CHAPTER
2
Atomic Structure
And
Interatomic Bonding
Chapter 2: Atomic Structure &
Interatomic Bonding
ISSUES TO ADDRESS...
• What promotes bonding?
• What types of bonds are there?
• What properties are inferred from bonding?
2
TABLE OF CONTENT
•
•
•
•
•
•
•
•
•
•
Introduction
Fundamental concepts
Electrons in atoms
The periodic table
Bonding forces and energies
Primary interatomic bonds
Secondary bonding or van der Waals bonding
Mixed Bonding
Molecules
Bonding type-Material classification correlations
3
Learning Objectives
• Name the two atomic models cited, and note the differences
between them.
• Describe the important quantum-mechanical principle that
relates to electron energies.
• Schematically plot attractive, repulsive, and net energies
versus interatomic for two atoms or ions.
• Note on the plot the equilibrium separation and the bonding
energy.
• Briefly describe ionic, covalent, metallic, hydrogen, and van
der Waal's bonds.
• Note which materials exhibit each of these bonding types.
4
TABLE OF CONTENT
•
•
•
•
•
•
•
•
•
•
Introduction
Fundamental concepts
Electrons in atoms
The periodic table
Bonding forces and energies
Primary interatomic bonds
Secondary bonding or van der Waals bonding
Mixed Bonding
Molecules
Bonding type-Material classification correlations
5
Atomic Structure and Interatomic Bonding
Gecko, harmless tropical lizards, have very sticky feet that cling
to virtually any surface. In fact, a gecko can support its body mass
with a single toe!
The secret to this remarkable ability is the presence of an extremely
large number of microscopically small hairs on each of their toe pads.
When the hairs come in contact with a surface, weak forces of
attraction are established between hair molecules and molecules on
the surface. The fact that these hairs are so small and so numerous
explains why the gecko grips surfaces so tightly.
To release its grip, the gecko simply curls up its
toes and peels the hairs away from the surface.
Using their knowledge of this mechanism of
adhesion, scientists have developed several
Ultra-strong synthetic adhesives.
6
Structure of Atoms
ATOM
Basic Unit of an Element
Diameter : 10 –10 m.
Neutrally Charged
Nucleus
Electron Cloud
–14
Diameter : 10
m
Accounts for almost all mass
Positive Charge
Mass : 9.109 x 10 –28 g
Charge : -1.602 x 10 –19 C
Accounts for all volume
Proton (2 up + 1 d quarks)
Neutron (1up+2d quarks)
Mass : 1.673 x 10 –24 g
Charge : 1.602 x 10 –19 C
Mass : 1.675 x 10 –24 g
Neutral Charge
7
Atomic Structure (Freshman Chem.)
• atom –
electrons – 9.11 x 10-31 kg
protons
1.67 x 10-27 kg
neutrons
}
• atomic number = # of protons in nucleus of atom
= # of electrons of neutral species
• A [=] atomic mass unit = amu = 1/12 mass of 12C
Atomic wt = wt of 6.022 x 1023 molecules or atoms
1 amu/atom = 1g/mol
C 12.011
H 1.008 etc.
8
Atomic Structure
• Valence electrons determine all of the
following properties
1)
2)
3)
4)
Chemical
Electrical
Thermal
Optical
9
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 1 mole atoms
6.02214179 x 1023 ( Avagadro Number)
 Example :- Carbon has 6 Protons and 6 Neutrons. Atomic Mass = 12.
•
One Atomic Mass unit is 1/12th of mass of carbon atom. (=
1.66 x 10-24 g)
• One gram mole = Gram atomic mass of an element.
 Example :-
One gram
Mole of
Carbon
12 Grams
Of Carbon
10
6.022 x 1023
Carbon
Atoms
Example Problem
• A 100 gram alloy of nickel and copper consists of 75 wt%
Cu and 25 wt% Ni. What are percentage of Cu and Ni
Atoms in this alloy?
