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Файл:Ординатура / Хирургия / Библиотека им академика М.И. Перельмана / Книга_5852_Библиотеки_им_академика_М_И_Перельмана.pdf
C H A P T E R 1 Structure and Properties
1s 1s
1s 1s
σ
1s
+
+
+
−
+ +
σ*
1s
+
+
−
−
+ -+
2p
z
2p
z
2p
z
−
−
+
−
+
+
2p
z
−
+
−
+
−
-
++
σ
2p
z
σ∗
2p
z
+
+−
+
−
+
π
2p
x
π
2p
x
Node
It is clear from the above discussion that the main difference between the two types of orbitals is in the
number of nodal planes. The σ orbital has no nodal plane while the π-orbital has one. Since the energy of a
molecular orbital is directly related to the number of nodal planes, the π-orbital is more energetic than the
σ-orbital. This explains why a π-bond is a weaker bond than σ-bond.
Figure 1.18 Formation of σ1s and σ*1s bonding and antibonding molecular orbitals.
41
Figure 1.19 Formation of σ
Figure 1.20 Formation of π
2pz
and σ*
and π*
2px
bonding and antibonding molecular orbitals.
2pz
bonding and antibonding molecular orbitals.
2pz
42
PHARMACEuTICAL ORgAnIC CHEMISTRy
1.11.1 Difference between s and p Molecular Orbitals
S. no.
1. It is formed from σ bonds by the end-on-end
overlapping of s–s, s–p, p–p orbitals.
2. Since the overlap is large, the bonds are quite strong. Since the overlap is not very large, the bonds formed are
3. The electron density is maximum along the
internuclear axis.
4. σ MO has no nodal plane while σ* has only one. π MO has one nodal plane while π* has two.
5. Bond σ and σ* are symmetric with respect to
rotation about the internuclear axis.
Sigma (s) MO Pi (p) MO
It is formed from bonds by the sidewise overlap of p–p
or p–d orbitals.
weak.
The electron density is maximum both above and below
the plane, which contains the internuclear axis.
Both π and π* are antisymmetric w.r.t. rotation about the
internuclear axis.
NOTEWORTHY POINTS
Using appropriate signs for the wave functions of the atomic orbitals, the overlap can be positive or
negative or zero.
σ1. bonds are always stronger than π bonds.
π2. bonds are formed only when σ bond has already formed.
So, single bond has 1σ bond.
Double bond has 1σ and 1π bond.
And triple bond has 1σ and 2π bonds.
H atom always form 3. σ bond.
π4. bonds are more reactive than σ bonds.
Free rotation around sigma bond is possible whereas rotation is restricted around 5. π bond.
1.12 COMPARISON OF VALENCE BOND THEORY AND MOLECULAR ORBITAL THEORY
1. Points of similarity: These are as given below.
a. In both the theories, the resultant electron distributions for given molecules are similar (rules for
llingof orbitals are similarand followAufbau principle,Pauli’sexclusion principle and Hund’s
rules).
b. Both require that in the description of normal covalent bond, the essential interaction is the sharing
of the bounding electron density by both nuclei and concentration of the electron density between the
nuclei.
c. Both require that in bond formation, the combining orbitals must be of nearly the same energy. Orbitals
of the bonded atoms overlap and have appropriate symmetry about the molecular axis in order to form
a bond.
C H A P T E R 1 Structure and Properties
d. Both theories are approximations because of impossibility of describing molecules with an exact
Schrodinger equation.
e. Both theories account for the directional nature of bond.
f. Both predict the nonexistence of helium molecule.
2. Points of difference: These are given in the table below.
S. no. Valence bond theory Molecular orbital theory
43
1. Only the electron present in the outermost (valence)
shell take part in the formation of covalent bond.
2. Atoms retain their individual identity.
All the electrons of all the atoms of the molecules are
affected and take part in the bond formation.
Atoms lose their identity.
