Chapter 2: Atomic bonding
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Chapter 2: Atomic bonding
Reading assignment
Ch. 2 and 3 of textbook
Homework No. 1
Problems 2-8, 2-9, 2-13.
Homework No. 2
Problems 3-13, 3-15, 3-17, 3-19, 3-20, 3-21, 3-27.
Effect of atomic bonding
Example: carbon exists as graphite (soft with greasy feeling) or diamond (hardest known material)
Atomic & electronic
configuration
Bonding
BOND STRENGTH
Mechanical & Physical
Properties
graphite diamond
Primary and secondary bonding primary bonds: strong atom-to-atom attractions produced by changes in electron position of the valence e– . Example : covalent atom between two hydrogen atoms
secondary bonds: much weaker. It is the attraction due to overall “electric fields”, often resulting from electron transfer in primary bonds. Example: intramolecular bond between H2 molecules gas
Electronic configurations
Valence electrons They represent the ability of an element to
enter into chemical combination with others. Valence es
− participate in the bonding between atoms.
Valence = # of electrons in outermost combined sp level.
Examples of the valence are:Mg: 1s22s22p63s2 valence = 2Al: 1s22s22p63s23p1 valence = 3 Ge: 1s22s22p63s23p63d104s24p2 valence = 4
Primary bonding types
Ionic bondingCovalent bondingMetallic bonding
Ionic bonding
Ionic bonding in NaCl
3s1
3p6
SodiumAtom
Na
ChlorineAtom
Cl
Sodium IonNa+
Chlorine IonCl -
IONIC
BOND
2-15
Ionic bondingIonic Bond:
The attractive bonding forces are coulombic (different polarities):
Interionic force
Force of attraction between Na+ and Cl- ions
Z1 = +1 for Na+, Z2 = -1 for Cl-
e = 1.60 x 10-19 C , ε0 = 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
NC
aeZZF attraction
910-212-
219
2
0
2
21 1002.3m) 10 x /Nm2)(2.76C 10 x 8.85(4
)1060.1)(1)(1(
4
+
2-18
Interatomic spacing The equilibrium distance between atoms is
caused by a balance between repulsive and attractive forces.
Equilibrium separation occurs when no net force acts to either attract or separate the atoms or the total energy of the pair of atoms is at a minimum.
For a solid metal the interatomic spacing is equal to the atomic diameter or 2r.
For ionically bonded materials, the spacing is the sum of the two different ionic radii.
Bond energyBonding force
and Energy curves for a Na+
& Cl- pair.
Since F = dE/da, the equilibrium bond length (ao) occurs @ F=0 and E is a minimum.
Coordination Coordination number (C.N.) = No. of nearest neighbors (radius R)
around (touching) a particular atom/ion (radius r).
C.N. depends on r/R ratio.
C.N. = 4 (in 2 dimensions)
Stable Stable Unstable(critical case)
r/R > min r/R = min r/R < min
C.N. = 2
Schematic drawing with nearest neighbors not in contact
C.N. = 3
Schematic drawing with nearest neighbors not in contact
C.N.= 43 dimensions
C.N. = 4 (3 dimensions)
C.N. = 8 (3 dimensions)
C.N. = 12 (3 dimensions)
C.N. = 63 dimensions
C.N. = 6 (3 dimensions)
C.N. = 4 (in 2 dimensions)C.N. = 6 (in 3 dimensions)
Stable Stable Unstable(critical case)
r/R > min r/R = min r/R < min
C.N. = 3
C.N. = 8 (3 dimensions)
C.N. = 4
Criteria of packing ions in a solid
Positive charge = negative charge Nearest neighbors of a cation are anions; nearest
neighbors of an anion are cations. (Nearest neighbors touch one another.)
The coordination number (CN) is determined by r/R, where r = radius of smaller ion (usually the cation), and R=radius of larger ion (usually the anion). The greater is r/R, the higher is CN.
The largest allowable CN is most favorable.
Summary on ionic bonding
The attractive force (energy) for two isolated ions is a function of distance.
Bonding is nondirectional.This is the predominant bonding type
in ceramics. 600-1500 kJ/mol (3-8 eV/atom)
bonding energies are large high Tm.
Covalent bonding
Directional bond due to the sharing of electrons
between atoms
Cl2 molecule
Planetary model
Actual electron density
Electron dot schematic
Bond line schematic
Example 1. Br2 (a bromine molecule)
Br has an outermost electronic configuration of 4s24p5, i.e.,
A single bond
A -bondEnd-to-end overlap
Example 2. O2 (an oxygen molecule)
O has an outermost electron configuration of 2s22p4, i.e.,
A double bond
A σ-bond (end-to-end overlap) together
with a π-bond (side-to-side overlap).
Double bond
Single bonds
Every carbon atom along the chain is four-fold coordinated.
The electronic configuration of carbon is 1s22s22p2, i.e.
