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CHAPTER
Atomic Structure
Bonding
2-1
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Structure of Atoms
ATOMBasic Unit of an Element
10
Neutrally Charged
Diameter : 10 14 m
Accounts for almost all mass
Positive Charge
Mass : 9.109 x 10 28 g
Charge : -1.602 x10 9
CAccounts for all volume
ProtonMass : 1.673 x 10 24 g
NeutronMass : 1.675 x 10 24 g
2-2
arge : . x eu ra arge
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Atomic Number and Atomic Mass
Atomic Number = Number of Protons in the nucleus
Unique to an element Example :- Hydrogen = 1, Uranium = 92
Relative atomic mass = Mass in grams of 6.203 x 1023
Example :- Carbon has 6 Protons and 6 Neutrons. Atomic Mass
= 12.
atom.
One gram mole = Gram atomic mass of an element. Example :-
One gram
12 Grams
6.023 x 1023
Mole of
Carbon
Of CarbonCar on
Atoms2-3
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Periodic Table
Source: Davis, M. and Davis, R., Fundamentals of Chemical Reaction Engineering, McGraw-Hill, 2003.
2-4
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Electron Structure of Atoms
Electron rotates at definite energy levels.
Ener is absorbed to move to hi her ener level.
Energy is emitted during transition to lower level.
Energy change due to transition = E =hc
h=Planks Constant
= 6.63 x 10-34 J.s
EmitAbsorb
c= Speed of light
= Wavelength of light
Energy
(Photon)
nergy
(Photon)
Energy levels
2-6
Photon = Electromagnetic radiation
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Quantum Numbers of Electrons of Atoms
Principal Quantum
Number (n)
Subsidiary Quantum
Number l
Represents main energylevels.
Represents sub energylevels (orbital).
Range 1 to 7.
Larger the n higher
Range 0n-1.
Represented by letters, , .
n=2
s orbital
(l = 0)
n=1
=n=1
p Orbital
(l =1)n=3
2-8
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Electron Structure of Multielectron Atom
Maximum number of electrons in each atomic shell is
given by 2n2.
Atomic size (radius) increases with addition of shells.
Electron Configuration lists the arrangement of electronsin orbitals.
Example :-Orbital letters
Number of Electrons
1s2 2s2 2p6 3s2
=
Principal Quantum Numbers
1s2 2s2 sp6 3s2 3p6 3d6 4s2
2-10
Z = atomic no. = no. of proton
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Electron Structure and Chemical Activity
Except Helium, mostnoble gasses (Ne, Ar, Kr, Xe, Rn)are chemically very stable
All have s
2
p6
configuration for outermost shell. Helium has 1s2 configuration
Electropositive elements give electrons during
Cations are indicated by positive oxidation numbers
Example:-
e : s s sp s p sFe2+ : 1s2 2s2 sp6 3s2 3p6 3d6
Fe3+ : 1s2 2s2 sp6 3s2 3p6 3d5
2-11
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Electron Structure and Chemical Activity (Cont..)
Electronegative elements accept electrons duringchemical reaction.
Some elements behave as both electronegative andelectropositive.
Electronegat v ty s the degree to wh ch the atom
attracts electrons to itself
Measured on a scale of 0 to 4.1
Example :- Electronegativity of Fluorine is 4.1
Electronegativity of Sodium is 1.
0 1 2 3 4K
Na N O Fl
W
Te
SeH
Electro-
positive
Electro-
negative
2-12
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Atomic and Molecular Bonds
Ionic bonds :- Strong atomic bonds due to transfer of
- oms n on e s a e are n a more s a e energy con on compare o un on e
condition. Net decrease of potential energy after bonding.
electrons non-directional
Covalent bonds :- Large interactive force due to
-
Metallic bonds :- Non-directional bonds formed by
sharing of electrons
Permanent Dipole bonds :- Weak intermolecular
bonds due to attraction between the ends of permanent
. Fluctuating Dipole bonds :- Very weak electric dipole
bonds due to asymmetric distribution of electron
densities.
2-12
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Ionic Bonding
Ionic bonding is due to electrostatic force of attraction
between cations and anions.
elements.
Electrons are transferred from electropositive to
electronegative atoms
ElectropositiveElement ElectronegativeAtomElectronrans er
Electrostatic (coulombic)
Cation
+ve charge
Anion
-ve charge
Attraction
IONIC BOND
2-14
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Ionic Bonding - Example
Ionic bonding in NaCl
1
Atomic no= 17
3p7
Sodium
Atom
ChlorineAtom
=
Na
Chlorine Ion
O
N
I
Na+ -
C
B
N
D2-15
Figure 2.10
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Ionic Force for Ion Pair
Nucleus of one ion attracts electron of another ion.
The electron clouds of ion repulse each other when
(Proton + neutron)
they are sufficiently close.
Force versus separation
oppositely charged ions
Figure 2.11
2-16
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Ion Force for Ion Pair (Cont..)
