Electrons and Chemical Bonding?

Bonding in Ovganic Molecules I75

The so-called experimental resonance energy of benzene is obtained by totaling the bond energies of the C-C, C=C, and C-H bonds present and comparing the total with the experimentally known value for the heat of formation of C6H6. The difference indicates that benzene is about 40 kcal/mole more stable than the sum of the bond energies for a system of six C-H, three C-C, and three isolated C==C units would suggest.

The value of p,, derived from the experimental resonance energy is therefore -20 kcal/mole. This value differs substantially from the value of

-

87 kcal/mole obtained from the absorption spectrum of C2H4. It is a general result that the resonance-energy p's are much smaller than the spectroscopic P's.

SUPPLEMENTARY PROBLEMS

1. Calculate the energies of the a molecular orbitals for C2H2.

2. Give the "bent-bond" descriptions of of HaCO; of HCN.

Bonds Involving d Valence Orbitals

T

here are many structures in which the central atom requires one or more d valence orbitals to complete a set of a bonding orbitals.

The most importa~lt of these structures are square planar, trigonal bipyramidal, square pyramidal, and octahedral; examples are shown in Fig. 9-1, Transition-metal ions have available a very stable set of d valence orbitals. The bonding in complexes formed between tran- sition-metal ions atld a large number of molecules and other ions un- doubtedly involves d orbitals. In this chapter we shall describe the bonding between metal ions and ligandsl in certain representative metal complexes.

9-2 THE OCTAHEDRAL COMPLEX Ti(HzO)G3+

The Ti3f ion forms a stable complex ion with six water molecules.

The structure around the Ti$+ ion is octahedral, as shown in Fig. 9-2.

1 Groups attached t o metal ions in complexes are called ligands.

176

Bonds Involving d Valence Orbitals

3 + F

F NH,

octahedral

C C1

square pyramidal 0

trigonal bipyramidal

square planar

Figure 9-1 Examples of structures in which d orbitals are used in bonding.

Electrons and Chemical Bonding

Z

41

I I I

OH, A

/

f OH*

I I I

.I

Figure 9-2 Coordinate system for Ti(Wz0)6".

The titanium has five 3d, one 4s, and three 4p valence orbitals that can be used in constructing molecular orbitals. Each water molecule must furnish one u valence orbital, which, in accord with the dis- cussion in Chapter VIII, is approximately an sp" hybrid orbital. We shall not specify the exact s and

p

character of the water ^u valence orbital, however, but simply refer t o i t as u.

The metal orbitals that can form a molecular orbitals are 3dz2-,2, 3dzz, 4s, 4p,, 4pY, and 4p.. Since the sign of the 4s orbital does not change over the boundary surface, the proper linear combination of ligand orbitals for 4s is

This is shown in Fig. 9-3. The wave function for the molecular orbital involving the metal 4s orbital is therefore

We find the other molecular orbitals by matching the metal-orbital

Bonds lnvolvtng d Valence O r b i ~ l s

0%

0%

(

4 s + o 1 + r a + ~ 1 + s r + o s + a 6

Figurc 9-3 Ovcrlap sf the titanium 45 orbital with the ^cr orbitals of the water molecules.

lobes with llgand a orbitals that have theproper sign and magnitude.

T h i s procedure is shown in Fig.

9-4. The

wave functions are:

$(vz) = c34pz

+

^{ca(a~- gas) (9-3)}

$ C F ~

= c34py

+

^c~(0-1

^-

^a41

^(9-4)

K c s ) = c34p8 $ C&(CS - ~ 6 ) cg-5)

$ ( ~ 2 4 ) " h 3 & 4

f

ca(al- ⁶²

+

^a8

^-

^{ad (9-6)}

t ( a 8 ) = 03d2

+

C ~ Z U S

+

^2as

-

^$1

^{- -}

^aa

^-

^a47

(9-7)

9-3 ENERGY LEVELS IN

Ti(Hq0)89+

Figure 9-4 shows 4p,, 4&, and 4p, to beequivalent 14 an octahedral complex; on this basis the c*, a=, and as molecular arbitals are de- generate in energy. Alrhough i t i~ not obv~ous from Fig.

9-4,

the

-

Electrons and Chemical Bonding

3d,-

+

^2s.

+

^{2 0 8}

^{- -}

^co

^-

^{o m}

( 3 d , = 2 9 - P - 4) Figure 9-4 Overlap of the tttanlum Ad and 41, o orbttals wath the s orbitals of the water molecules.

