation that also includes the				ionic	function,
	1 ψI(1, 2)	= N	[1s a (1)1s	a (2)	+ 1s	b (1)1s b (2)].		(2.31)
When this is done we obtain the curve labeled (d) in Fig. 2.1, which, we see,
represents a small improvement in the energy.
2	The quantity we have calculated here is appropriately	compared	to			D e ,	the bond energy from the	bottom of
	the curve. This differs from the experimental bond energy,			D	0 , by the	amount	of energy due to the zero	point
	motion of the vibration. There is no vibration in our system, since the nuclei are infinitely massive. We use the
	theoretical result for comparison, since it is today considered more accurate than experimental numbers.

28	2 H 2 and localized orbitals

	−	1.00
	−	1.05
(hartree)
Energy	−	1.10	(e)
			(d)
			(c)
	−	1.15	(b)
			(a)
		0	2	4			6	8	10
			H	---	H	distance	(bohr)

Figure 2.1.

Energies of H

2

for various calculations using the H-atom

1 s orbital functions;

(b) covalent

+ ionic,

scaled;

(c) covalent only,

scaled; (d) covalent

+ ionic, unscaled;

(e) covalent only, unscaled. (a) labels the curve for the accurate function due to Kolos and

Wolniewicz[31]. This is included for comparison purposes.

A much larger improvement is obtained if we follow the suggestions of Wang[29]

and Weinbaum[30] and use

the 1

s (1) of Eq. (2.12), scaling

the

1

s

H function

at

each internuclear distance to give the minimum energy according to the variation

theorem. In Fig. 2.1 the covalent-only energy is labeled (c) and the two-state energy

is labeled (b). There is now more difference between the oneand two-state energies

and the better binding energy is all but 15% of the total.

3 The scaling

factor,

α,

shows a smooth rise from

≈1

at large

distances to a value

near

1.2

at

the energy

minimum.

Rhetorically, we might

ask, what

is it

about

the

ionic function that

produces

the energy lowering, and just how does it differ from the covalent function? First,

we note that the normalization constants for the two functions are the same, and,

indeed, they represent exactly the same charge density. Nevertheless, they differ in

their two-electron properties.

1 ψC

(adding a subscript

C

to indicate covalent) gives

a higher probability that the electrons

be far from

one another, while

1 ψI gives

3Perhaps we should note that in this relatively simple case, we will approach the binding energy from below as we improve the wave function. This is by no means guaranteed in more complicated systems where both the equilibrium and asymptotic wave functions will be approximate. One of the most important problems of bond

energy calculations is to have a “balanced” treatment. This happens if the accuracies at the equilibrium and asymptotic geometries are sufficiently high to give an accurate difference. More realistically in larger systems, one must have any errors at the two geometries sufficiently close to one another to obtain a useful value of the difference. We will take up this question again later.

2.4 Extensions to the simple Heitler–London treatment	29
just the opposite. This is only true, however, if the electrons are close to one or the
other of the nuclei. If both of the electrons are near the mid-point of the bond, the
two functions have nearly the same value. In fact, the overlap between these two
functions is quite close to 1, indicating they are rather similar. At the equilibrium
distance the basic orbital overlap from Eq. (2.19) is
S = 1s s |1s b = 0.658 88.	(2.32)
A simple calculation leads to

	=		C			2 S	=
					= 1 + S 2
		1	ψ	1 ψI				0.918 86 .
(We consider these relations further below.) The covalent function has been char-
acterized by many workers as “overcorrelating” the two electrons in a bond.
Presumably, mixing in a bit of the	ionic		function ameliorates the overage, but
this does not really answer the questions at the beginning of this paragraph. We
take up these questions more fully in the next section, where we discuss physical
reasons for the stability of H			2 .
At the calculated energy minimum (optimum							α) the total wave function is found
to be
= 0.801 981					1 ψC + 0.211 702				1 ψI.

