1. CHAPTER 2: BONDING AND PROPERTIES
ISSUES TO ADDRESS...
• What promotes bonding?
• What types of bonds are there?
• What properties are inferred from bonding?

2. Atomic Structure (Freshman Chem.)
atom – electrons – x kg  protons 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 x 1023 molecules or atoms   1 amu/atom = 1g/mol
C H etc.
} 1.67 x kg

3. Atomic Structure
Valence electrons determine all of the following properties
Chemical
Electrical
Thermal
Optical

4. Electronic Structure
Electrons have wavelike and particulate properties.
This means that electrons in orbitals are defined by a probability.
Each orbital at discrete energy level determined by quantum numbers.   Quantum # Designation
n = principal (energy level-shell) K, L, M, N, O (1, 2, 3, etc.)
l = subsidiary (orbitals) s, p, d, f (0, 1, 2, 3,…, n -1)
ml = magnetic 1, 3, 5, 7 (-l to +l)
ms = spin ½, -½

5. Electron Energy States
Electrons...
• have discrete energy states
• tend to occupy lowest available energy state. 1s 2s 2p
K-shell n = 1
L-shell n = 2 3s 3p
M-shell n = 3 3d 4s 4p 4d
Energy
N-shell n = 4
Adapted from Fig. 2.4, Callister 7e.

6. SURVEY OF ELEMENTS • Most elements: Electron configuration not stable.
... 1s 2 2s 2p 6 3s 3p 3d 10 4s 4p
Atomic # 18 36
Element 1
Hydrogen
Helium 3
Lithium 4
Beryllium 5
Boron
Carbon
Neon 11
Sodium 12
Magnesium 13
Aluminum
Argon
Krypton
Adapted from Table 2.2, Callister 7e.
• Why? Valence (outer) shell usually not filled completely.

7. Electron Configurations
Valence electrons – those in unfilled shells
Filled shells more stable
Valence electrons are most available for bonding and tend to control the chemical properties
example: C (atomic number = 6)
1s2 2s2 2p2
valence electrons

8. Electronic Configurations26
1s2 2s2 2p6 3s2 3p6
3d 6 4s2
valence
electrons
ex: Fe - atomic # = 1s 2s 2p
K-shell n = 1
L-shell n = 2 3s 3p
M-shell n = 3 3d 4s 4p 4d
Energy
N-shell n = 4
Adapted from Fig. 2.4, Callister 7e.

9. The Periodic Table • Columns: Similar Valence Structure inert gases
give up 1e
give up 2e
accept 2e
accept 1e
give up 3e O Se Te Po At I Br He Ne Ar Kr Xe Rn F Cl S Li Be H Na Mg Ba Cs Ra Fr Ca K Sc Sr Rb Y
Adapted from Fig. 2.6, Callister 7e.
Electropositive elements:
Readily give up electrons
to become + ions.
Electronegative elements:
Readily acquire electrons
to become - ions.

10. Electronegativity • Ranges from 0.7 to 4.0,
• Large values: tendency to acquire electrons.
Adapted from Fig. 2.7, Callister 7e. (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.
Smaller electronegativity
Larger electronegativity

11. Ionic bond – metal + nonmetal
donates accepts electrons electrons  
Dissimilar electronegativities  
ex: MgO Mg 1s2 2s2 2p6 3s O 1s2 2s2 2p4
[Ne] 3s2 
Mg2+ 1s2 2s2 2p O2- 1s2 2s2 2p6
[Ne] [Ne]

12. Ionic Bonding - + • Occurs between + and - ions.
• Requires electron transfer.
• Large difference in electronegativity required.
• Example: NaCl
Na (metal)
unstable
Cl (nonmetal)
electron + -
Coulombic
Attraction
Na (cation)
stable
Cl (anion)

13. Ionic Bonding Energy – minimum energy most stable r A B EN = EA + ER =
Energy balance of attractive and repulsive terms r A n B
EN = EA + ER = + -
Attractive energy EA
Net energy EN
Repulsive energy ER
Interatomic separation r
A & B are constants that depend on the specific ionic system, while n is typically = 8. Repulsion diminishes more quickly with distance due to the exponential factor n.
The magnitude of this bonding energy and the shape of the energy-versus interatomic separation curve vary from material to material, and they both depend on the type of atomic bonding. Furthermore, a number of material properties depend on Eo, the curve shape, and bonding type.
Adapted from Fig. 2.8(b), Callister 7e.

