2. Atomic Structure and Bonding
Chapter 2
Atomic Structure and Bonding
What holds atoms together anyway?

3. Atomic Bonding Bonds are in equilibrium when forces are equal o zero.
His occurs at some equilirbiurm distance, ro, or as r infinity.

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

5. Bohr Atomic Model (Niels Bohr, 1915 - 1962)
Nucleus
Bohr Atomic Model (Niels Bohr, )
Electrons have fixed energy and are resticed to specific “shells”
Bohr Atomic Model: very early
Electrons aranged in orbitals around nucleas.
Orbitals define positions of electrons
General QM: each electron has an energy level/state assoc. with orbital.
“What do we know about electrons?”
Heisenberg uncertainty - you can never know where an electron is at a given instant.
dual wave/particle nature
Must use notion of Electron clouds.
Bohr model insufficient
Electron shells

6. Electronic Structure
Electrons have wavelike and particulate properties.
This means that electrons are in orbitals 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

7. Electron Clouds (Orbitals)
Heisenberg – the Orbital is just a state of energy, a level of energy but even with the energy state known, the position of the electron is only known as a probability of its positions.
s,p,d and f orbitals give probability densities of positions in increasingly higher order energy states

8. Wave-mechanical Model and Quantum Numbers
Principal Quantum Number, n - 1,2,3,4,... => K.L.M.N, ... => Bohr Shells
Second Quantum Number, l - Subshells => s,p,d, and f
Third and Fourth - spin moment and orientation
Points of interest:
Farther from the nucleus, the energy level goes up.
Change in shell or subshell position requires energy loss/absorption
Move out - must have absorbed energy
Move in - gives off energy
Electrons on outer shell are called- valence electrons
v.e. take part in forming bonds with other atoms
Atoms like to have their outer shells filled.
Shell may loose, gain, or share electrons to complete shell.
Types of Chemical Bonds in Between Atoms

9. 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.

10. 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.

11. 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

12. 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.

13. Chem4kids.com

14. The Periodic Table • Columns: Similar Valence Structure give up 1e
inert gases
accept 1e
accept 2e 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.

15. The Periodic Table • Columns: Similar Valence Structure give up 1e
inert gases
accept 1e
accept 2e 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.

16. Types of Chemical Bonds in Between Atoms
Primary
Ionic
Covalent
Metallic
Secondary
Van der Waals Forces

17. 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)
Example: NaCl
Na's outer shell has one extra electron, which it gives up.
Cl's outer shell has a one empty location, that accepts from Na.
Na  Na+ and Cl  Cl-
The attractive forces between the two form the ionic bond.

18. Atomic Bonding Bonds are in equilibrium when forces are equal o zero.
His occurs at some equilirbiurm distance, ro, or as r infinity.

19. Ionic Bonding r A B EN = EA + ER =
Attractive and repulsive terms exist.
The _______ energy state is most stable r A n B
EN = EA + ER = -
Attractive energy EA
Net energy EN
Repulsive energy ER
Interatomic separation r
Force versus seperation distance curve Interatomic Seperation
Str depends w/ type of bond
Str. Vary with distance
Net = attract + repel
FN=FA+FR
Can also define bond energy.
A is found from the valences of atom and from physical constants.
Adapted from Fig. 2.8(b), Callister 7e.

20. 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.

21. 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
Electronegativities
are comparable.
Adapted from Fig. 2.10, Callister 7e.

22. Metallic Bonds
“Sea of Electrons”

23. 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

24. 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

25. 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.

26. 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.

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