(2.33)

(2.34)

The relative values of the coefficients indicate that the variation theorem thinks

better of the covalent				function, but	the other appears fairly	high	at	first glance.
If,	however,	we apply the EGSO process described in Section				1.4.2,		we obtain
0.996 50		1 ψC	+	0.083 54	I
					1 ψ , where, of course, the covalent function is unchanged,
but	1 ψI	is the new ionic function orthogonal to					1 ψC . On the basis of this calculation
we conclude that the the covalent character in the wave function is (0								.996 50) 2 =

0.993 (99.3%) of the total wave function, and the ionic character is only 0.7%. Further insight into this situation can be gained by examining Fig. 2.2, where a

geometric representation of the basis vectors and the eigenfunction is given. The overlap integral is the inner product of the two vectors (basis functions) and is the

cosine of the angle between them. Since arccos(	) = 23.24 ◦, some care was taken
with Fig. 2.2 so that the angle between the vectors	representing the covalent and

ionic basis functions is close to this value. One conclusion to be drawn is that these two vectors point, to a considerable extent, in the same direction. The two smaller segments labeled (a) and (b) show how the eigenfunction Eq. (2.34) is actually put together from its two components. Now it is seen that the relatively large coefficient

of 1 ψIis required because it is poor in “purely ionic” character, rather than because the eigenvector is in a considerably different direction from that of the covalent basis function.

30 2 H 2 and localized orbitals

"Ionic" orthonormal vector

Ionic

Eigenvector

(b)

Covalent

(a)

"Covalent" orthonormal vector

Figure 2.2. A geometric representation of functions for H	2 in terms of vectors for
The small vectors labeled (a) and (b) are, respectively, the covalent and ionic components of
the eigenvector. The vectors with dashed lines are the symmetrically orthogonalized basis
functions for this case.
It is important to realize that the above geometric representation of	the H
Hilbert space functions is more than formal. The overlap integral of two normalized
functions is a real measure of their closeness, as may be seen from

R = R eq .

2

1 ψC

−

1 ψC

−

=

−

1 ψI

1 ψI

2(1

),

(2.35)

and, if the two functions were exactly

the

same,

would be 1. As pointed out

above,

is a dependent upon

S , the orbital overlap. Figure 2.3 shows the relation

between these two quantities for the possible values of

S .

In addition, in Fig. 2.2 we have plotted with dashed lines the symmetrically

orthogonalized basis functions in this treatment. It is simple to verify that

C

−

−

=

1

ψ

S

1 ψI 1 ψI

S 1 ψC

0 ,

(2.36)

where

S is the orbital overlap. Therefore, the

vectors given in Fig. 2.2 are just the

normalized versions of those in Eq. (2.36). Since they must be at right angles, they

must move out 33.38

◦ from the vector they are supposed to approximate. Thus, the

real basis functions are closer together than their orthogonalized approximations

are to the functions they are to represent. Clearly,

writing the eigenfunction

in

terms of the two symmetrically orthogonalized basis functions will require nearly

equal coefficients, a

situation giving

a very overblown view of

the importance

of

2.5 Why is the H

2 molecule stable?

31

	1.0
	0.8
	0.6	ψ	C	ψ
				I
overlap					1s	a 1s	b
−ionic	0.4
Covalent
	0.2
	0.0
	0.0	0.2		0.4	0.6		0.8	1.0
				Orbital	overlap

Figure 2.3. The relation between the orbital and covalent–ionic function overlaps.

the “ionic” component of the eigenfunction.

√

◦

=

√

4

Since cos 30

√3/2, this situation

will be the same for any pair of vectors with

S

>

3

/2. Only when

S

3

/2, is it

likely that the symmetrical orthogonalization will, in the two-vector case, produce

coefficients that are qualitatively useful.

Another interesting conclusion can be drawn from Fig. 2.2. The linear combina-

tion of AOs-molecular orbital (LCAO-MO) function for H

2 is

M O

= 0.510 46

1 ψC

+ 1 ψI ,

(2.37)

and, although it has not been shown in the figure to reduce

crowding, the

vector

corresponding to

M O

would appear halfway between the two for

1 ψC and

1 ψI. It

is thus only a little farther from the optimum eigenfunction than

1 ψC

alone.

2.5 Why is the H

2

molecule stable?

Our discussion so far has not touched upon the origin of the stability of the H

2

molecule. Reading from the articles of the early workers, one obtains the impression

that most of them attributed the stability to “resonance” between the 1

s a (1)1s b (2)

form of the wave function and the 1

s a

(2)1s

b

(1) form. This phenomenon was new

to physicists and chemists at the time and was frequently invoked in “explaining” quantum effects. Today’s workers prefer explanations that use classical language,

4Although the function is more complicated, this is actually what happened in the third column of Table 1.1, where the first two entries are nearly equal compared to those in other columns.

32	2 H 2 and localized orbitals

insofar as possible, and attempt to separate the two styles of description to see as clearly as possible just where the quantum effects are.

