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Файл:Ординатура / Хирургия / Библиотека им академика М.И. Перельмана / Книга_5574_Библиотеки_им_академика_М_И_Перельмана.pdf
C H A P T E R 1 Structure and Properties
N
O
O
O
O
O
. . . .
H
O
. . . .
H
O
. . . .
H
N
O
N
O
. . . .
H – O
. . . .
H – O
H
N
. . . .
H – O
. . . .
O
HO
Boiling point and solubilities of nitrophenols
Due to intermolecular hydrogen bonding in p-nitrophenol, association of molecules takes place and thereby
their bonding points increases.
o
-nitrophenol due to intramolecular hydrogen bonding (i.e. chelation prevents) prevents the association
of molecules and therefore has a low boiling point.
A high solubility of p-nitrophenol in water is due to greater extent of hydrogen bonding between –OH and
nO2. The low solubility of o-nitrophenol in water is due to suppression of hydroxyl character as a result of
intramolecular hydrogen bonding (i.e. chelation).
61
1.15.4 Hydrogen Bonding in Ice in Water
In water, hydrogen bonds are formed between oxygen atom of one molecule and hydrogen atoms of another
moleculeatdeniteangles.
In ice, water molecules are tetrahedrally linked in which two hydrogen atoms form covalent bonds with
oxygen atom while the other two hydrogen atoms form hydrogen bonds (Fig. 1.32). This results in the
formation of ‘open cage-like structure’ as revealed by X-ray. In ice, covalently bonded hydrogen are at
lesser distance (10 Å), whereas other two hydrogen bonds lie at a distance of 1.76 Å. It results in vacant
spaces in crystal structure and, therefore, gives rise to large volume for a given mass of water. Therefore,
density is mass per unit volume.
Density (= mass/volume) is less than that of water
When ice melts, the open cage-like structure (Fig. 1.33) partially breaks, which brings the water molecules
closer by entering into the vacant spaces. As a result, volume decreases and density increases in the liquid
state of water.
62
H
H
H
H
H
H
H
H
O
O
O
O
H
H
1.76Å
2.76Å
1.0Å
H
HH HH H
H
HH H
Vacant
space
Vacant
space
Vacant
space
H
H
HH
HH
OO
OOO
O
O
O
PHARMACEuTICAL ORgAnIC CHEMISTRy
Figure 1.32 Tetrahedral arrangement of water molecules in ice.
Density of Water is Maximum at 4°C
As water is heated form 0°C (melting point of ice) onwards, more and more of hydrogen bonds are broken
and the molecules come more and more closer in the liquid state. As a result, the volume decreases and hence
density increases. This process continues up to 4°C where water attains the maximum density. On heating
further, the kinetic energy of the molecules increases and the molecules move apart causing expansion. So
up to 4°C, contraction effect predominates and above 4°C, the expansion dominates over contraction and
overall density decreases. So the value of density is maximum at 4°C.
Figure 1.33 Open cage-like structure of ice.
63
C H A P T E R 1 Structure and Properties
1.15.5 Importance of Hydrogen Bonding
1. Without hydrogen bonding, water would have existed as a gas and no life would have been
possible.
2. Hydrogen bonding exists in all living organisms like various tissues, blood, skin muscles and
bones.
3. Hydrogenbondinggivesextrarigiditytowood,cottonandsyntheticbres.
4. Carbohydrates and sugar, which are an integral part of our diet, consist of hydrogen-bonded mol-
ecules.
5. Honey and glycerol are viscous due to hydrogen bonding.
1.16 BOND DISSOCIATION ENERGIES
Bond energy or bond strength:Itisdenedastheamountofenergyrequiredtobreakonemoleofthebond
(i.e. Avogadro’s number of bonds) and separate the bonded
atoms in the gaseous state. It is also known as the bond
dissociation energy (D) of that particular bond. Each bond in
a molecule has a characteristic value of energy . In general, the
greater the bond dissociation energy, the stronger and hence
more stable (less reactive) is that bond.
It is important to note that bond dissociation energy is
characteristic of a particular bond (not a particular type of
bond). For example, the four C–H bonds in the methane (CH4)
have different values of D. These are 104 (for H3C–H), 106 (for H2C–H), 106 (for HC–H) and 81 kcal/mole
(for C–H). It should be noted that the dissociation bond energy (denoted by D) should not be confused with
bond energy (denoted by E). Actually the bond energy is the average of the various values of dissociation
energies of the same type of bonds present in the molecule. Thus the bond energy of C–H bond in methane is
99.2 kcal/mol (104 + 106 + 106 + 81 = 397; 397/4 = 99.2).