An excited state of carbon with electronic configuration
sp3 hybridization
Mixing of an s electron cloud with three p
electron clouds
Methane molecule
Methane molecule
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
2-25
SiO4
tetrahedron in silicate glass
Silicon dioxide (SiO2)
Covalent network solids
Network of covalent bondsExamples: diamond, silicon,
etc.
Properties of covalent network solids
Materials have poor ductility. Poor electrical conductivity. Many ceramic, semiconductor and polymer materials
are fully or partly covalent.
Mixed bonding (ionic + covalent)
Few compounds exhibit pure ionic or pure
covalent bonding. the bond type degree
depends on their position in the Periodic Table. The greater the difference in
electronegativity, the more ionic is the bond.
Conversely, the smaller the difference, the larger is the degree of covalency.
Example of mixed ionic-covalent bonding: HF (a hydrogen fluoride molecule)
The electronic configuration of H is 1s1, i.e.,
There is one unpaired electron.
The electronic configuration of F is 1s22s22p5, i.e.
There is also one unpaired electron.
Metallic Bonding
Found in metallic elements (low electronegativities).
Give up their valence electrons to form a “sea or cloud” of electrons.
The valence electrons move freely within the electron sea and become associated with the ion cores. The free electrons shield the (+) charged ion cores from repulsion.
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. These free electrons are reason for electric conductivity and ductility Since outer electrons are shared by many atoms, metallic bonds are Non-directional
Positive Ion
Valence electron charge cloud2-28
Metallic Bonding
Good thermal &
electrical
conductors
The free electrons in the “cloud” move freely under an applied voltage.
Bond energy
It is the energy required to create or break the bond.
Materials with high bond energy high strength and high melting point.
Ionic materials have a large bond energy due to the large difference in electronegativities between the ions.
Metals have lower bond energies, because the electronegativities of the atoms are similar.
Bond energy and melting temperature
Secondary bonding
Van der Waals bonding
Secondary bonding exists between virtually all Secondary bonding exists between virtually all molecules, but its presence is diminished if any molecules, but its presence is diminished if any primary bond is present. primary bond is present.
Dipoles are created when positive andnegative charge centers exist.
-q
Dipole moment=μ =q.d
q= Electric charged = separation distance
2-30
+q d
Skewed electron cloud
Secondary bonds are due to attractions of electric dipoles in atoms or molecules.
Electric dipole types
Permanent dipolesInduced dipoles
Permanent dipoles
Dipoles that do not fluctuate
with time
Hydrogen fluoride HF
Example of a permanent dipole
Permanent dipoles in water
Attraction between positive oxygen pole and negative hydrogen pole.
Water
105 0O
H
HDipole-dipole
interaction
2-33
Dipole-dipole interaction in water
Methane
Vector sum of four C-H dipoles is zero.
Symmetricalarrangement
of 4 C-H bondsCH4
No dipolemoment
CH3ClAsymmetrical
tetrahedralarrangement
Createsdipole
2-32
Methane
Methyl chloride
Cl- ions in green
Hydrogen bonding
Hydrogen bonds are dipole-dipole interaction between polar bonds
containing hydrogen atoms. It is a particularly strong type of secondary
bonding, due to the almost bare proton. Examples: water, HF, etc.
Induced dipolesNo permanent dipole momentStatistical fluctuation in electron
density distributionLondon dispersion forces (weak)Examples: argon (an inert gas),
methane (CH4 - a symmetric molecule)
Weak secondary bonds in noble gasses. Dipoles are created due to asymmetrical distribution of electron
charges. Electron cloud charge changes with time.
Symmetricaldistribution
of electron charge
Asymmetricaldistribution
(Changes with time)
2-31
London forces between methane molecules
Another example of mixed bonding types in a material
Graphite
Crystal forms of carbon
GraphiteDiamond Fullerene
Fullerene (a molecule)
Diamond
Graphite
The electronic configuration of carbon is 1s22s22p2, i.e.
An excited state of carbon with electronic configuration
sp2 hybridization
Mixing of an s electron cloud with
two p electron clouds
Bonding in graphite
In-plane bonding: covalent + metallic
Out-of-plane bonding: van der Waal’s bonding
Consequent properties of graphite
Van der Waal’s bonding between layers - ease of sliding between layers (application as lubricant)
Metallic bonding within a layer – high in-plane thermal and electrical conductivity
Bonding in benzene molecule
sp2 hybridization of the carbon atoms
Chemical composition of benzene is C6H6.
The carbon atoms are arranged in hexagonal ring. Single and double bonds alternate between the atoms.
CC
CC
C
CH
H
H
H
H
HStructure of benzene Simplified notations
2-27
Effect of atomic bonding on material properties
Modulus of elasticity
Related to the material stiffness.Defined as the amount that a material will
stretch when a force is applied.It is related to the slope of the force-
distance curve.
A steep dF/da slope gives a high modulus.
Coefficient of thermal expansion
Describes how much a material expands or contracts when its temperature changes.
Asymmetric energy trough resulting in thermal expansion phenomenon
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