2.5.1 Electrostatic interactions

Equation (2.23) is not well adapted to looking at the nature of the bonding. We rearrange it so that we see the terms that decrease the energy beyond that of two H atoms. This gives

E

( R )

=

2

E

H

+

J ( R ) + K

( R

)

.

(2.38)

)

1

+

T

( R

Here

E H is the energy of

a

normal

hydrogen atom,

J

( R

),

K

( R ), and

T ( R ) were

called by Heitler and London the

Coulomb integral,

the

exchange

integral,

and

the overlap integral, respectively. The reader should perhaps be cautioned that the

terms “Coulomb”, “exchange”, and “overlap” integrals have been used by many

other workers in ways that differ from that initiated by Heitler and London. In the

present section we adhere to their original definitions, as follows:

J

( R

)

= 1 s a

(1)1

s b

(2)

|V (1 , 2)

|1 s a

(1)1

s b

(2)

,

(2.39)

K

( R

)

= 1 s a

(1)1

s b

(2)

|V (1 , 2)

|1 s b

(1)1

s a

(2)

,

(2.40)

T

( R

)

= 1 s a

(1)1

s b

(2)

|1 s b

(1)1

s a

(2) ,

(2.41)

= 1 s a

(1) |1 s b (1)

2 ,

and

V (1 , 2) = −1/r 2 a

− 1/r 1 b

+ 1/r 12 + 1/R ab

.

(2.42)

These equations are obtained by assigning electron 1 to proton

a

and

2

to

b , so

that the kinetic energy terms and the Coulomb attraction terms

−1/r 1 a

− 1/r 2 b

give

rise

to the 2

E H

term in Eq. (2.38).

V (1 , 2)

in

Eq.

(2.42) is

then

that part of

the

Hamiltonian that goes to zero for the atoms at long distances. It is seen to consist of

two attraction terms and two repulsion terms. As observed by Heitler and London,

the

bonding in the

H

2

molecule

arises

from the way these terms

balance

in

the

J

and

K integrals. We show a graph of these integrals in Fig. 2.4. Before continuing,

we discuss modifications of Eq. (2.38) when scaled 1

s

orbitals are used.

With the 1

s

function of Eq. (2.12), we obtain

E (α, R

)

=

2

E

H

+

(α

−

1) 2

+

α

J (α R

)

+ K

(α R )

,

(2.43)

1

+

T (α R )

which reduces to the energy expression of Eq. (2.38) when

α = 1. The

changes

brought by including the scale factor are only quantitative in nature and leave the

qualitative conclusions unmodified.

2.5 Why is the H

2 molecule stable?

33

	3
	2
	1
	0
(eV)	−1	J(R)
Energy	−2
	−3
	−4	K(R)
	−5
	1	2		3					4	5
		Internuclear	distance	(bohr)
	Figure 2.4.	The relative sizes of the		J	( R	) and	K	( R	) integrals.

It is important to understand why the			J ( R	) and	K ( R	) integrals have the sizes they
do. We consider		J ( R ) first. As we have seen from Eq. (2.42),				V (1, 2) is the sum of
four different Coulombic terms from the Hamiltonian. If these are substituted into
Eq. (2.39), we obtain
		J ( R ) = 2 j1 ( R ) + j2 ( R ) + 1/ R ,
		j1 ( R ) = 1s a | − 1/ r b |1s a = 1s b | − 1/ r a |1s b ,
		j2 ( R ) = 1s a (1)1s	b (2)|1/	r 12 |1s a	(1)1s	b (2) .
The quantity	j1 ( R	) is seen to be the energy of Coulombic attraction between a point
charge and a spherical charge distribution,				j2 ( R ) is the energy of Coulombic repul-
sion between two spherical charge distributions, and 1/						R is the energy of repulsion
between two point charges.		J ( R ) is thus the difference between two attractive and
two repulsive terms	that	cancel to a considerable	extent. The	magnitude		of the

charges is 1 in every case. This is shown in Fig. 2.5, where we see that the resulting difference is only a few percent of the magnitudes of the individual terms.

This is to be contrasted with the situation for the exchange integral. In this case we have

K ( R ) = 2k 1 ( R ) S ( R ) + k 2 ( R ) + S ( R )2 /R ,

k 1 ( R ) = 1s a | − 1/ r b |1s b = 1s a | − 1/ r a |1s b , k 2 ( R ) = 1s a (1)1s b (2)|1/ r 12 |1s a (2)1s b (1) .