Hence in polyatomic molecules, bond dissociation energy is considerably different from that of the average
bond energy, but even then for practical purposes, the average bond energy is taken as the dissociation bond
energy. However, in case of diatomic molecules the two values are identical.
Important features among bond energies: Bond energy depends upon the following factors:
1. Atomic size of bonded atoms: Since we know that bond distance
∝ atomic size of the bonded atoms.
Hence smaller atoms form shorter bonds whose energy will be large. This is evident from the following
order of bond energies of the halogens:
Cl–Cl > Br–Br > I–I
Learning Plus
Bond energy: Single bond < Double bond < Triple bond
Rigidity: Single bond < Double bond < Triple bond
Stability: Single bond < Double bond < Triple bond
Reactivity: Single bond < Double bond < Triple bond
Bond length: Single bond > Double bond > Triple bond
64
C
81.6 39 34.2 33.3 kcal mol
−1
CN
:
:
:
:
::
: :
::::
NOOFF
PHARMACEuTICAL ORgAnIC CHEMISTRy
Recall that the atomic size of the three halogens follows the order:
Cl < Br < I
2. Electronegativity of the bonded atoms (bond polarity): greater the electronegativity difference,
greater is the bond polarity and hence greater will be the bond strength, i.e. bond energy. This is evident
from the following order of bond energies of the different hydrogen halides:
H – F > H – Cl > H – Br > H – I
3. Extent of overlapping of atomic orbitals: A larger extent of overlapping of the component atomic
orbitals imparts greater strength to the bond.
4. Hybridization: Hybrid orbitals form stronger bonds because they provide more extent of overlapping
than the pure atomic orbitals. Thus sp3–sp3overlapping(i.e.σbond)resultsinastrongerbondthanthe
p–poverlapping(i.e.πbond).
5. Percentage of s-character in a hybrid bond: The bond energy increases with the increases in the
percentage of s-character in a hybrid orbital. Thus bond energy increases in the following order:
% of s-character 25 33.3 50
sp
3
<
2
sp
<
sp
6. Bond order: Since bond energy ∝ bond order, bond energy increases from a single bond to a triple
bond, i.e.
7. The bond energy decreases with the increase in number of lone pairs on the bonded atom. For example,
bond energies of the following single bonds having zero, one, two and three lone pairs of electrons
follow the following order.
This is due to the presence of electrostatic repulsion between lone pairs of electrons on the two bonded
atoms.
The homolytic and heterolytic bond dissociation energies are considerably different; generally the
valuesfortheformerarelow.Forexample,theC–Brbondenergyinhomolyticandheterolyticssion
of C2H5Br is 67.2 and 183 kcal mol
C – C < C = C < C C
-1
, respectively.
65
C H A P T E R 1 Structure and Properties
MEMORY FOCUS
Bond energy may be defined as the amount of energy required to break the bond between two 1.
atoms.
Bond energy is directly proportional to bond strength and thus bond energy gives a measure of the 2.
strength of the bond.
Bond energy is also known as 3. bond dissociation energy and is measured in kJ mole
–1
.
The smaller the size of the atoms, greater is the value of bond energy. 4.
(Cl - Cl > Br - Br > I - I)
The shorter the bond length, greater is the bond energy ( C 5. ≡ C > C = C > C – C). The strength of the
bond indicates the stability of the bond. Thus C = C bond is more stable than C-C.
Nitrogen molecule is more stable than oxygen molecule. F6. -F bond is less stable than Cl-Cl bond,
which indicates that flourine is more reactive than chlorine.
Greater the number of bonds or higher the bond order, the greater the bond energy. 7.
Greater the number of lone pairs at the two bonding atoms, greater is the repulsion between them 8.
and weaker is the bond.
REVISION QUESTIONS
1. Dene or explain the following terms:
(a) Octet rule (b) Ionic bond
(c) Covalent bond (d) Coordinate covalent bond
(e) Polar covalent bond (f) VSEPR theory
2. (a)Whatdoyouunderstandbystableconguration?Whatarethewaysbywhichanatomcanattain
stableconguration?