Bonds

Involving d Vale~ce Orbitals

3$2-,2 and 3d,2 orbitals are also equivalent in an octahedral complex, and a,z-,z and az2 are degenerate in energy. We shall solve a problem a t the end of this chapter to prove the equivalence of 3dX2-,2 and 3dZ2. Finally, we see that, including the a, orbital, there are three sets of u molecular orbitals in an octahedral complex: a,; a,, a,, a,;

and a,z-,z, a,2.

We have used all but three of the metal valence orbitals in the a molecular orbitals. We are left with 3d,,, 3d,,, and 3dx, These orbitals are situated properly for a bonding in an octahedral com- plex, as will be discussed later. However, since water is not a good

T bonding ligand, we shall consider that the 3dx,, 3d,,, and 3dx, orbit- als are essentially nonbonding in T i ( H ~ t 3 ) ~ ~ + . The three d, orbitals are clearly equivalent in an octahedral complex,. and we have the degenerate set: a,,, a,,, a,,.

In order t o construct an energy-level diagram for Ti(H20)63+, we must know something about the relative energies of the starting orbitals 3d, 4s, 4p, and aE20. In this case, aa20 is more stable than any of the metal valence orbitals. This is fairly general in metal complexes, and in energy-level diagrams the ligand a valence orbitals are shown to be more stable than the corresponding metal valence orbitals. I t is also generally true that the order of increasing energy for the metal valence orbitals in transition-metal complexes is nd

<

(n

+

^1)s

^<

^{( n}

+

^1)p.

The energy-level diagram for Ti(I-120)83+ is shown in Fig. 9-5 There are three sets of bonding orbitals and three sets of antibonding orbitals. The virtually nonbonding a ( d ) orbitals are less stable than the bonding a(d) set but more stable than the antibonding a(d) set.

The relative energies of the three bonding a sets are not known. The order given in Fig. 9-5 was obtained from a calculation that is beyond the level of our discussion.

We must count every electron in the valence orbitals used to con- struct the diagram in Fig. 9-5. The complex is considered to be com- posed of Ti3+ and six water molecules. Each of the six a valence orbitals of the water molecules furnishes two electrons, for a total of

182

Ti orbitals

Electrons and

Chem~cal

Bondtng

TI(H,O);+ orbitals H,O orbitals

twelve. S~nce the electronic structure of Tla+ is (3d)l, we have a total of thirteen electrons to place in the molecular orbitals shown in Fig. 9-5. The ground state of Ti(H20)ea+ is therefore

.

^{, ,..}

Bonds Inuolving d Valence Orbitals 1 ~ 3

There is one unpaired electron in the ~ ( d ) level. Consistent with this ground state, Ti(HzO)$f is paramagnetic, with S =

a.

The electrons in o bonding orbitals are mainly localized on the water molecules, since the u valence orbital of HzO is more stable than the metal orbitals. The nonbonding and antibonding orbitals, on the other hand, are mainly located on the metal. We shall focus our attention in the sections to follow on the molecular orbitals that are mainly based on the metal and derived from the 3d valence orbitals.

The difference in energy between 6 ( d ) and a(d) is called A or 10Dq. Excitation of the electron in =(dl to a*(& occurs with ab- sorption of light in the visible region of the spectrum, and Ti(HzO)i+

is therefore colored reddish-violet. The electronic spectrum of Ti(HzO)sa+ is shown in Fig. 9-6. The maximum absorption occurs at 4930 A, or 20,300 cm-l. The value of the splitting A is usually ex- pressed in cm-' units; thus we say that Ti(H@)s3+ has a A of 20,300 cm-'

.

The colors of many other transition-metal complexes are also due to such "d-d" transitions.

v, cm - 5 t

( ll Figure 9-6 The absorption spectrum of Ti(H2O)a3+ i n the

: i Q ' visible region.

184 Electvons and Chemical Bonding

9-6 VALENCE-BOND THEORY FOR Ti(K20)63+

The localized bonding scheme for Ti(IJ[20)63+ is obtained by first constructing six equivalent hybrid orbitals that are octahedrally directed. We use the six a valence orbitals of Ti for this purpose:

3d,2-y2, 3d,2, 4s, 4pZ, 4pU, and 4 p . Thus we want to construct six d2sp3 hybrid orbitals, each with one-third d character, one-sixth s char- acter, and one-half

p

character.

