Добавил:
Sekretar
kiopkiopkiop18@yandex.ru
t.me/Prokururor I Вовсе не секретарь, но почту проверяю
Опубликованный материал нарушает ваши авторские права? Сообщите нам.
Вуз:
Предмет:
Файл:Ординатура / Хирургия / Библиотека им академика М.И. Перельмана / Книга_5574_Библиотеки_им_академика_М_И_Перельмана.pdf
111 1
++=
=
=
maximum multiplicity
.
21
p
1.6 EXTRA STABILITY OF HALF-FILLED AND COMPLETELY FILLED SUBSHELLS
1. Symmetry of orbitals
22
d
4
Config.
4 Exchanges by
electron 1
43
2
1
3 Exchanges by
electron 2
2 Exchanges by
electron 3
1 Exchanges by
electron 4
s d
sd
Exchange energy:
sd
dd5
=+ + 1 = 6.
=+ + + 1 = 10.
sspspsdpsdps f d psf d<< << << < << < << < < << <
MEMORY FOCUS
Electronic Configuration Rules
Electronic configuration of atoms is based on certain rules, namely Aufbau principle, Pauli’s exclusion 1.
principle and Hund’s rule.
x
Electronic configurations are written as 2. nl
x is the number of electrons.
Electrons enter that sublevel first that has lower energy. 3.
The order of energies of various sublevels is:4.
5. (n + l) Rule
a. Electrons enter that sublevel which has least (n + l) value.
b. When two or more sublevels with same (n + l) value are available electron enter that sublevel that
has the least value of n.
c. Among 3d and 4s electrons occupy 4s first.
d. Among 3d, 4p and 5s, electrons after 3d first followed by 4p and then 5s.
6. Pauli’s Exclusion Principle
a. No two electrons in an atom can have all the four quantum numbers identical.
b. There is a room for only two electrons in an orbital and they should have opposite spin.
c. In an atom if electrons have the same n, l and m values, they must differ in spin quantum number.
d. The capacity of any sublevel is determined on the basis of Pauli’s principle.
s sublevel – 1 orbital and 2 electrons
p sublevel – 3 orbitals and 6 electrons
d sublevel – 5 orbitals and 10 electrons
f sublevel – 7 orbitals and 14 electrons
g sublevel – 9 orbitals and 18 electrons
e. The capacity of a main shell is also determined on the basis of Pauli’s principle K shell – 2 electrons,
L shell – 8 electrons, M shell – 18 electrons and N shell – 32 electrons.
7. Hund’s Rule of Maximum Multiplicity
a. Pairing of electrons takes place in a subshell only after all the orbitals are filled with one electron each.
b. Hund’s rule is not valid for s sublevel as it has only one orbital.
c. In p sublevel pairing of electrons takes place with 4th electron.
d. Similarly pairing of electrons starts with 6th electron and 8th electron in d and f sublevels,
respectively.
e. N, P, As, Sb and Bi atoms have three unpaired electrons in their p orbitals because of Hund’s rule.
f. When all the orbitals of the given sublevel are filled with one electron each (half-filled) or two
electrons each (completely filled) that electronic configuration gets extra stability due to symmetry
and exchange energy.
, where n is the main shell, l represents the subshell and
23
24
REVISION QUESTIONS
² in it?
n
=
sp
sp
pp
Z
=Z = 50).
25
(a) H
–
– +
+
+
+
– – +
+
(a) 1s
sp6sp6d
ssp6sp
6
MULTIPLE CHOICE QUESTIONS
d
DxDp = hp
(a) A
.
(a) d Zero
s
(a) A
A
n
(a) A
A
A
A
26
(a) n l
m s
8. the px py pz
(a) s
s
l l + 1
l l
n l m s
+
+
+
N
(Z =
n l m s
+
spdf
p
d f
s d
.
A
s
n is
n n
n n + 1
ssp5s1 shows
(a)
ion
(a) 1ssp6sp6sd1 ssp6sp6s1d
ssp6sp6s0d
ssp6sp6sd
27
l =
(a) s
d
1
1
ssp6p
x
y
p
1
z
(a)
s s
(a) t t
H
28
(a) s
s
s
(a) s
s
s
s
ANSWERS
1. (a) 5. 6. 8. (a) 10. 11.