(b) Writetheelectroniccongurationofanytwoofthefollowingcompounds:
(i) Phosphorus pentachloride
(ii) Sulphuric acid
(iii) Lithiumuoride
3. What type of bonds do you expect in the following cases? give reason.
(a) Between a very small cation and a large anion
(b) Between atoms having a very large difference in electronegativities
(c) Between atoms of the same element.
4. Explain qualitatively the valence bond theory with reference to hydrogen molecule.
5. Compare the properties of ionic and covalent compounds. give two examples of each type of
compounds.
VI
66
PHARMACEuTICAL ORgAnIC CHEMISTRy
6. Write Lewis dot formulae of
(a) HOCl (b) BF3 (c) nH
+
4
7. Why bond angles of H2O and nH3 are 104º and 107º, respectively, although central atoms are sp
hybridized.
8. Explain the following terms:
(a) Hybridization (b) Bonding molecular orbital
(c) Antibonding molecular orbital (d) Bond order
9. What is meant by hybridization? Explain the hybridization involved in
(a) beryllium dichloride
(b) borontriuoride
10. The molecules of CH4 (methane), nH3 (ammonia) and H2O (water) all involve sp3 hybridization
of the central atom but the bond angles in these molecules are 109º28′, 107º18′ and 104º30′,
respectively. Explain.
11. Explain the following on the basis of molecular orbital theory
(a) Br2 molecule is not stable.
3
(b) n2 has a stronger bond than B2.
(c) F2 is diamagnetic while O2 is paramagnetic.
12. (a) What is a molecular orbital? In what ways is a molecular orbital different from a valence bond
orbital?
(b) Sketch and describe the molecular orbital resulting from the following overlaps of atomic
orbitals:
(i) Two s orbitals
(ii) Head-on overlap of two p orbitals
13. (a) What is hybridization? Write down the rules for hybridization of orbitals.
(b) CH4, nH3 and H2O have tetrahedral geometry yet their bond angles are different. Why?
14. Describe sp3, sp2 and sp hybridization of carbon in organic molecules. Draw the model of acetylene
andlabelthebondassigma(σ)andpi(π)bonds.
15. (a) Explain the terms bonding and molecular orbitals with examples.
(b) Explain with the help of MO theory, the paramagnetic character of oxygen.
16. (a) Whatareσandπmolecularorbitalsandgivetheircharacteristics?
(b) Explain the formation of HF molecule with MO diagram.
17. What is pi bond? Explain the formation of pi bond in ethylene and acetylene.
18. (a) What are the characteristics of hybrid orbitals?
67
C H A P T E R 1 Structure and Properties
(b) What is tetrahedral hybridization? Explain the formation of C–H bonds in CH4.
(c) Whataresigma(σ)andpi(π)bonds?Distinguishbetweenthem.
19. (a) Compare VBT and MOT.
(b) What are the differences between bonding and antibonding molecular orbitals?
(c) Explain the structure of SF
4
on the basis of hybridization.
20. What do you understand by LCAO in molecular orbital theory? Differentiate bonding, antibonding
and nonbonding molecular orbital.
MULTIPLE CHOICE QUESTIONS
1. The total number of electron pairs in a nitrogen molecule is
(a) 2 (b) 3
(c) 5 (d) 7
2. The covalent compounds are soluble in
(a) All acids (b) All bases
(c) All solvents (d) nonpolar solvents
3. The compounds which contain both ionic and covalent bonds are
(a) CHCl
3
and CCl4 (b) KClandAlCl
3
(C) KCNandNaOH (d) H2 and CH
4
4. The polarity of a covalent bond is due to
(a) Lesser electronegativity difference between two atoms
(b) greater electronegativity difference between two atoms
(c) Lesser bond energy
(d) greater bond energy
5. A CO
2
molecule contains two polar bonds but the net dipole moment is zero . It is because