Given:- 75g Cu
Atomic Weight 63.54
25g Ni
Atomic Weight 58.69
• Number of gram moles of Cu =
• Number of gram moles of Ni =
75g
 1.1803mol
63.54 g/mol
25g
 0.4260mol
• Atomic Percentage of Cu =
58.69 g/mol
1.1803
 100  73.5%
(1.1803  0.4260)
• Atomic Percentage of Ni =
0.4260
100  26.5%
(1.1803  0.4260)
11
TABLE OF CONTENT
•
•
•
•
•
•
•
•
•
•
Introduction
Fundamental concepts
Electrons in atoms
The periodic table
Bonding forces and energies
Primary interatomic bonds
Secondary bonding or van der Waals bonding
Mixed Bonding
Molecules
Bonding type-Material classification correlations
12
Materials Science and Engineering
Electrons in Atoms
魏茂國
 Two atomic models
- Bohr atomic model
- Wave-mechanical model
 Bohr atomic model
- Electrons are assumed to revolve around the atomic
nucleus in discrete orbitals, and the position of any
particular electron is more or less well defined in
terms of its orbital.
Figure 2.1 Schematic
- The Bohr model represents an early attempt
representation of the Bohr atom.
to describe electrons in atoms, in terms of
both position (electron orbitals) and energy
(quantized energy levels).
13
Materials Science and Engineering
Electrons in Atoms
 Quantum-mechanical
魏茂國
(量子力學) principle
- The energies of electrons are quantized; that is, electrons are
permitted to have only specific values of energy.
- An electron may change energy, but in doing so it must make a
quantum jump either to an allowed higher energy or to a lower
energy.
- It is convenient to think of these allowed electron energies as
being associated with energy levels (能階) or states (能態).
- In Fig 2.2a, the energies are taken to be negative, whereas the
zero reference is the unbound or free electron.
14
The first three electron energy state in hydrogen Atom
Materials Science and Engineering
魏茂國
Figure 2.2 (a) the Bohr hydrogen atom. (b) the wave-mechanical hydrogen atom.
15
Materials Science and Engineering
 Wave-mechanical
Electrons in Atoms
魏茂國
(波動力學) model
- The electron is considered to exhibit both wave-like and
particle-like characteristics.
- An electron is no longer treated as a particle moving in a
discrete orbital; rather, position is considered to be the probability
of an electron’s being at various locations around the nucleus.
In the other words, position is described by a probability
distribution or electron cloud.
16
Materials Science and Engineering
Electrons in Atoms
魏茂國
Figure 2.3 (a) the Bohr and (b) the wavemechanical models in terms of electron
distribution
Figure 2.3 Comparison of the (a) Bohr
and (b) wave-mechanical atom models
in terms of electron distribution.
17
Electronic Structure
• Electrons have wavelike and particulate properties.
– This means that electrons are in orbitals defined by a
probability.
– Each orbital at discrete energy level is determined by
quantum numbers.
Quantum #
Designation
n = principal (energy level-shell)
l = subsidiary (orbitals)
ml = magnetic
K, L, M, N, O (1, 2, 3, etc.)
s, p, d, f (0, 1, 2, 3,…, n -1)
1, 3, 5, 7 (-l to +l)
ms = spin
½ , -½
18
Materials Science and Engineering
Electrons in Atoms
魏茂國
Table 2.1 The number of available electron states in some of the electron shells and subshells.
Principal
quantum
number
1
2
3
4
Shell
designation
K
L
M
N
n
Number of electrons
Subshells
Number
of states
Per subshell
Per shell
s
1
2
2
s
1
2
p
3
6
s
1
2
p
3
6
d
5
10
s
1
2
p
3
6
d
5
10
f
7
14
l = 0~n-1
ml = -l~l
ms = ½
8
18
32
2n2
19
Electron Energy States
Electrons...
• have discrete energy states
• tend to occupy lowest available energy state.
4d
4p
N-shell n = 4
3d
4s
Energy
3p
3s
M-shell n = 3
Adapted from Fig. 2.4,
Callister & Rethwisch 8e.