15. 16. 18. (a)
(a) (a)
29
C H A P T E R 1 Structure and Properties
Chapter Outline (Part B)
1.7 Introduction, 1.8 Valence Bond (VB) Theory, 1.8.1 Principles of VB Theory, 1.8.2 Applications of VB Theory—Formation of H
2
Molecule,
1.8.3 Limitations of VB Theory, 1.9 Molecular Orbital (MO) Theory, 1.9.1 Conditions of Formation of Molecular Orbitals, 1.9.2 Difference Between
Atomic and Molecular Orbitals, 1.10 Formation of Bonding and Antibonding Molecular Orbitals (ICAO Method), 1.10.1 Bonding and Antibonding
Molecular Orbitals in Terms of Wave Functions, 1.10.2 Differences between Bonding and Antibonding Molecular Orbitals, 1.11 Combinations
of Atomic Orbitals—Sigma (σ) and Pi (π) Molecular Orbitals, 1.11.1 Differences between σ and π Molecular Orbitals, 1.12 Comparison
of Valence Bond Theory and Molecular Orbital Theory, 1.13 Hybridization, 1.13.1 Conditions for Hybridization, 1.13.2 Characteristics of
Hybridization, 1.13.3 Types of Hybridization, 1.14 Intermolecular Forces, 1.15 Hydrogen Bonding, 1.15.1 Factors Responsible for the Formation of
Hydrogen Bond and Strength of Hydrogen Bond, 1.15.2 Consequences of Hydrogen Bonding, 1.15.3 Types of Hydrogen Bonds, 1.15.4 Hydrogen
Bonding in Ice in Water, 1.15.5 Importance of Hydrogen Bonding, 1.16 Bond Dissociation Energies
1.7 INTRODUCTION
We know that most of the substances exist in the form of aggregates of atoms called molecules. The attractive
force which holds together the constituents (atoms or ions) in a molecule is called a chemical bond. The
concept of covalent bond formation on sharing of electrons (Lewis concept) was purely a qualitative
approach. It could not answer a number of basic questions such as the following:
1. Why is a covalent bond formed at all?
2. How are electrons arranged in the molecule?
3. What are the forces of attractive interaction in a molecule?
4. What is the bond energy in a covalent bond?
5. How does the sharing of electrons explain the shapes of resulting molecules?
The development of wave mechanics has provided answers to these questions and led to answers to these
questions, which ultimately paved the way for the development of more comprehensive theories of covalent
chemical bonding. These theories are discussed below:
1. Valence Bond Theory (VBT)
2. Molecular Orbital Theory (MOT)
The theory of chemical bond … is still far from
perfect. Most of the principles that have been
developed are crude, and only rarely can they
be used in making an accurate quantitative
prediction. However, they are the best that we
have, as yet, and I agree with Poincare that ‘it
is far better to foresee even without certainty
than not to foresee at all’.
–Linus Pauling
STRUCTURE AND
PROPERTIES
1
PART - B
Formation of Molecules
sp
sp
2
sp
3
dsp
2
dsp
2
3
30
1s atomic orbital
of H atom
1s atomic orbital
of H atom
Orbital overlap
Overlapping of 1s atomic
orbitals to form H
2
HH
HH
+
PHARMACEuTICAL ORgAnIC CHEMISTRy
1.8 VALENCE BOND (VB) THEORY
According to valence bond theory of chemical bonding, developed by Linus Pauling (1930), a covalent bond
is formed by the overlapping of half-lled valence orbitals of two atoms with their electrons having spins in
opposite directions. It is assumed that the atoms interact to form the molecule in which the atomic orbitals
retain a large part of their individual character, i.e. electrons in the molecule are localized as if they are in
isolated atoms.
The main features of this theory are as follows:
1. Acovalentbondisformedwhenahalf-lledvalenceorbitalofoneatomoverlaps(partiallymerges)
withahalf-lledvalenceorbitalofanotheratomofthesameordifferentelement.
2. The two electrons of overlapping atomic orbitals must have their spins in opposite directions.
3. The overlapping of atomic orbitals leads to the formation of a new orbital known as molecular orbitals.
4. During overlapping of atomic orbitals, the magnitude of attractive forces exceeds the magnitude of
repulsive forces. Thus the potential energy of the system decreases, which accounts for the stability
of the molecule.
5. The extent of overlapping of atomic orbitals is a measure of the strength of the covalent bond. The
greater the overlapping of atomic orbitals, the stronger is the covalent bond formed.
6. As a result of overlapping, the electron density becomes maximum in the overlapping space between
the two orbitals. As such, any electron may belong to any one of the atoms at any time. Apart from it,
electrons are attracted to both the nuclei instead of one.
For example, the formation of H
moleculetakesplaceby theoverlapofhall-lled1s atomic orbitals
2
of two hydrogen atoms as shown in Fig. 1.9. After overlapping, the electron density becomes high at the
region of overlap. It means that each of the two electrons spends maximum time in the region of overlap and
is now attracted simultaneously by the two nuclei, resulting in the formation of a new bond between them.
Figure 1.9 Formation of chemical bond by overlap of atomic orbitals.
1.8.1 Principles of VB Theory
If 1.
ψ
and
ψ
A
are the wave functions for two isolated independent systems A and B having energies
B
EA and EB, respectively, then the total wave function, ψ, of the system as a whole can be given by a
product of the two wave functions:
ψ
=
ψA ψ
B