( )( )
( ) ( )a
eZZ
a
ZZF
ee
attractive2
2
21
2
21
44
==
Z1,Z2 = Number of electrons removed oradded during ion formation
=
a = Interionic seperation distance
= Permeability of free space (8.85 x 10-12c2/Nm2)
(n and b are constants)Fn
repulsive
nb
1+
=
eZZF
nnet
nb
12
2
21
+
=
aa0
2-17
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Interionic Force - Example
Force of attraction between Na+ and Cl- ions
Z1 = +1 for Na+
, Z2 = -1 for Cl-
e = 1.60 x 10-19 C , 0 = 8.85 x 10-12 C2/Nm2
a0 = Sum of Radii of Na+ and Cl- ions
= 0.095 nm + 0.181 nm = 2.76 x 10-10
m
Na+ Cl-
CeZZ 92192
21)1060.1)(1)(1(
+
a0
( )aattraction
10-212-2
0
.
m)10x/Nm2)(2.76C10x8.85(44
2-18
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Ion Arrangements in Ionic Solids
Ionic bonds areNon Directional
Geometric arrangements are present in solids to
maintain electric neutrality.
Example:- in NaCl, six Cl- ions pack around central Na+ Ions
Ionic packing
In NaCl
and CsCl
Figure 2.13
As the ratio of cation to anion radius decreases, fewer
CsCl NaCl
anion surround central cation.
2-20
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Bonding Energies
Lattice energies and melting points of ionically
bonded solids are high.
(Measure of bonding strength)
increases (because bonding electrons in larger ions are farther away fromthe attractive influence of the +ve nucleus).
.
Example :-NaCl Lattice energy = 766 KJ/mol
Melting point = 801oC
CsCl Lattice energy = 649 KJ/mol
Melting Point = 646oCBaO Lattice energy = 3127 KJ/mol
Melting point = 1923oC
2-21
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Covalent Bonding
In Covalent bonding, outer s and p electrons are
shared between two atoms to obtain noble gas
con gura on.
Takes place between elements
electronegativity and close by
in eriodic table.
In Hydrogen, a bond is formed between 2 atoms by
sharing their 1s1 electronsElectron Overlapping Electron Clouds
H + H H HPair
H H
1s1
Electrons
Hydrogen
Molecule2-22
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Covalent Bonding - Examples
In case of F2, O2 and N2, covalent bonding is formed bysharing p electrons
uor ne gas u er or a s p s are one p e ec ron o
attain noble gas configuration.
F + F F FH
F FBond Energy=160KJ/mol
Oxygen (Outer orbital - 2s2 2p4) atoms share two p electrons
O + O O O O = O
Nitrogen (Outer orbital - 2s2 2p3) atoms share three p electrons
Bond Energy=28KJ/mol
H H N + N Bond Energy=54KJ/molN N N N
2-23
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Covalent Bonding in Carbon
Carbon has electronic configuration 1s2 2s2 2p2
roun tate arrangement Ind cates
carbon
Forms two1s 2s 2p
Two filed 2p orbitals
Covalent
bonds
Hybridization causes one of the 2s orbitals promoted to
2p orbital. Result four sp3 orbitals.
Indicatesfour covalent
bonds ares 2p
Four filled sp3 orbitals
formed
2-24
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Covalent Bonding in Benzene
Chemical composition of Benzene is C6H6.
The Carbon atoms are arranged in hexagonal ring.
Single and double bonds alternate between the atoms.H
CH H
C C
CH
H
Structure of Benzene Simplified Notations
2-27
.
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Metallic Bonding
Atoms in metals are closely packed in crystalstructure.
towards nucleus of other atoms. Electrons spread out among atoms forming electron
clouds.
These free electrons arePositive Ion
conductivity and ductility*
Since outer electrons are
shared by many atoms,metallic bonds are
-*a mechanical property used to describe the extent to which materials can be deformed plastically without
fracture Valence electron charge cloud2-28
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Metallic Bonds (Cont..)
Overall energy of individual atoms are lowered by
metallic bonds (i.e. from unstable atoms to stable bonding).
Minimum energy between atoms exist at equilibrium
distance a0
,
metallic the bond is (i.e. valence electrons are freer to move).
Example:- Na Bonding energy 108KJ/mol,Melting temperature 97.7oC
Higher the number of valence electrons involved,
.
Example:- Ca Bonding energy 177KJ/mol,
Melting temperature 851oC
2-29
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Secondary Bonding
Secondary bonds are due to attractions of electricdipoles in atoms or molecules.
centers exist.
-
Dipole moment= =q.d
nucleus
q= ec r c c arge
d = separation distance
There two types of bonds permanent and.
2-30
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Fluctuating Dipoles
Weak secondary bonds in noble gasses.
Dipoles are created due to asymmetrical distribution of
They have complete outer-
valence-electron shells
electron charges.
Electron cloud charge changes with time.
nucleus
+_
Symmetrical
AsymmetricalFigure 2.27
of electron charge (Changes with time)
2-31
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Permanent Dipoles
Dipoles that do not fluctuate with time are
called Permanent dipoles (weak bonding forces among covalently.
Examples:-
Arrangement
Of 4 C-H bondsCH4
moment
(Methane)
Asymmetrical
TetrahedralCreates
arrangementpo e
(Chloromethane)
2-32
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Hydrogen Bonds
Hydrogen bonds areDipole-Dipole interaction
between polar bonds containing hydrogen
spec a case o permanen po e- po e n erac on e w. po ar mo ecu es
atom. Example :-
In water, dipole is created due to asymmetrical
arrangement of hydrogen atoms.
Attraction between positive oxygen pole andnegative hydrogen pole.
H
105 0O
H dro enFigure 2.28
H Bond
2-33