Referring back to Fig. 9-2, let us form linear combinations of the d, s, and

p

valence orbitals that direct large lobes at the six ligands. We first construct the orbitals that are directed toward ligands @ and @. We shall call these orbitals ^$5 and ^{$ 6 ,} respec- tively. The metal orbitals that can n bond with @ and @ are 3dz2, 4s, and 4p,. Choosing the coefficients of the 3dz2, 4s, and 4p, orbitals so that ^$6 and ^$6 have the desired d , s , and

p

character, we obtain the following hybrid-orbital wave functions:

The positive coefficient of 4pz in $5 directs a large lobe toward @, and the negative coegcient of 4p, in $6 directs a large lobe toward @.

The orbitals directed toward @ and @ are constructed from the 3d&,2,3d22, 4s and 4p2 metal orbitals. The orbitals directed toward @ and @ are constructed from the 3dz2-y2, 3dz2, 4s, and 4py orbitals.

The coefficients of 4s and 4p pose no problem, but we have to divide the one-third d character in each hybrid orbital between 3d,z and 342-4. We see from Eqs. (9-8) and (9-9) that we have "used up"

two-thirds of the 3d,Z orbital in $S and $6. Thus we must divide the remaining one-third equally among $1, $2,

&,

and rl.,. This means that each of $1, $2, $3, and t,b4 has one-twelfth 3d,2 character and one- fourth 3d,2,2 character. Choosing the signs of the coefficients so that a large lobe is directed toward each ligand in turn, we have:

@ndsInvplving d Valence Orbitals 185

1

$4 =

-%

3daxy2

- -

^{1 1}

.\/iz

^3d,3

+ ^-

^4s

^-

^-

^4pv

(9-13)

4 v5

These six localized

&p3

orbitals are used to form electron-pair bonds with the six water molecules. The valence-bond description of the ground state of Ti(HzO)t+ is shown in Fig. 9-7. The unpaired elec-

r'igure 9-7 Valence-bond representations of T ~ ( H S O ) ~ ~ + .

186

Electrons and

Chemtcal

Bonding

tron is placed in one of the d orbitals that has not been used to con- struct hybrid bond orbitals. This simple valence-bond orbital dia- gram is also shown in Fig. 9-7.

' ^I

9 7

CRYSTAL-FIELD THEORY FOR

Ti(H20)63+

In the crystal-field-theory formulation of a metal complex, we consider the ligands as point charges or point dipoles. The crystal- field model is shown in Fig. 9-8. The point charges or point dipoles constitute an electrostatic field, which has the symmetry of the com- plex. The effect of this electrostatic field on the ener ies of the metal d orbitals is the subject of our interest.

/"

Let us examine the energy changes in the 3d orbitals of Tia+ that result from placement in an octahedral field of point dipoles (the water molecules). First, all the d orbital energies are raised, owing to the proximity of the negative charges. More important, however, the two orbitals (3d,n, 3 d n 3 4 ) that point directly at the negative charges are raised higher in energy than the three orbitals (3d,,

'I

Figure 9-8 An octahedral field of point charges.

Bonds Involving

d

Valence

Orbitals

free ion mtahcdral crystal ficld

Figure 9-9 Splitting of the metal d orbitala in an octahedral crystal ficld.

3&, 3dm)

that are directed at points between the negative charges.

Thus we have a splitting of the five

d

orbitals in an octahedral crystal field as shown in Fig. 9-9. It is convenient to use the grouptheoret- ical symbols for the split

d

levels. The

3d2

and

3d.14

orbitals fonn the degenerate set called 8, and the

3d,, 3&,

and

3dw

orbitals form the degenerate set called h. The separation of ⁸ and h is again desig- nated A or 10Dg.

The one

d

electron in Tia+ is placed in the more stable h orbitals in the ground state. The excitation of this electron from tz to a is re- sponsible for the spectral band shown in Fig. 9-6.

9-8 RELATIONSHIP OF THE GENERAL MOLECULAR-ORBITAL TREATMENT TO THE VALENCE-BOND AND CRYSTAL-FIELD THEORIES

The valencebond and crystal-field theories describe different parts of the general molecular-orbital diagram shown in Fig. 9-5. The a bonding molecular orbitals are related to the six $spa bonding orbit- als of the valence-bond theory. The valence-bond theory does not include the antibonding orbitals, and therefore does not provide an explanation for the spectral bands of metal complexes. The A and

e levels of the crystal-field theory are related to the ~ ( d ) and a*(d) molecular orbitals. A diagram showing the relationship between the three theories is given in Fig. 9-10.

:.

^{, . .} . . I , . c ,