14. Examples: Ionic Bonding
• Predominant bonding in Ceramics
NaCl
MgO
Give up electrons
Acquire electrons
CaF 2
CsCl
Adapted from Fig. 2.7, Callister 7e. (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.

15. Covalent Bonding similar electronegativity  share electrons
bonds determined by valence – s & p orbitals dominate bonding
Example: CH4
shared electrons
from carbon atom
from hydrogen
atoms H C CH 4
C: has 4 valence e-,
needs 4 more
H: has 1 valence e-,
needs 1 more
Covalent bonds are directional, meaning that atoms so bonded prefer specific orientations relative to one another; this in turn gives molecules definite shapes, as in the angular (bent) structure of the H2O molecule.
Electronegativities
are comparable.
Bonding is directional.
Adapted from Fig. 2.10, Callister 7e.

16. Metallic Bonding
Metallic Bond -- delocalized as electron cloud  

17. Mixed Bonding Ionic-Covalent Mixed Bonding % ionic character =
Ionic-Covalent Mixed Bonding
% ionic character =  
where XA & XB are Pauling electronegativities %)
100 ( x
Ex: MgO XMg = XO = 3.5

18. SECONDARY BONDING + - Arises from interaction between dipoles
• Fluctuating dipoles
asymmetric electron
clouds + -
secondary
bonding H 2
ex: liquid H
Adapted from Fig. 2.13, Callister 7e.
• Permanent dipoles-molecule induced
Adapted from Fig. 2.14,
Callister 7e. H Cl
secondary
bonding + -
-general case:
-ex: liquid HCl
secondary bonding
-ex: polymer
secondary bonding

19. IMPORTANCE OF SECONDARY BONDING
In some cases the type of bond explains materials’
properties
Gecko’s have very sticky feet that cling to any surface!
There are a large number of microscopic hairs on their
toe-pads. Which are self cleaning!
Weak intermolecular forces are established between the
hairs and any surface; resulting in adhesion.
Self cleaning reversible adhesives!

20. http://webecoist. momtastic

22. ANAMOLOUS BEHAVIOUR OF WATER
exhibits the anomalous and familiar expansion upon freezing—approximately 9 volume percent expansion.
its single O atom can bond to two hydrogen atoms of other water molecules.
1 water molecule has 4 H-bonds; leading to a non-close packed or open structure, hence expansion on freezing.
This structure is destroyed on melting so in liquid state the water molecules are more closely packed 1 molecule has 4.5 H-bonds leads to increase in density.
A watering can that ruptured along a side panel-bottom panel seam. Water that was left in the can during a cold late-autumn night expanded as it froze and caused the rupture.
The arrangement of water molecules in (a) solid ice, and (b) liquid water.

23. CONSEQUENCES OF ANAMOLOUS BEHAVIOUR OF WATER
This phenomenon explains why icebergs float.
• Why, in cold climates, it is necessary to add antifreeze
to an automobile’s cooling system (to keep the engine
block from cracking)
• Why freeze-thaw cycles break up the pavement in streets
and cause potholes to form.
• If there were no H-bonding water will boil at -800C!
Life on earth as we know it will not be possible!

24. Properties From Bonding: Tm
• Bond length, r
• Melting Temperature, Tm r o
Energy r
• Bond energy, Eo Eo =
“bond energy”
Energy r o
unstretched length
smaller Tm
larger Tm
Tm is larger if Eo is larger.

25. Properties From Bonding : a
• Coefficient of thermal expansion, a
coeff. thermal expansion D L
length, o
unheated, T 1
heated, T 2 ( - 1 = a T 2 ) D L o
• a ~ symmetry at ro r o
smaller a
larger a
Energy
unstretched length Eo
a is larger if Eo is smaller.

26. Summary: Bonding Type Bond Energy Comments Ionic Large!
Nondirectional (ceramics)
Covalent
Variable
large-Diamond
small-Bismuth
Directional
(semiconductors, ceramics
polymer chains)
Metallic
Variable
large-Tungsten
Nondirectional (metals)
small-Mercury
Secondary
smallest
Directional
inter-chain (polymer)
inter-molecular

27. Summary: Primary Bonds
Ceramics
Large bond energy
large Tm
large E
small a
(Ionic & covalent bonding):
Metals
Variable bond energy
moderate Tm
moderate E
moderate a
(Metallic bonding):
Polymers
Directional Properties
Secondary bonding dominates
small Tm
small E
large a
(Covalent & Secondary):
secondary bonding