1. Today's Agenda ISSUES TO ADDRESS... • What promotes bonding?
• What types of bonds are there?
• What properties are inferred from bonding? 1

2. GECKO
WHY ?

3. Here is the answer

4. Chapter 2: Atomic Structure and Interatomic Bonding
Electron Configuration
Periodic Table
Primary Bonding
Ionic
Covalent
Metallic
Secondary Bonding or van der Waals Bonding
Three types of Dipole Bonding
Molecules

5. REVIEW OF ATOMIC STRUCTURE (FRESHMAN CHEMISTRY)
ATOMS = (PROTONS+NEUTRONS) + ELECTRONS
NUCLEUS
BONDING
Mass of an atom:
Proton and Neutron: ~ 1.67 x kg
Electron: 9.11 x kg
Charge:
Electrons and protons: (±) 1.60 x C
Neutrons are neutral
The atomic mass (A): total mass of protons + total mass of neutrons
Atomic weight ~ Atomic mass
# of protons are used to identify elements (Z)
# of neutron are used to identify isotopes ( e.g. 14C6 and 12C6 )
Isotopes are written as follows: AXZ , i.e. 1H1, 2H1, 3H1
Actually neutron is a little heavier than proton but no problem with that…
Electrons are so much lighter so neglect in calculating the weight of an atom…
Hydrogen, deuterium, tritium
How is Carbon 14 technique used to measure the age of old artifacts ?

6. AMU and Mole Concept Nav = 1 gram/1 amu.
The atomic mass unit (amu) :
1 amu is defined as the 1/12 of the atomic mass of the most common isotope of carbon, carbon 12 (12C6).
Atomic mass of 12C6 is 12 amu : Carbon= 6 protons (Z=6) + 6 neutrons (N=6)
Mproton ~ Mneutron = 1.67 x g = 1 amu
Atomic Mass = Z + N
The atomic weight is often expressed in mass per mole
A mole is the amount of matter that has a mass in grams equal to the atomic mass in amu of the atoms (A mole of carbon has a mass of 12 grams).
The number of atoms in a mole is called the Avogadro number, Nav = × 1023.
Nav = 1 gram/1 amu.
Example:
Atomic weight of iron = amu/atom = g/mol
Inverse of mass of neutron + protons = Avagadro’s number

7. MOLE CONCEPT
ANY IDEAS WHY ?

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

9. Atomic Models
Towards the end of 19th century physicists realized Newtonian physics has serious difficulty in explaining many phenomena involving electrons => quantum mechanics
Bohr atomic model
Electrons assume very well defined orbits around the nucleus (protons + neutrons)
Electrons in each shell orbit assumes the same energy level
Severe issues when considering events involving electrons such as emission spectra and photoelectrons…)
See figure 2.1 in Callister page 18 (page 13 for 6th ed.)
Wave mechanical model
Electrons in an atom or molecule are permitted to have only specific values of energy, energy is quantized…
Electrons do not move in circular orbits but in “fuzzy” orbits. At any given time we can only talk about the probability of finding an electron at a radius from the orbit.
Every electron is characterized by four quantum numbers. The size, shape, spatial orientation of an electrons probability density are specified by these numbers
BOHR model represents an early attempt to describe electrons in atoms, interms of both position (electron orbital) and energy (quantized energy levels)

10. BOHR ATOM Nucleus: Z = # protons = 1 for hydrogen to 94 for plutonium
Adapted from Fig. 2.1, Callister 6e.
Nucleus: Z = # protons
= 1 for hydrogen to 94 for plutonium
N = # neutrons
Atomic mass A ≈ Z + N 2

11. Beyond Bohr’s Model In 1924 de Broglie : dual character of electrons
In 1927 Heisenberg : uncertainity, it is not possible to measure simultaneously both the momentum (or velocity) and the position of a microscopic particle with absolute accuracy.
Schrodinger, math expression for the behavior of an electron around an atom

12. FUZZY ORBITS

13. ELECTRON ENERGY STATES
Electrons...
• have discrete energy states
• tend to occupy lowest available energy state.
Adapted from Fig. 2.5, Callister 6e. 3

14. A product of Schrodinger’s Equation
Quantum Numbers (I)
A product of Schrodinger’s Equation
n, l , ml , ms
n principal quantum number, distance of an electron from the nucleus
l subshell, describes the shape of the subshell
ml number of energy states in a subshell
ms spin moment
Pauli’s exclusion principle: only one electron can have a given set of four quantum numbers.

15. Quantum Numbers (II) l ml
ms = ±½

16. Quantum Numbers (III)
Electrons fill quantum levels in order of increasing energy ( only n and l make significant differences in energy configurations).
1s, 2s, 2p, 3s,3p,4s,3d,4p,5s,4d,5p,6s,4f,5d,….
When all electrons are at the lowest possible energy levels => ground state
Excited states do exist such as in glow discharges etc…
Valence electrons occupy the outermost filled shell.
Valence electrons are responsible for all bonding !