3. It considers the interaction between atoms. It starts with the nuclei of the constituent atoms.
4. Resonance phenomenon is a part and parcel of V.B.T. It does not consider the resonance phenomenon at all.
5. It does not predict the presence of bonding and
antibonding molecular orbitals.
It does consider bonding and antibonding molecular
orbitals; the bonding molecular orbital causes
attraction while the antibonding molecular orbital
causes repulsion.
6. It cannot explain the paramagnetic nature of O
7.
Itdoesnotexplainthebondinginelectron-decient
compounds.
8. The properties like bond order and electronic
congurationintheexcitedstatecannotbepredicted.
9. It cannot explain the existence of species like H
+
He
.
2
. It explains the paramagnetic nature of O2.
2
Itdoesexplainthebondinginelectron-decient
compounds.
It does predict the bond order and the electronic
congurationintheexcitedstate.
+
and
2
It can explain their existence.
10. The wave functions include the ionic contribution. The wave functions do not include the ionic
contribution.
11. It is simple to apply.
Itisdifculttoapply.
1.13 HYBRIDIZATION
According to orbital concept of overlapping, a covalent bond is formed by the overlap of two atomic
orbitals, each containing a single electron with opposite spin. Thus the number of covalent bonds which
an atom can form is equal to the number of unpaired electrons in it or in other words, the number of
half-lled orbitals present in its valence shell. However, the atom of carbon is an exception to this
generalization.Theelectroniccongurationofcarbonatom(atomicno.=6)intheground state is given
below:
1
C (ground state) : 1s22s22p
1
2p
x
y
Aglanceattheelectroniccongurationrevealsthatthecovalencyofcarbonatomshouldbetwoasit
hashalf-lledatomicorbitals.But,inactualpractice,carbonshowsacovalencyoffoursuchasinCH4,
CCl
, etc.
4
44
PHARMACEuTICAL ORgAnIC CHEMISTRy
To account for the anomalous behaviour of carbon atom, it is assumed that prior to the bond formation,
one of the 2s electrons in this atom gets promoted to the vacant higher energy 2p orbital. Consequently, the
excited stateelectroniccongurationofcarbonatomisasfollows:
1
1
1s22s12p
1
2p
2p
z
x
y
Thus, carbon has four unpaired electrons in its excited state, thereby explaining the tetravalent nature of
carbon.
no doubt, such a process of excitation of electrons requires energy. However, this is more than compensated
by the energy released in the formation of two additional covalent bonds.
It is evident that one 2s and three 2porbitalscanoverlapwiththehalf-lledorbitalsoftheotheratoms
involved and thereby form four covalent bonds. But these four bonds should not be equivalent since
the 2s orbital is spherically symmetrical while the 2p orbitals are dumbbell shaped. As such, one bond
(derived from 2s orbital) should be of one type and three bonds (derived from three 2p orbitals) should
be of other type. But actually all the four bonds of carbon are equivalent in all respects. For example,
all the four C–H bonds in methane (CH4 ) are equivalent and are directed towards the four corners of a
regular tetrahedron.
To explain this abnormal behaviour, the concept of hybridization was introduced. According to this
concept, it is postulated that one s and three p orbitals of carbon, having a slight difference in energies mix
up to give four new identical orbitals, i.e. orbitals which have the same energy. This concept of intermixing
of two or more atomic orbitals of slightly different energies to form an equal number of new orbitals which
have same energy, identical shapes and are symmetrically oriented in space is called hybridization. The new
orbitals thus obtained are known as hybridized or hybrid orbitals.
1.13.1 Conditions for Hybridization
necessary conditions for hybridization are as follows:
1. The orbitals of the valence shell of an atom only undergo hybridization.
2. The orbitals undergoing hybridization must have a small energy difference.
3. Bothhalf-lledandcompletelylledorbitalscanundergohybridization.
1.13.2 Characteristics of Hybridization
1. The number of hybrid orbitals formed is equal to the number of orbitals involved in hybridization.
2. The hybrid orbitals have equivalent energy and identical shape.
3. Hybrid orbitals are more effective in forming strong directional bonds than the pure atomic
orbitals.