(a) The molecule has symmetrical linear geometry
(b) The molecule is nonlinear
(c) The electronegativity difference between the two atoms is too large
(d) The electronegativity difference between the two atoms is too small
6. Methanol is soluble in water due to
(a) Covalent bond nature (b) Ionic bond nature
(c) Hydrogen bonding (d) Its poisonous nature
68
PHARMACEuTICAL ORgAnIC CHEMISTRy
7. According to VSEPR theory
(a) The lone pairs only decide the structure of the molecule
(b) The bond pairs only decide the structure of the molecule
(c) The lone pairs and bond pairs both decide the structure of the molecule
(d) none of these
8. In which of the following, the central atom is surrounded by four electron pairs:
(a) H2O (b) nH
3
(c) CH4 (d) All
9. CO2 has zero dipole moment whereas H2O has a dipole moment. It is because
(a) H2O is linear while CO2 is a bent molecule
(b) Of intermolecular H-bonding in H2O molecules
(c) CO2 is linear while H2O is a bent molecule
(d) CO2 is a gas while H2O is a liquid at room temperature
10. Which of the following does not obey the octet rule?
(a) PCl5 (b) H2O
(c) nH3 (d) CCl
4
11. In a bond between two atoms X and y, the shared electron pair does not lie in the centre. The bond
is
(a) Single bond (b) nonpolar bond
(c) Polar bond (d) Coordinate bond
12. According to valence bond theory, a bond between two atoms is formed when
(a) Half-lledatomicorbitalsoverlap (b) Fullylledatomicorbitalsoverlap
(c) nonbonding atomic orbitals overlap (d) Electrons of the two atoms overlap
13. The axial overlap between the two orbitals leads to the formation of a
(a) Sigma bond (b) Pi bond
(c) Multiple bond (d) none of these
14. In a n2 molecule there are
(a) One sigma and one pi bond (b) Two sigma and one pi bond
(c) One sigma and two pi bond (d) Three sigma bonds
15. Out of the following which statement is incorrect?
(a) Half-lledorbitalsonlytakepartinhybridization
C H A P T E R 1 Structure and Properties
(b) The orbitals taking part in hybridization should have very little energy difference
(c) The number of hybrid orbitals formed is equal to the number of orbitals intermixed
(d) The energy of hybrid orbitals is less than that of atomic orbitals
16. When one s and two p orbitals hybridize we get
(a) Three new orbitals at 90º to each other
(b) Three new orbitals at 120º to each other
(c) Two new orbitals at 180º to each other
(d) Two new orbitals at 90º to each other
17. In H2O, nH3 and CH4 molecules the oxygen, nitrogen and carbon atom are
(a) sp3 hybridized (b) sp3, sp2 and sp, respectively
(c) sp, sp2, sp3, respectively, hybridized (d) sp2 hybridized
18. In dsp2 hybridization, the new orbitals have the following geometry:
(a) Square planar (b) Tetrahedral
(c) Trigonal (d) Trigonal bipyramidal
69
19. In PCl5 molecule the phosphorus atom is
(a) sp3 hybridized (b) sp3d hybridized
(c) sp3d2 hybridized (d) sp2 hybridized
20. Valence bond theory was proposed by
(a) Rutherford (b) neils Bohr
(c) Heitler and London (d) Hund and Mulliken
21. The molecule orbitals formed as a result of LCAO method obey
(a) Pauli’s exclusion principle (b) Hund’s rule of maximum multiplicity
(c) Aufbau principle (d) All of these
22. The energy of atomic orbitals taking part in molecular orbital formation
(a) is equal to the molecular orbitals formed
(b) is less than the molecular orbitals formed
(c) is greater than the molecular orbitals formed
(d) cannot be predicted
23. A molecule or ion is stable if
(a) Nb= Na (b) Nb< N
a
(c) Na< Nb (d) Na - Nb = +ve
70
PHARMACEuTICAL ORgAnIC CHEMISTRy
24. Which of the following is not true regarding LCAO method?
(a) The energies of atomic orbitals should be comparable
(b) The atomic orbitals should overlap to a considerable extent
(c) The symmetry of the combining orbitals should be the same
(d) The energy of resulting antibonding orbitals is less than that of bonding orbital
25. Which is not true for bonding MO?
(a) It is formed by the addition overlap of atomic orbitals
(b) The wave function of a bonding MO is given by
ψ
=
ψ
-
ψ
MO
A
B
(c) The lobes of atomic orbitals should have the same signs
(d) Every electron in bonding MO contributes towards the attractive force
ANSWERS
1. (d) 2. (d) 3. (c) 4. (b) 5. (a) 6. (c) 7. (c) 8. (d) 9. (c) 10. (a) 11. (c) 12. (a) 13. (a)
14. (c) 15. (a) 16. (b) 17. (a) 18. (d) 19. (b) 20. (c) 21. (d) 22. (c) 23. (b) 24. (d) 25. (b)