Chapter 2
Chapter 2
Chapter 2
ISSUES TO ADDRESS...
• What promotes bonding?
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Introduction
• Properties of materials depend on the ATOMIC STRUCTURE
and INTERATOMIC BONDING (interactions existing
between constituent atoms or molecules)
• GRAPHITE Vs DIAMOND:
– Both are forms of CARBON
– Graphite: Soft and Greasy whereas Diamond: Hardest known Material
– Different behavior is due to different bonding characteristics
• Atomic Structure:
Subatomic Particles:
Electrons: -vely Charged
Protons: +vely Charged
Neutrons: Electrically Neutral
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Definitions
• Atomic Number Z: Number of protons in nucleus
• Atomic Mass A: Sum of masses of protons and neutrons
within nucleus
• Isotopes: Number of neutrons N may be different forming
different isotopes
• Atomic Mass Unit amu:
– 1 amu = 1/12 atomic mass of the most common isotope of Carbon
– 1 amu/atom (or molecule) = 1 g/mole
– In one mole of a substance we have 6.023 x 1023 atoms
• ELECTRONS IN ATOMS:
– Bohr Atomic model based on the principle of quantum mechanics
– Electrons revolve in discrete orbitals with specific value of energy
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BOHR ATOM
orbital electrons:
n = principal
quantum number 1
n=3 2
Nucleus: Z = # protons
= 1 for hydrogen to 94 for plutonium
N = # neutrons
Atomic mass A ≈ Z + N
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Quantum Numbers
• Size, Shape and Spatial Orientation of an Electron probability
density are specified by FOUR Quantum Numbers
1. Principal Quantum Number n: Specifies distance from Nucleus
• K, L, M, N, O or 1, 2, 3, 4
2. Sub-shell Quantum Number
: Specifies shape of electron sub-
shell
• s, p, d, f
3. Number of Energy States ml: Each electron state cannot hold
more than two electrons with s=1, p=3, d=5 and f=7 energy
states
4. Spin Moment Quantum Number ms : Oriented either up or
down
• ms =1/2 or -1/2
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ELECTRON ENERGY STATES
Electrons...
• have discrete energy states
• tend to occupy lowest available energy state.
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• VALENCE ELECTRONS:
– Electrons occupying the outermost filled shell
– Very important since they participate in the bonding between atoms to
form atomic and molecular aggregates
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STABLE ELECTRON CONFIGURATIONS
Stable electron configurations...
• have complete s and p subshells
• tend to be unreactive.
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SURVEY OF ELEMENTS
• Most elements: Electron configuration not stable.
Electron configuration
1s1
1s2 (stable)
1s22s1
1s22s2
1s22s22p1
1s22s22p2
...
1s22s22p6 (stable)
1s22s22p63s1
1s22s22p63s2
1s22s22p63s23p1
...
1s22s22p63s23p6 (stable)
...
1s22s22p63s23p63d104s246 (stable)
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Bonding Forces and Energies
• As two isolated atoms are brought closer from an infinite
separation following effects will be observed:
– At large distance no interaction
– As they come closer two forces are experienced
– (1) Attractive Forces FA and (2) Repulsive Forces FR
– In beginning Attractive Forces are important, however, Repulsive
Forces become significant when outer electron shells of two atoms
overlap
• Net Force FN =FA + FR
• When FA+FR = 0 No Net Force State of equilibrium; the
two atoms are separated by an equilibrium spacing ro (ro 3A)
• Distance less than ro Repulsive forces will dominate, distance
more than ro attractive forces will dominate
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Bonding Forces and Energies
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Bonding Energies
• In terms of Potential energy instead of Forces:
E N FN dr
r r
E N FAdr FR dr E A E R
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Properties from Bonding: TM
• Bond length, r • Melting Temperature, Tm
F
F
• Bond energy, Eo
Tm is larger if Eo is larger.
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Properties from Bonding: E
• Elastic modulus, E
Elastic modulus
F L
=E
Ao Lo
• E ~ curvature at ro
Energy
unstretched length
ro E is larger if Eo is larger.
r
smaller Elastic Modulus
• a ~ symmetry at ro
a is larger if Eo is smaller.
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Types of Bonds
• PRIMARY OR CHEMICAL BONDS:
– Ionic
– Covalent
– Metallic
(Based on the atoms assuming stable electronic structure by filling
outer most electron shell; Strong Bonds)
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IONIC BONDING
• Occurs between + and - ions.
• Requires electron transfer.
• Large difference in electronegativity required.
• Example: NaCl
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Ionic Bonding
• Present in compounds between metallic and non-metallic
elements
• Metals gives up electrons and non-metal gains electrons
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COVALENT BONDING
• Requires shared electrons
• Example: CH4
C: has 4 valence e,
needs 4 more
H: has 1 valence e,
needs 1 more
Electronegativities
are comparable.
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Covalent Bonding
• Stable Electron Configuration assumed by sharing of electrons
between adjacent atoms
• The bonded atoms will share at least one electron and shared
electron may be considered to belong to both atoms
• Covalent bonds are directional
• Present in dissimilar molecules such as CH4 and elemental
solids such as Diamond (Carbon) and Silicon, and in Semi-
conductor Materials
• Number of covalent bonds present depend on number of
valence electrons
• Covalent bonds can be very strong and weak depending on
atomic arrangements
• Polymer have covalent bonds in their main chain
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Examples: Covalent Bonding
H2O
column IVA
H2 F2
C(diamond)
H He
2.1
SiC - Cl2
Li Be C O F Ne
1.0 1.5 2.5 2.0 4.0 -
Na Mg Si Cl Ar
0.9 1.2 1.8 3.0 -
K Ca Ti Cr Fe Ni Zn Ga Ge As Br Kr
0.8 1.0 1.5 1.6 1.8 1.8 1.8 1.6 1.8 2.0 2.8 -
Rb Sr Sn I Xe
0.8 1.0 1.8 2.5 -
Cs Ba Pb At Rn
0.7 0.9 1.8 2.2 -
Fr Ra
0.7 0.9 GaAs
• PERMANENT DIPOLES:
– Due to unsymmetrical arrangements of positive and negatively charged
regions
– Bonding energies are greater than Fluctuating Dipoles. E.g: HCl
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Van Der Waals BONDING
Arises from interaction between dipoles
• Fluctuating dipoles
-ex: polymer
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• HYDROGEN BONDS:
– A special case of secondary polar molecule bonding
– E.g: HF, H2O, NH3
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• Intramolecular bonds within a polymer chains are covalent (e.g:C-C and C-H)
• Intermolecular bonds betweens two polymer chains are weak hydrogen and
van der Waals 33
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SUMMARY: BONDING
Type Bond Energy Comments
Ionic Large! Nondirectional (ceramics)
Variable Directional
Covalent large-Diamond semiconductors, ceramics
small-Bismuth polymer chains)
Variable
Metallic large-Tungsten Nondirectional (metals)
small-Mercury
Directional
Secondary smallest inter-chain (polymer)
inter-molecular
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SUMMARY: PRIMARY BONDS
Ceramics Large bond energy
(Ionic & covalent bonding): large Tm
large E
small a
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