4. The type of hybridization represents geometry of molecules.
5. Like atomic orbitals, a hybrid orbital can accommodate at the most two electrons with opposite
spins.
C H A P T E R 1 Structure and Properties
1.13.3 Types of Hybridization
Depending upon the number and types of the atomic orbitals involved, many types of hybridization are
possible. However, few main types of hybridization are as follows:
1. Diagonal or sp hybridization
2. Trigonal or sp2 hybridization
3. Tetrahedral or sp3 hybridization
3
4. sp
d hybridization
5. sp3d2 hybridization
6. sp3d3 hybridization
7. dsp2 hybridization
1. Diagonal or sp hybridization: If one s and one p orbital of the same energy level of an atom intermix to
form two new equivalent orbitals, the phenomenon is said to be sp hybridization or diagonal hybridization.
The orbitals thus formed are called sp orbitals.
Each sp hybrid orbital contains one electron and has 50% s character and 50% p character. Because of
interelectronic repulsions, the two hybrid orbitals tend to remain as far apart as possible from each other.
This is feasible only if they are inclined to each other at an angle of 180º, i.e. the two sp hybrid orbitals
are collinear (Fig. 1.21). It is due to this reason that sp hybridization is also called diagonal or linear
hybridization.
45
180º
sp
Hybridization
2s 2p
Figure 1.21 Diagonal (sp hybridization).
This type of hybridization occurs in the following:
a. All compounds of beryllium such BeF2, BeCl2, etc.
b. All compounds of carbon containing a triple bond between carbon atoms (C C), e.g.
acetylene.
2. Trigonal or sp
2
hybridization: This type of hybridization occurs when one s and two p orbitals of the
same energy level of an atom intermix to form three new equivalent orbitals. The new orbitals thus formed
are called sp2 hybrid orbitals. Because of interelectronic repulsions, the three sp2 hybrid orbitals tend to
remain farthest apart from one another. This is possible only if they are inclined to each other at an angle
of 120º, i.e. if all the three hybrid orbitals remain in the same plane and are directed along the corners of an
equilateral triangle (Fig. 1.22). It is due to this reason that sp2 hybridization is also referred to as trigonal
hybridization.
46
sp2 Hybridization
2s
2
p
x
Three sp2 hybrid orbitals
sp3 Hybridization
2p
y
2s
x
2p
z
Four sp3 hybrid orbitals
109°28¢
PHARMACEuTICAL ORgAnIC CHEMISTRy
Each of the sp2 hybrid orbitals has 33–34% s character and 66–67% p character. This type of hybridization
is found in the following compounds:
a. All compounds of boron such as BF3, BCl3, BH3, etc.
b. All compounds of carbon having a double bond between carbon atoms (C C), e.g. ethylene.
Figure 1.22 Trigonal (sp2 hybridization).
3. Tetrahedral or sp
3
hybridization: If one s and three p orbitals of the same energy of an atom intermix
to form four new equivalent orbitals, the hybridization is called sp3 hybridization and the new orbitals
formed are called sp3 hybrid orbitals. Due to interelectronic repulsions, these orbitals tend to remain
farthest apart and are, therefore, directed towards the four corners of a regular tetrahedron forming an
angle of 109º28′ (Fig. 1.23).
Figure 1.23 Tetragonal (sp3 hybridization).
It is on this account that sp3 hybridization is also referred to as tetrahedral hybridization.
3
Each sp
hybrid orbital has 25% s character and 75% p character. This type of hybridization occurs in all
compounds of carbon that contain carbon–carbon single bond (C–C) only, e.g. methane, ethane, etc. The
characteristics of sp, sp2 and sp3 hybridization are given in Table 1.3.
