‘’ATOMIC STRUCTURE AND
BONDING’’
IE-114 Materials Science and General Chemistry
Lecture-2
Outline
• Atomic Structure
(Fundamental concepts, Atomic models (Bohr and Wave-Mechanical Atomic Model),
Electron configurations)
• Periodic Table
(Classification of elements, their characteristics)
• Atomic Bonding in Solid Materials
(Primary Bonding:Ionic, covalent, metallic bonds)
(Secondary bonding: Fluactuating Induced Dipole, Polar Molecule-Induced
Dipole, Permanent Dipole Bonds)
Why atomic structure is important?
 Some properties of solid materials depend on atomic nature
and its arrangement.
 Type of atomic arrangements : crytalline or amorphous
 Type of bonding (interactions among the atoms) determines the melting temperature (Tm),
coefficient of thermal expansion(α), mechanical properties, i.e. Elastic modulus, E
Graphite and Diamond
graphite
Graphite: Soft and greasy feel
Diamond: Hardest material known
diamond
different type of interatomic
bonding in graphite and diamond.
Atomic Structure
Atom = nucleus (protons+neutron) and electrons
Electrons are negatively(-) charged,
Protons are positively(+) charged
Neutrons are electrically neutral “particle
Charge of an electron = - 1.60x10-19 C
Charge of a proton = +1.60x10-19 C
 Atomic Number (Z): Number of protons in the nucleus
 Electrically neutral atom; # protons = #electrons
Mass of an Atom
 Atomic mass (A) for an atom = masses of protons(Z)+masses of neutrons(
(electrons are not considered, because ...?)
Mass of proton = mass of neutrons=1.67x10-27 kg
Mass of electron= 9.11x10-31 kg
 Atoms with two or more atomic masses (ISOTOPES)
For all atoms of an element the number of protons are the same, but the number
of neutrons may vary, which vary the atomic mass,
Example: 12C, 13C, 14C
 Atomic Weight:
Weighted average of the atomic masses of the atom’s naturally
occuring isotopes
The atomic mass unit (amu)
It is used for the computation of atomic weight.
Scale: 1 amu= 1/12 of the atomic mass of the Carbon (C)
(A=12,00000 for carbon 12 isotope)
1 amu/atom= 1 g/mol ( 1 mol of a substance=6.023x1023
atoms)
For example: Fe
A=55.85 amu/atom or 55.85 g/mol (this is most commonly used
form)
Structure of Atom
1) Bohr Atomic Model
2) Wave-Mechanical Model
 BOHR ATOMIC MODEL:
(Used hydrogen atom)
1) Electrons are assumed to be positioned around the nucleus in discrete
orbitals
2) Position of the electron is more or less well defined in its orbital.
Nucleus: Z = # protons
N = # neutrons
Energy of Electrons
 Bohr Atomic model describe the electrons in terms of their
positions (orbitals) and energy (quantized energy levels by Rydberg
equation).
E= - (2π2me4/n2h2) = - (13.6/n2) eV
e: electron charge
m: electron mass
n: principal quantum number or principal energy
levels(1,2,3,….)
h:Planck’s constant
An electron can change its energy level
To a higher level by absorbing energy, to a lower level by emitting
energy
Example:
n=1  E1 = -13.6 eV
n=2  E2 = - 3.4 eV
E2>E1
If electron changes its energy level from 1 to 2 (from lower to higher energy level)
It must absorb energy, and the amount of energy absorbed;
E =E2 - E1 = -3.4 – (-13.6) = 10.2 eV
Allowed energy levels for hydrogen electron in Bohr Model
Ionization energy:
Energy required to remove the electron completely from the atom
Ionization energy for hydrogen electron is 13,6 eV
Figure.The first three electron energy states for the Bohr hydrogen aton
* Bohr’s model was not able to explain quantitatively the spectra of the atoms more complex
than hydrogen and the model could not have been modified.
 WAVE-MECHANICAL MODEL:
 Limitations of Bohr model was resolved by this model and electrons are
considered to behave both wave-like and particle-like.
