1. Chapter 2: Atomic bonding

2. Reading assignment
Ch. 2 and 3 of textbook

3. Homework No. 1
Problems 2-8, 2-9, 2-13.

4. Homework No. 2
Problems 3-13, 3-15, 3-17, 3-19, 3-20, 3-21, 3-27.

5. 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

6. 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

7. Electronic configurations

8. 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 = 2
Al: 1s22s22p63s23p1 valence = 3
Ge: 1s22s22p63s23p63d104s24p2 valence = 4

9. Primary bonding types
Ionic bonding
Covalent bonding
Metallic bonding

10. Ionic bonding

12. Ionic bonding in NaCl 3s1 3p6 Chlorine Sodium Atom Atom Cl Na
Chlorine Ion
Cl -
Sodium Ion
Na+
2-15

14. Ionic bonding Ionic Bond:
The attractive bonding forces are coulombic (different polarities):

15. Interionic force Force of attraction between Na+ and Cl- ions
Z1 = +1 for Na+, Z2 = -1 for Cl-
e = x C , ε0 = 8.85 x C2/Nm2
a0 = Sum of Radii of Na+ and Cl- ions
= nm nm = 2.76 x m
2-18

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.

21. Bond energy Bonding force and Energy curves for a Na+ & Cl- pair.
Since F = dE/da, the equilibrium bond length (ao) F=0 and E is a minimum.

22. 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.

24. C.N. = 4 (in 2 dimensions) Unstable Stable Stable (critical case)
r/R > min
r/R = min
r/R < min

25. C.N. = 2
Schematic drawing with nearest neighbors
not in contact

26. C.N. = 3
Schematic drawing with nearest neighbors
not in contact

27. C.N.= 4
3 dimensions

28. C.N. = 4 (3 dimensions)

29. C.N. = 8
(3 dimensions)

30. C.N. = 12 (3 dimensions)

31. C.N. = 6
3 dimensions

32. C.N. = 6 (3 dimensions)

33. C.N. = 4 (in 2 dimensions) C.N. = 6 (in 3 dimensions) Unstable Stable
(critical case)
r/R > min
r/R = min
r/R < min

35. C.N. = 3

36. C.N. = 8 (3 dimensions)

39. C.N. = 4

42. 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.

44. 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.
kJ/mol (3-8 eV/atom) bonding energies are large  high Tm.

45. Directional bond due to the sharing of electrons between atoms
Covalent bonding
Directional bond due to the sharing of electrons between atoms

46. Cl2 molecule
Planetary model
Actual electron density
Electron dot schematic
Bond line schematic

47. Example 1. Br2 (a bromine molecule)
Br has an outermost electronic configuration of 4s24p5, i.e.,

49. A single bond
A -bond
End-to-end overlap

50. Example 2. O2 (an oxygen molecule)
O has an outermost electron configuration of 2s22p4, i.e.,

52. A double bond
A σ-bond (end-to-end overlap) together with a π-bond (side-to-side overlap).

53. Double bond
Single bonds

55. Every carbon atom along the chain is four-fold coordinated.

56. The electronic configuration of carbon is 1s22s22p2, i.e.

57. An excited state of carbon with electronic configuration

58. Mixing of an s electron cloud with three p electron clouds
sp3 hybridization
Mixing of an s electron cloud with three p electron clouds

60. Methane molecule

61. Methane molecule

63. 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

65. SiO4 tetrahedron in silicate glass

66. Silicon dioxide (SiO2)

68. Covalent network solids
Network of covalent bonds
Examples: diamond, silicon, etc.

69. Properties of covalent network solids
Materials have poor ductility.
Poor electrical conductivity.
Many ceramic, semiconductor and polymer materials are fully or partly covalent.

70. 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.

71. There is one unpaired electron.
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.

72. The electronic configuration of F is 1s22s22p5, i.e.
There is also one unpaired electron.

76. 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.

78. 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 cloud
2-28

80. Metallic Bonding Good thermal & electrical conductors
The free electrons in the “cloud” move freely under an applied voltage.
Good thermal &
electrical
conductors

81. 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.

82. Bond energy and melting temperature

83. Secondary bonding Van der Waals bonding
Secondary bonding exists between virtually all molecules, but its presence is diminished if any primary bond is present.

84. Dipoles are created when positive and negative charge centers exist.
Dipole moment=μ =q.d
q= Electric charge
d = separation distance +q -q d
Skewed electron cloud
2-30

85. Secondary bonds are due to attractions of electric dipoles in atoms or molecules.

86. Electric dipole types
Permanent dipoles
Induced dipoles

87. Dipoles that do not fluctuate with time
Permanent dipoles
Dipoles that do not fluctuate
with time

88. Example of a permanent dipole
Hydrogen fluoride HF

89. Permanent dipoles in water

90. Water
Attraction between positive oxygen pole and negative hydrogen pole.

91. Dipole-dipole interaction
Dipole-dipole interaction in water H O
105 0
Dipole-dipole interaction H
2-33

92. Methane
Vector sum of four C-H dipoles is zero.

93. Methane Methyl chloride CH4 CH3Cl Symmetrical arrangement
of 4 C-H bonds
No dipole
moment
CH4
Methyl chloride
Asymmetrical
tetrahedral
arrangement
Creates
dipole
CH3Cl
2-32

94. Cl- ions in green

95. 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.

96. Induced dipoles No permanent dipole moment
Statistical fluctuation in electron density distribution
London dispersion forces (weak)
Examples: argon (an inert gas), methane (CH4 - a symmetric molecule)

98. 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
Asymmetrical
distribution
(Changes with time)
2-31

100. London forces between
methane molecules

101. Another example of mixed bonding types in a material
Graphite

102. Crystal forms of carbon
Graphite
Diamond
Fullerene

103. Fullerene
(a molecule)

104. Diamond

105. Graphite

106. The electronic configuration of carbon is 1s22s22p2, i.e.

107. An excited state of carbon with electronic configuration

108. Mixing of an s electron cloud with two p electron clouds
sp2 hybridization
Mixing of an s electron cloud with two p electron clouds

111. Bonding in graphite In-plane bonding: covalent + metallic
Out-of-plane bonding: van der Waal’s bonding

112. 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

113. Bonding in benzene molecule
sp2 hybridization of the carbon atoms

114. C C C C C C Chemical composition of benzene is C6H6.
The carbon atoms are arranged in hexagonal ring.
Single and double bonds alternate between the atoms. H C H H C C C C H H C H
Structure of benzene
Simplified notations
2-27

116. Effect of atomic bonding on material properties

117. 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.

118. A steep dF/da slope gives a high modulus.

119. Coefficient of thermal expansion
Describes how much a material expands or contracts when its temperature changes.

120. Asymmetric energy trough resulting in thermal expansion phenomenon