C H A P T E R 1 Structure and Properties
3s 3p
(Ground state)
3d
3s 3p
sp
3
d Hybridization
(Excited state)
3d
Table 1.3 Characteristics of sp, sp2 and sp3 hybrid orbitals
47
Type of
hybridization
sp s + p
2
sp
3
sp
3
4.
sp
d Hybridization: Let us explain the formation of sp3d hybrid orbitals, taking the case of
Atomic
orbitals
s + 2 (p)
s + 3 (p)
Shape
Bond
angle
Linear 180º
Trigonal planar 120º
Tetrahedral
109º 28'
Example
BeH
, HgCl2, BeCl2, ZnCl2, C2H
2
, BCl3, AlCl3, C2H4, CO
BF
3
CH4, SnCl4, SiCl4, SiF4, nH
23-
+
, ClO
4
, nO
2
-
3
-
4
phosphorus.
Thegroundstateelectroniccongurationofphosphorusatommayberepresentedas:
1
1
1
3s23p
3p
x
0
3p
3d
y
z
or
In order to explain the pentavalency of phosphorus, one of the 3s electrons is promoted to the vacant 3d
orbital, which is comparable in size and energy to 3s and 3porbitals.Theexcitedstatecongurationof
phosphorus may be represented as follows:
Nowthereareveorbitals:onces, three p (px, py and pz) and one d (d
Theseorbitalshybridizetoyieldasetofvesp3d hybrid orbitals, giving a trigonal bipyramidal structure.
This is made up of two processes: an initial hybridization of s, px and py orbitals to yield three sp2 orbitals
at an angle of 120º in xy plane and then hybridization of pz and d
angle (90º) to this plane, one above and one below (Fig. 1.24). It is quite clear from Fig. 1.25 that the entire
vehybridorbitalsarenotequivalent.Therearetwosetsofequivalenthybridorbitals:onesetconsistsof
three coplanar (120º) equivalent equatorial (e) sp2 orbitals and the other consisting of two equivalent axial
(a) dp orbitals.
Shape of PCl
chlorine atoms (3s2, 3p
2
)orbital,whicharesinglylled.
z
2
orbitals to yield two d–p orbitals at right
z
: Thevesp3dhybridorbitalsofphosphorusoverlapwithvehalf-lledporbitalsfromve
5
2
2
3p
3p
x
y
1
) and a trigonal bipyramidal PCl5 results.
z
48
z
z
z
z
z
p
y
p
x
p
z
Five sp3d
hybrid orbitals
d
z
2
y
y
y
y
y
x
x
x
x
s
x
z
y
x
CI
CI
CI
CI
CI
P
120°
90°
a
e
e
a
e
PHARMACEuTICAL ORgAnIC CHEMISTRy
Figure 1.24 sp3d hybridization.
5.
sp 3d
In order to explain the hexavalency of sulphur two electrons (one 3s and one 3p) are promoted to empty
and d
2
Hybridization: Let us explain the formation of sp3d2 hybrid orbitals, taking the case of sulphur.
2
2
orbitals.
–d
x
y
Figure 1.25 Shape of PCl5 molecule (trigonal bipyramidal).
2
d
z
49
C H A P T E R 1 Structure and Properties
Theexcitedstatecongurationofsulphurmayberepresentedasfollows:
now there are six orbitals: one s, three p (p
x
, py, pz) and two d (d
z
2
and d
x
2
– d
y
2
) orbitals, which are singly
lled.Theseorbitalshybridizetoyieldasetofsixsp
3d2
hybridized orbitals directed towards the corners of
a regular octahedron as shown in Fig. 1.26.
3s 3p
(a) P (Ground state)
3d
sp3d2 hybridization
FFFF FF
(c) SF
6
(b) Excited state
Figure 1.26 Formation of SF6 molecule involving sp3d2 hybridization.