2p
2s
L-shell n = 2
1s
K-shell n = 1
20
fig_02_04
SURVEY OF ELEMENTS
• Most elements: Electron configuration not stable.
Element
Hydrogen
Helium
Lithium
Beryllium
Boron
Carbon
...
Neon
Sodium
Magnesium
Aluminum
...
Argon
...
Krypton
Atomic #
1
2
3
4
5
6
Electron configuration
1s 1
1s 2
(stable)
1s 2 2s 1
1s 2 2s 2
1s 2 2s 2 2p 1
1s 2 2s 2 2p 2
...
Adapted from Table 2.2,
Callister & Rethwisch 8e.
10
11
12
13
1s 2 2s 2 2p 6
(stable)
1s 2 2s 2 2p 6 3s 1
1s 2 2s 2 2p 6 3s 2
1s 2 2s 2 2p 6 3s 2 3p 1
...
18
...
36
1s 2 2s 2 2p 6 3s 2 3p 6
(stable)
...
1s 2 2s 2 2p 6 3s 2 3p 6 3d 10 4s 2 4p 6 (stable)
• Why? Valence (outer) shell usually not filled completely.
22
Electronic Configurations
ex: Fe - atomic # =
4d
4p
26 1s2 2s2 2p6 3s2 3p6
3d 6 4s2
N-shell n = 4
valence
electrons
3d
4s
Energy
3p
3s
M-shell n = 3
Adapted from Fig. 2.4,
Callister & Rethwisch 8e.
2p
2s
L-shell n = 2
1s
K-shell n = 1
23
TABLE OF CONTENT
•
•
•
•
•
•
•
•
•
•
Introduction
Fundamental concepts
Electrons in atoms
The periodic table
Bonding forces and energies
Primary interatomic bonds
Secondary bonding or van der Waals bonding
Mixed Bonding
Molecules
Bonding type-Material classification correlations
24
accept 2eaccept 1einert gases
give up 1egive up 2egive up 3e-
The
Periodic
Table
• Columns: Similar Valence Structure
H
He
Li Be
O
F
Na Mg
S
Cl Ar
Se
Br Kr
K Ca Sc
Rb Sr
Y
Cs Ba
Te
Po
I
Ne
Adapted from
Fig. 2.6,
Callister &
Rethwisch 8e.
Xe
At Rn
Fr Ra
Electropositive elements:
Readily give up electrons
to become + ions.
Electronegative elements:
Readily acquire electrons
to become - ions.
25
Materials Science and Engineering
Periodic Table
魏茂國
 Periodic table (週期表)
- All the elements have been classified according to electron configuration in the
periodic table. Here, the elements are situated, with increasing atomic number, in
seven horizontal rows called periods (週期). All elements arrayed in a given
column or group (族) have similar valence electron structures, as well as chemical
and physical properties.
- The elements positioned in Group 0 are the inert gases, which have filled electron
shells and stable electron configurations.
- The Group VII elements are termed the halogen.
- The alkali and the alkaline earth metals are labeled as Groups IA and IIA,
respectively.
- Group IIIB through IIB are termed the transition metals, which have partially
filled d electron states.
26
Relative sizes of some atoms and ions (in nm)
Animation: Atomic radii
27
Electron Structure and Chemical Activity (Cont..)
•
Electronegative elements accept electrons during chemical
reaction.
• Some elements behave as both electronegative and
electropositive.
• Electronegativity is the degree to which 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.
Te
Na
Electropositive 0
K 1
W
2H
28
N
Se
3
O
Fl
4
Electronegative
Electronegativity
• Ranges from 0.7 to 4.0,
• Large values: tendency to acquire electrons.
Smaller electronegativity
Larger electronegativity
Adapted from Fig. 2.7, Callister & Rethwisch 8e. (Fig. 2.7 is adapted from Linus Pauling, The Nature of the Chemical
Bond, 3rd edition, Copyright 1939 and 1940, 3rd edition. Copyright 1960 by Cornell University.