Electrons are no longer treated as a particle moving in discrite orbitals.
Position of electron is described by a probability distribution or electron cloud
Heisenberg’s uncertainty principle;
Position and momentum of a small particle such as
an electron can not be determined simultaneously.
BOHR MODEL
WAVE-MECHANICAL
MODEL
Since the position of an electron can not be precisely
determined, an electron charge cloud density distribution is
used
Motion of electron around its nucleus and its energy is
characterized by 4 QUANTUM NUMBERS (n, l, ml, ms)
1) Principal quantum number,n:
Represents main energy levels for the electrons or shells (n=1 to 7 )
n=1 (first shell, K) n= 2 second shell(L), so forth... or
2) Secondary quantum number, l:
Specifies subenergy levels within the main energy levels(subshells)
and related to shape of the electron subshell
l=0..n-1
Number designation of l: 0 1 2 3 4 5....
Letter designation of l : s p d f g h
Principle quantum
number, n
Shell Designation
Subshells
1
K
s
2
L
s,p
3
M
s,p,d
4
N
s,p,d,f
3) Magnetic quantum number, ml:
Number of orbitals or energy states for each subshell
ml=2l+1
Example: For a given l, ml can range from +l to –l
l=0 (s subshell)  ml = 1 energy state (0)
l=1 (p subshell)  ml = 3 energy states (+1, 0, -1)
l=2 (d subshell)  ml = 5 energy states (+2,+1, 0, -1,-2)
Pauli’s Exclusion Principle:
No two electrons can have the identical values for all four of their quantum numbers
4) Electron spin quantum number, ms:
Specifies two allowed spin directions for an electron.
ms = +1/2 and -1/2
Orbitals
Orbital Shapes
Electron density in s and p- orbitals
Comparison of electron energy states in
Bohr and Wave-mechanical Models
BOHR MODEL
WAVE-MECHANICAL MODEL
The maximum number of electrons in each shell in an atom is 2n2
***Electrons fill up the lowest
possible energy states in the
electron shells and subshells
Electron Configurations
 Represents the manner in which the states are occupied
 The number of electrons in each subshell is indicated by a superscript
after the shell-subshell designation.
Example:
H 1s1
He 1s2
Na 1s22s22p63s1
Na 1s22s22p63s1
p subshell has three orbitals:
Principal quantum number,n
ml=-1,0,+1, each of these orbitals
contains 2 electrons
(SHELL K,L,M,..)
Secondary quantum number,l
(SUBSHELL; s,p,d,f)
s subshell has one orbital: ml=0
Orbital contains 2 electrons
Valance Electrons:The electrons occupying the outermost shell
These electrons participate in bonding. Many of the physical and chemical
properties of solids are based on these valence electrons.
Stable electron configurations
• have complete s and p subshells
• tend to be unreactive. (inert, or noble, gases)
Survey of Elements
* Most of the elements are not stable.
Electron configuration
1s1
1s2
(stable)
1s22s1
1s22s2
1s22s22p1
1s22s22p2
...
1s22s22p6
(stable)
1s22s22p63s1
1s22s22p63s2
1s22s22p63s23p1
...
1s22s22p63s23p6
(stable)
...
1s22s22p63s23p63d10 4s246
(stable)
Periodic Table
 Elements are classified according to electron configuration in this
table.
 Same column, or group, have similar chemical and physical properties
due to similar valence electron configurations
 These properties change gradually and systematically across each
period moving horizontally.
Atomic Number(Z)
+
 Groups are designated at the top by the numbers 0-7 and by the letters
A and B.
A group elements- Representative or main group elements
B group elements- Transition elements
*Group 0: inert gases (filled electron shells)
*Group IA and IIA are alkali (except H) and alkaline earth
metals
*Group IIIA, IVA and VA elements have characteristics between
metal and nonmetals because of their valence electron
configurations.
*Group VIIA (halogens) and VIA elements= one and two
electrons deficient respectively from having stable
configurations.
*Groups from IIIB to IIB are transition metals, with partially
filled d electron states and in some cases one or two electrons in
the next higher shell.
