Chapter 2. Atomic structure and
interatomic bonding
2.1. Atomic structure
2.1.1.Fundamental concepts
2.1.2. Electrons in atoms
2.1.3. The periodic table
2.2. Atomic bonding in solids
2.2.1. Bonding forces and energies
2.2.2. Primary interatomic bonds
2.2.3. Secondary bonding
2.2.4. Molecules
2.1. Atomic structure
2.1.1.Fundamental concepts
Atom consists of a nucleus (protons and neutrons) and electrons
that move in their orbits.
e : 1.6 X 10-19 C
Mp = mn = 1.67 X 10-27 kg
Me = 9.11 X 10-31 kg
A chemical element has
Atomic number (Z) : the number of proton in the nucleus,
Atomic mass (A) : the sum of the masses of protons and neutrons
in the nucleus,
Neutron number (N) : the number of neutron may vary for a given
elements.
Isotopes : elements with two or more different atomic masses
Atomic weight (amu)
A≈Z+N
2.1.2. Electrons in atoms
Bohr atomic model
An electron is an particle
Electrons revolve around
an atomic nucleus in
discrete orbitals
Wave-mechanical model
Electrons are considered as both wave-like and particle-like
Every electron in an atom is characterized by four parameters
called quantum numbers.
Principal quantum number, n
n = 1 → K; n = 2 → L; and so on
Second quantum number, l
s, p, d, f
Third quantum number, ml (number of energy states)
s → ml = 1; p → ml = 3 ; d → ml = 5 ; f → ml = 7
Fourth quantum number, ms (spin orientation when a magnetic
field is applied)
ms = + ½ and – ½
The schematic of the relative energies of the electrons
for various shells and subshells
Comparison between the two models
a. Bohr atomic model
Electrons are in fixed positions
and energy (quantized energy
levels).
b. Wave-mechanical model
Electrons’ position is considered
to be the probability of an
electron’s being at various
locations around the nucleus.
Electron configurations
Pauli exclusion principle
Each electron state can hold no more than two electrons,
which must have opposite spins
Ground state is a level where electrons occupy the lowest
possible energy according to Pauli exclusion principle
Electron configurations represent the manner in which these
states are occupied.
Examples:
Hydrogen (H) 1s1
Helium (He)
1s2
Sodium (Na)
1s2 2s2 2p6 3s1
Calcium (Ca) 1s2 2s2 2p6 3s2 3p6 4s2
2.1.3. The periodic table
The periodic table consists of all elements that have been
classified according to electron configuration.
Groups IA and IIA (alkali and alkaline earth metals)
Elements with one and two electrons in excess of stable structures.
Groups IIIA, IVA, and VA
Elements between metals and nonmetals by virtue of their valence
electron structures.
Groups VIA and VIIA
Elements with two and one electron in deficient of stable structures.
(Group VIIA is called the halogens)
Group VIIIA (Inert gases)
Elements which have filled electron shells and stable configurations.
2.2. Atomic bonding in solids
2.2.1. Bonding forces and energies
Interatomic forces determine the physical properties of
materials.
There are two types of forces:
the attractive and the repulsive forces
Most of time, it is more convenient to work with potential
energies, which are attractive and the repulsive energies.
Both terms depend on the distance between the centre of
two atoms.
The equilibrium distance, ro
For many atoms, ro = 0.3 nm (3Å )
Bonding force
The net force, FN
FN = FA + FR
In equilibrium:
FA + FR = 0
Bonding energy
The potential energy
between two atoms
E = ∫ Fdr
The net energy, EN
EN = EA + ER
Types of bonding:
A. Primary bonding or chemical bonding
This bonding is found in solids and involves the valence electrons.
This type of bonding is strong (» 100 kJ/mol)
Examples: ionic, covalent, and metallic bonds
B. Secondary bonding or physical bonding or van der Waals
This bonding is found in most solids and arises from atomic or
molecular dipoles.
This type of bonding is weak ( ≅ 10kJ/mol)
Examples: fluctuating induced dipole bonds, polar moleculeInduced dipole bonds, and pemanent dipole bonds
A. Primary bonding or chemical bonding
Ionic bonding
It is always found in compounds that are composed of both metallic
and nonmetallic elements. Atoms of a metallic element easily give
up their valence electrons to the nonmetallic atoms.
This bonding is a nondirectional bonding, the magnitude of the bond
is equal in all directions around an ion.
Coulombic bonding force
Attractive energy:
E
A
=−
r
A
Repulsive energy: E =
B
B
A, B, n = constants,
n~8
r
n
A. Primary bonding or chemical bonding
Covalent bonding
It is usually found in many nonmetallic elemental molecules (H2, Cl2, F2)
and molecules containing dissimilar atoms (CH4, H20, HNO3, HF)
This bonding is formed on stable electron configurations by sharing of
electrons between adjacent atoms.
A very strong covalent bond
Diamond with a very high
melting temperature
(713 kJ/mol; 3550 ºC)
A very weak covalent bond
Bismuth with a very low
melting temperature
(270 ºC)
A. Primary bonding or chemical bonding
Metallic bonding
It is found in many metals and their alloys (group IA and IIA).
Metallic materials have 1, 2 or at most 3 valence electrons.
These valence electrons are not bound to any particular atom to any
Particular atom in the solid and
are free to drift throughout the
entire metal.
“sea of electrons”
or “electron cloud”
Net negative charge
Ion cores
Net positive charge
Weak metallic bond
Hg (68 kJ/mol; -39 ºC)
Strong metallic bond
W (850 kJ/mol; 3410 ºC)
B. Secondary bonding or physical bonding or van der Waals
Fluctuating induced dipole bonds
All atoms have constant vibrational motion and it causes electrical
symmetry and creates small electric dipoles
B. Secondary bonding or physical bonding or van der Waals
Polar molecule-induced dipole bonds
It causes by virtue of an asymmetrical arrangement of positively
and negatively charged regions
B. Secondary bonding or physical bonding or van der Waals
Permanent dipole bonds
It exist between adjacent polar molecules. The hydrogen bond is
the strongest secondary bonding type.