29
Oxidation numbers of the elements
30
TABLE OF CONTENT
•
•
•
•
•
•
•
•
•
•
Introduction
Fundamental concepts
Electrons in atoms
The periodic table
Bonding forces and energies
Primary interatomic bonds
Secondary bonding or van der Waals bonding
Mixed Bonding
Molecules
Bonding type-Material classification correlations
31
Materials Science and Engineering
Bonding Forces and Energies
魏茂國
 Interaction between two isolated atoms
- Two types of force, attractive and repulsive, and the magnitude of each
depends on the separation or interatomic distance.
- The origin of an attractive force FA depends on the particular type of bonding
that exists between the two atoms.
- Repulsive force FR arise from interactions between the negatively charged
electron clouds for the two atoms and are important only at small values of r as
the outer electron shells of the two atoms begin to overlap.
- The net force FN between the two atoms is just the sum of both attractive and
repulsive components.
FN  FA  FR
- At the equilibrium state, the centers of the two atoms will remain separated by
the equilibrium spacing r0, as indicated in Fig. 2.8a.
- For many atoms, r0 is approximately 0.3 nm.
32
Materials Science and Engineering
Bonding Forces and Energies
魏茂國
- Mathematically, energy (E) and
force (F) are related as
E   Fdr
For atomic systems,
r
E N   FN dr

r
r


 E N   FAdr   FR dr
 EN  E A  ER
EN: net energy, EA: attractive
energy, ER: repulsive energy.
Figure 2.8 (a) The dependence of
repulsive, attractive, and net forces
on interatomic separation for two
isolated atoms. (b) The dependence
of repulsive, attractive, and net
potential energies on interatomic
separation for two isolated atoms.
bonding
energy
33
Atomic and Molecular Bonds
• Ionic bonds :- Strong atomic bonds due to transfer of
electrons
• Covalent bonds :- Large interactive force due to sharing of
electrons
• Metallic bonds :- Non-directional bonds formed by sharing
of electrons
• Permanent Dipole bonds :- Weak intermolecular bonds due
to attraction between the ends of permanent dipoles.
• Fluctuating Dipole bonds :- Very weak electric dipole bonds
due to asymmetric distribution of electron densities.
Animation: Ionic vs. covalent bond
34
Ionic Bonding
•
•
•
•
Occurs between + and - ions.
Requires electron transfer.
Large difference in electronegativity required.
Example: NaCl
Na (metal)
unstable
Cl (nonmetal)
unstable
electron
Na (cation)
stable
-
+
Coulombic
Attraction
Cl (anion)
stable
35
Ionic Bonding - Example
• Ionic bonding in NaCl
3s1
3p6
Sodium
Atom
Na
Sodium Ion
Na+
I
O
N
I
C
B
O
N
D
36
Chlorine
Atom
Cl
Chlorine Ion
Cl -
Figure 2.10
Ionic Bonding
• Energy – minimum energy most stable
– Energy balance of attractive and repulsive terms
EN = EA + ER =
-
A
r

B
rn
Repulsive energy ER
Interatomic separation r
Net energy EN
Adapted from Fig. 2.8(b),
Callister & Rethwisch 8e.
Attractive energy EA
37
Examples: Ionic Bonding
• Predominant bonding in Ceramics
NaCl
MgO
CaF 2
CsCl
Give up electrons
Acquire electrons
Adapted from Fig. 2.7, Callister & Rethwisch 8e. (Fig. 2.7 is adapted from Linus Pauling, The Nature of the Chemical
Bond, 3rd edition, Copyright 1939 and 1940, 3rd edition. Copyright 1960 by Cornell University.
38
Covalent Bonding
• similar electronegativity  share electrons
• bonds determined by valence – s & p orbitals dominate
bonding
• Example: CH4
C: has 4 valence e-,
needs 4 more
H: has 1 valence e-,
needs 1 more
Electronegativities
are comparable.
H
CH 4
H
C
H
shared electrons
from carbon atom
H
shared electrons
from hydrogen
atoms
Adapted from Fig. 2.10, Callister & Rethwisch 8e.