17. SURVEY OF ELEMENTS • Most elements: Electron configuration not stable.
Adapted from Table 2.2, Callister 7e.
• Why? Valence (outer) shell usually not filled completely. 5

18. STABLE ELECTRON CONFIGURATIONS
• have complete s and p subshells
• tend to be unreactive.
Adapted from Table 2.2, Callister 6e. 4

19. 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

20. THE PERIODIC TABLE
• Columns: Similar Valence Structure, Similar Properties
Electropositive elements:
Readily give up electrons
to become + ions.
Electronegative elements:
Readily acquire electrons
to become - ions. 6

21. Periodic Table
Draft of the first periodic table, Mendeleev, 1869

22. ELECTRONEGATIVITY • Ranges from 0.7 to 4.0,
• Large values: tendency to acquire electrons; reactivity
Metals like to give up, halogens like to acquire electrons !
Smaller electronegativity
Larger electronegativity 7

23. Concept Checks
Question: Why are the atomic weights of the elements generally not integers? Cite two reasons.
Answer: The atomic weights of the elements ordinarily are not integers because:
(1) The atomic masses of the atoms normally are not integers (except for 12C), and
(2) the atomic weight is taken as the weighted average of the atomic masses of an atom's naturally occurring isotopes.
Question: Give electron configurations for the Fe3+and S2- ions.
Answer: The Fe3+ ion is an iron atom that has lost three electrons. Since the electron
configuration of the Fe atom is 1s22s22p63s23p63d64s2 (Table 2.2), the configuration for Fe3+ is 1s22s22p63s23p63d5.
The S2- ion a sulfur atom that has gained two electrons. Since the electron configuration of the S atom is 1s22s22p63s23p4 (Table 2.2), the configuration for S2- is 1s22s22p63s23p6.

24. Atomic Bonding in Solids
Start with two atoms infinitely separated
Attractive component is due to nature of the bonding (minimize energy thru electronic configuration)
Repulsive component is due to Pauli exclusion principle; electron clouds tend to overlap
Essentially atoms either want to give up (transfer) or acquire (share) electrons to complete electron configurations; minimize their energy
Transfer of electrons => ionic bond
Sharing of electrons => covalent
Metallic bond => sea of electons r

25. IONIC BONDING (I) • Occurs between + and – ions (anion and cation).
• Requires electron transfer.
• Large difference in electronegativity required.
• Example: Na+ Cl-
Easiest to describe…. 8

26. 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]

27. IONIC BONDING (II)
Easiest to describe….
Oppositely charged ions attract, attractive force is coulombic.
Ionic bond is non-directional, ions get attracted to one another in any direction.
Bonding energies are high => 2 to 5 eV/atom,molecule,ion
Hard materials, brittle, high melting temperature, electrically and thermally insulating 8

28. 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
Adapted from Fig. 2.8(b), Callister 7e.

29. 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.

30. COVALENT BONDING (I) • Requires shared electrons • Example: CH4
C: has 4 valence e,
needs 4 more
H: has 1 valence e,
needs 1 more
Electronegativities
are comparable.
In covalent bonding, electrons are shared between the
molecules, to saturate the valency. The simplest example is
the H2 molecule, where the electrons spend more time in
between the nuclei than outside, thus producing bonding.
Adapted from Fig. 2.10, Callister 6e. 10

31. COVALENT BONDING (II) Diamond, sp3
Covalent bonds are formed by sharing of the valence electrons
Covalent bonds are very directional
Covalent bond model: an atom can have at most 8-N’ covalent bonds, where N’ = number of valence electrons
Covalent bonds can be very strong, eg diamond, SiC, Si, etc, also can be very weak, eg Bismuth
Polymeric materials do exhibit covalent type bonding.
In covalent bonding, electrons are shared between the
molecules, to saturate the valency. The simplest example is
the H2 molecule, where the electrons spend more time in
between the nuclei than outside, thus producing bonding. 10

32. Primary Bonding Metallic Bond -- delocalized as electron cloud
Ionic-Covalent Mixed Bonding
% ionic character =  
where XA & XB are Pauling electronegativities %)
100 ( x
Ex: MgO XMg = XO = 3.5

33. COVALENT BONDING (III)
Very few materials have pure ionic or covalent bonding; electronegativity inpart defines how much time electrons spend between ion cores…
In covalent bonding, electrons are shared between the
molecules, to saturate the valency. The simplest example is
the H2 molecule, where the electrons spend more time in
between the nuclei than outside, thus producing bonding. 10

34. EXAMPLES: COVALENT BONDING
Adapted from Fig. 2.7, Callister 6e. (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.
• Molecules with nonmetals
• Molecules with metals and nonmetals
• Elemental solids (RHS of Periodic Table)
• Compound solids (about column IVA) 11

35. METALLIC BONDING • Arises from a sea of donated valence electrons
(1, 2, or 3 from each atom).
Ion cores in the “sea of electrons”.
Valance electrons belong no one particular atom but drift throughout the entire metal.
“Free electrons” shield +’ly charged ions from repelling each other…
Adapted from Fig. 2.11, Callister 6e.
• Primary bond for metals and their alloys 12

36. 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 + -
-general case:
secondary
bonding
Adapted from Fig. 2.14,
Callister 7e. Cl Cl
-ex: liquid HCl
secondary H H
bonding
secondary bonding
-ex: polymer
secondary bonding

37. Bonding Energies

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

39. 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.

40. PROPERTIES FROM BONDING: E
• Elastic modulus, E
• E ~ curvature at ro
E is larger if Eo is larger. 16

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

42. 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

43. ANNOUNCEMENTS
Read Chapter 3!!!