Shape of SF6: Six sp3d2 hybridorbitalsofsulphuroverlapwithsixhalf-lledporbitalsofsixuorineatoms
(2s2 2p
x
2
2p
y
2
2p
z
1
) and an octahedral SF6 molecule results.
F
F
F
F
F
F
S
6. sp 3d 3 Hybridization: It involves the intermixing and
redistribution of seven atomic orbitals, namely one s, three
p and three d orbitals forming seven sp3d3 hybrid orbitals having pentagonal bipyramidal geometry. The
geometry of IF7 molecule can be explained on the basis of this type of hybridization.
Geometry of IF
7
molecule:Theouterelectroniccongurationofiodineatomis5s25p5. There is only one
half-lledatomicorbitalavailableintheoutermostshellinthegroundstateoftheatom.SotheIatomisin
the excited state in its compounds. To make seven bonds, one electron from 3s and two electrons from 3p
are promoted to the higher vacant 5d orbitals as shown in Fig. 1.26 (a). These seven orbitals are hybridized
to give seven sp3d3 hybrid orbitals. Each of these hybrid orbitals overlaps with 2porbitalsofuorinetoform
IF7 molecule having pentagonal pyramidal geometry.
Learning Plus
PCl3 dissociates to give PCl3 and Cl, and the
equatorial and axial bond length in PCl5 are not
equal.
The P–Cl equatorial bond length of the three
bonds is larger than that of the two axial P–Cl
bonds in PCl
5
molecules. The reason for the larger
axial bond length is that the equatorial bonds are
repelled by one or two electron pairs, but the axial
P–Cl bonds are repelled by the three electron
pairs. Thus the repulsion exerted on the axial
bond pairs is more and hence their bond length is
larger. These two axial bonds are easily broken in
comparison to the other equatorial bonds.
50
F
F
F
F
F
F
F
sp
3d3
hybridization
5p 5d5s
F FFFFFF
I
(a) I (Ground state)
(c) IF
7
(b) Excited state
dsp
2
3d 4s
CN
¯CN ¯CN ¯CN ¯
NC CN
NC CN
2−
4p
hybridization
Ni
(a) Ni
(c) [Ni(CN)
4
]¯
2
(b) Ni
2+
NC CN
NC CN
2−
Ni
PHARMACEuTICAL ORgAnIC CHEMISTRy
TheveF-atomsdirectedtowardstheverticesofaregularpentagonmakeanangleof72ºtooneanother
in one plane. The other two F-atoms are directed at right angle (90º) to this plane. Two axial bonds in the
perpendicularplanehavedifferentbondlengththantheveequatorialbondsinthesameplane.Theaxial
bonds are longer than the equatorial bonds (Fig. 1.27).
2
dsp
7.
Hybridization: Another type of hybridization involving one d orbital of a lower subshell and
one s and two p orbitals of higher subshell are common in case of transition metal ions. The d orbital
involved is d
-d
. These four orbitals intermix and redistribute their energies to give size to four
2
2
x
y
equivalent hybrid orbitals (Fig. 1.28).
Geometry of [Ni(CN)
4]2–
:TheelectroniccongurationofNi(Z = 28) atoms in the ground state is [Ar]18
3d 84s2. The oxidation state of nickel is +2,i.e.theotherelectroniccongurationofNi2+ is 3d 8 as given in
Fig. 1.27 (a). In the complexion, the [ni(Cn)4]2- has square planar geometry as shown in Fig. 1.28 (b).
Figure 1.27 (a) Formation of IF7 molecule involving sp3d 3 hybridization.
Figure 1.28 (a) Formation of [Ni(CN)4]2- involving dsp2 hybridization.
The hybridizations and molecular shapes of some molecules involving d orbitals are summarized in
Table 1.4.
(b) Pentagonal bipyramidal geometry of IF7 molecule.
(b) Structure of [Ni(CN)4]2-.