39
Covalent Bonding - Examples
•
In case of F2, O2 and N2, covalent bonding is formed by sharing p
electrons
•
Fluorine gas (Outer orbital – 2s2 2p5) share one p electron to attain noble gas
configuration.
F + F
•
H
F F
F F
Bond Energy=160KJ/mol
Oxygen (Outer orbital - 2s2 2p4) atoms share two p electrons
O + O
O
O
O=O
Bond Energy=220KJ/mol
•
HH
Nitrogen (Outer orbital - 2s2 2p3) atoms share three p electrons
N N
N + N
40
N
N
Bond Energy=945KJ/mol
Bond Hybrization
• Carbon can form sp3 hybrid
orbitals
Fig. 2.14, Callister & Rethwisch 9e.
(Adapted from J.E. Brady and F. Senese, Chemistry:
Matter and Its Changes, 4th edition. Reprinted with
permission of John Wiley and Sons, Inc.)
Fig. 2.13, Callister & Rethwisch 9e.
41
Covalent Bonding: Carbon sp3
• Example: CH4
C: has 4 valence e-,
needs 4 more
H: has 1 valence e-,
needs 1 more
Electronegativities of C and H
are comparable so electrons
are shared in covalent bonds.
Fig. 2.15, Callister & Rethwisch 9e.
(Adapted from J.E. Brady and F. Senese, Chemistry:
Matter and Its Changes, 4th edition. Reprinted with
permission of John Wiley and Sons, Inc.)
42
Covalent Bonding in Carbon
•
Carbon has electronic configuration 1s2 2s2 2p2
Ground State arrangement
1s
2s
2p
Two ½ filed 2p orbitals
Indicates
carbon
Forms two
Covalent
bonds
• Hybridization causes one of the 2s orbitals promoted to 2p
orbital. Result
four sp3 orbitals.
Indicates
four covalent
bonds are
1s
2p
formed
Four ½ filled sp3 orbitals
43
Structure of Diamond
• Four sp3 orbitals are directed symmetrically toward corners
of regular tetrahedron.
• This structure gives high hardness, high bonding strength
(711KJ/mol) and high melting temperature (3550oC).
Carbon Atom
Tetrahedral arrangement in diamond
Figure 2.18
Figure 2.19
44
Metallic Bonding
• Atoms in metals are closely packed in crystal structure.
• Loosely bounded valence electrons are attracted towards
nucleus of other atoms.
• Electrons spread out among atoms forming electron clouds.
Positive Ion
• These free electrons are
reason for electric
conductivity and ductility
• Since outer electrons are
shared by many atoms,
metallic bonds are
Non-directional
Valence
electron charge cloud
45
Figure 2.24
Metallic Bonds (Cont..)
• Overall energy of individual atoms are lowered by metallic
bonds
• Minimum energy between atoms exist at equilibrium
distance a0
• Fewer the number of valence electrons involved, more
metallic the bond is.
 Example:- Na
Bonding energy 108KJ/mol,
Melting temperature 97.7oC
• Higher the number of valence electrons involved, higher is
the bonding energy.
 Example:- Ca
Bonding energy 177KJ/mol,
Melting temperature 851oC
46
TABLE OF CONTENT
•
•
•
•
•
•
•
•
•
•
Introduction
Fundamental concepts
Electrons in atoms
The periodic table
Bonding forces and energies
Primary interatomic bonds
Secondary bonding or van der Waals bonding
Mixed Bonding
Molecules
Bonding type-Material classification correlations
47
Secondary Bonding: van der Waals bonds
• Secondary bonds are due to attractions of electric dipoles in
atoms or molecules.
• Dipoles are created when positive and negative charge
centers exist.