Electropositive and Electronegative Elements
Electropositive elements: Elements capable of giving up their
electrons to become positively charged ions (located on the left of the
table.)
Electronegative elements: Elements ready to accept electrons to form
negatively charged ions or to share their electrons.
Electropositive elements:
Readily give up electrons
to become + ions.
Electronegative elements:
Readily acquire electrons
to become - ions.
Electronegativity
 The degree to which an atom attracts electrons to itself
• Ranges from 0.7 to 4.0
• Large values: tendency to acquire electrons.
Smaller electronegativity
Larger electronegativity
Atomic Bonding in Solids
 Atomic bonding can be
explained by interaction of two
isolated atoms
Net force is zero (equilibrium state)
r0=equilibrium spacing
 Physical properties are
related to interatomic forces
that bind the atoms together
Bonding energy, Eo
Energy required to seperate these two atoms to an infinite seperation
Bonding energy (Eo); Solids >Liquids > Gases
1) Bonding arises from the tendency of the atoms to assume stable
electron structures
2) Valance electrons are involved
3) Nature of bond depends on the electron structure
Properties from bonding
Bond energy, Eo
1) Melting Temperature, Tm
Tm is larger if Eo is larger.
2) Coefficient of thermal expansion,
α
coeff. thermal expansion
L
= a(T2-T1)
Lo
• α ~ symmetry at ro
a is larger if Eo is smaller.
3) 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
larger Elastic Modulus
Types of Bondings
1)Primary Bonding:
 Ionic bonding
 Covalent bonding
 Metallic bonding
2)Secondary bonding:
 Fluactuating Induced Dipole
 Polar Molecule-Induced Dipole
 Permanent Dipole Bonds
Ionic Bonding
 Found in compounds formed by metallic and nonmetallic elements
(occurs between + and – ions)
 Requires electron transfer.
Large difference in electronegativity is required.
• Example: NaCl
Enet= Eatt. + Erep.
Enet= - (A/r) + B/rn
B, and n are constants. n is
approximately 8.
(A = (Z1Z2e2/4πo) + B/rn)
The ionic bonding is nondirectional, that is the magnitude of the bond is
equal in all directions.
The predominant bonding in ceramics is ionic.
Give up electrons
Acquire electrons
Covalent Bonding
 Stable electron configurations are assumed by sharing of
electrons between adjacent atoms.
• Example: CH4(methane)
H feels like helium electron configuration, while C feels like neon electron configuration.
C: has 4 valence e, needs 4
more
H: has 1 valence e, needs 1 more
Electronegativities are
comparable.
 Covalent bonding is directional
 It forms between two specific atoms and may exist only in the direction
between one atom and another.
• Molecules with nonmetals
• Molecules with metals and nonmetals
• Elemental solids (RHS of Periodic Table)
• Compound solids (about column IVA)
Number of covalent bonds
Number of covalent bonds= 8-N’
N’= number of valence electrons
Example: Cl atom
7 valence electrons, an atom can have maximum 1 more bond (completing the valence
orbital electron number to eight)
 Covalent bonds may be extremely strong (like in diamonds) or may be
weak (like in Bismuth).Polymeric materials are covalently bonded
materials.
%Ionic character
 Some bonds are partially ionic and partially covalent. The degree of
either bond is controlled by the electronegativities of the composing
atoms.
%ionic character = (1-e-(0.25)(X -X )2)x100
A
B
(XA and XB are the electronegativities of the respective elements)
As the electronegativity difference gets higher, the bonding becomes
more ionic.
Metallic Bonding
 Valence electrons are not bound to any particular atom in the solid
and they are more or less free to move throughout the entire metal.
 Primary bond for metals and their alloys
 Metallic bond is nondirectional.
 Metalling bonding explains the heat and electric conductivity of the
metallic materials as well as their ductility.
Secondary Bonding
 Arises from interaction between dipoles
• Fluctuating dipoles
• Permanent dipoles-molecule induced
-general case:
-ex: liquid HCl
-ex: polymer