Dipole moment=μ =q.d
-q
+q
Figure 2.26
q= Electric charge
d = separation distance
d
• Two types of bonds: permanent and fluctuating
• Three types of bonds: London, Debye and Keesom
48
SECONDARY BONDING
Arises from interaction between dipoles
• Fluctuating dipoles
ex: liquid H 2
H2
H2
asymmetric electron
clouds
+
-
+
secondary
bonding
-
H H
H H
secondary
bonding
Adapted from Fig. 2.13,
Callister & Rethwisch 8e.
• Permanent dipoles-molecule induced
-general case:
-ex: liquid HCl
-ex: polymer
+
-
H Cl
secondary
bonding
+
secondary
bonding
H Cl
Adapted from Fig. 2.15,
Callister & Rethwisch 8e.
secondary bonding
49
Fluctuating Dipoles
• Weak secondary bonds in noble gasses.
• Dipoles are created due to asymmetrical distribution of
electron charges.
• Electron cloud charge changes with time.
Symmetrical
distribution
of electron charge
Figure 2.27 Asymmetrical
Distribution
(Changes with time)
50
Hydrogen Bonds (Keesom force)
• Hydrogen bonds are Dipole-Dipole interaction
between polar bonds containing hydrogen atom.
 Example : In water, dipole is created due to asymmetrical
arrangement of hydrogen atoms.
 Attraction between negative oxygen pole and positive
hydrogen pole.
H
O
105 0
Figure 2.28
H
51
Hydrogen
Bond
TABLE OF CONTENT
•
•
•
•
•
•
•
•
•
•
Introduction
Fundamental concepts
Electrons in atoms
The periodic table
Bonding forces and energies
Primary interatomic bonds
Secondary bonding or van der Waals bonding
Mixed Bonding
Molecules
Bonding type-Material classification correlations
52
Mixed Bonding
• Ionic-Covalent
– % of ionic character = (1 – e
(-1/4)(Xa-Xb)2
) (100%)
• Metallic-Covalent
– dsp orbital bonding,
– covalent-metallitivities:
– % of covalent bonding = C-M / 4 (100%)
• Metallic-Ionic
– Intermetallic compound
• Ionic-Covalent-Metallic
• Primary + Secondary
53
Fig_2-25
Copyright © 2014 John Wiley & Sons, Inc. All rights reserved.
Solid water
Liquid water
Fig_2-24
Copyright © 2014 John Wiley & Sons, Inc. All rights reserved.
 Molecules
- Many of the common molecules are composed of groups of
atoms that are bound together by strong covalent bonds. (F2,
O2, N2, C2H6, C6H6, …)
- In the condensed liquid and solid states, bonds between
molecules are weak secondary ones. Consequently, molecular
materials have relatively low melting and boiling temperatures
- Polymers of large molecules are mostly solid – their
properties affected by van der Waals and hydrogen bonds
56
Summary: Bonding
Comments
Type
Bond Energy
Ionic
Large!
Nondirectional (ceramics)
Covalent
Variable
large-Diamond
small-Bismuth
Directional
(semiconductors, ceramics
polymer chains)
Metallic
Variable
large-Tungsten
small-Mercury
Nondirectional (metals)
Secondary
smallest
Directional
inter-chain (polymer)
inter-molecular
57
Properties From Bonding: Tm
• Bond length, r
• Melting Temperature, Tm
Energy
r
• Bond energy, Eo
ro
Energy
r
smaller Tm
unstretched length
ro
r
Eo =
“bond energy”
larger Tm
Tm is larger if Eo is larger.
58
Properties From Bonding : a
• Coefficient of thermal expansion, a
length, L o
coeff. thermal expansion
unheated, T 1
DL
DL
= a (T2 -T1 )
Lo
heated, T 2
• a ~ symmetric at ro
Energy
unstretched length
ro
Eo
Eo
r
a is larger if Eo is smaller.
larger a
smaller a
59
Summary: Primary Bonds
Ceramics
(Ionic & covalent bonding):
Metals
(Metallic bonding):
Polymers
(Covalent & Secondary):
Large bond energy
large Tm
large E
small a
Variable bond energy
moderate Tm
moderate E
moderate a
Directional Properties
Secondary bonding dominates
small Tm
small E
large a
60