Atomic and Molecular
Structure
By: Kevin “Diminutive Vocabulary” Mospan
and Liz “Bottle Cap Tea Cup” Schultz
•Lewis Structures: Do you know how to draw them?
•Molecular Shapes: That fun process when you try to
represent 3D things on 2D paper
•Formal charge: Remember that? Closer to 0 is favorable,
and means that you’re Lewis structure is correct…probably
•Bond polarity: The difference in the electronegativity needs
to be between 0.4 and 1.7 for the bond to be polar
•Molecular polarity: symmetrical means things cancel and
are therefore not polar
All of these are things you can practice on our quiz! How exciting!
•Electromagnetic radiation
•
(p.295)
•. = c = 2.99792458 x 108 m/s
•Wavelength ()– the distance between successive peaks or troughs
•Frequency – the number of complete waves or cycles passing thru a point in
a certain amount of time
•Amplitude – the maximum height of the wave
Relating Wave Properties and
Energy
• Useful equation: E = h . 
•Energy of a photon is inversely proportional to the wavelength.
i.e. ↑  means ↓ Energy
• Energy of a photon is directly proportional to the frequency.
i.e. ↑  means ↑ Energy
• There are 6.02 x 1023 photons of energy per mole of light
Avogadro's Number strikes again!
• Planck’s constant is h ≈ 6.6260755 x 10-34 J.s
•Useful equation: E = h.
•Photoelectric effect:
•Phenomenon explained by Einstein using Planck’s ideas
•Occurs when light strikes the surface of a metal, and electrons
are ejected.
•Another useful equation: E = h . c / 
•Electrons lie in circular orbits around molecule
•Radius of the circular orbits increases as n increases
•An atom with its electrons in the lowest possible energy levels
is said to be in its “ground state”
•When an electron occupies an orbit greater than the lowest
possible energy level, it is said to be in an “excited state”
•Somewhat useful equation: ΔE = -Rhc
•Rhc ≈ 1312 kJ/mol
(
)
1 - 1
n f2
n i2
Not with swords. Or
guns. Wave/Particle
duality applies what
Einstein and Planck
said about light photons
to massless electrons.
Electrons also exhibit
both wave and particle
properties. This is why
we can use those
“useful” light equations
on the previous slides
for electrons as well.
Actually, it’s DeBroglie and he had some valid things
to say. For example:
•He utilized the idea of wave/particle duality
to create yet another “useful” equation:
=
h
m .
Just make sure m (mass) is in kg
There are 4 Quantum Numbers: n , l , ml , ms
•n is the principle quantum number
•it can be any integer from 1 to 
•l is the angular momentum
quantum number. It characterizes
the subshell of an orbital, like in the table
•It starts at 0 and goes to n-1
Value of l
Corresponding subshell
label
0
s
1
p
2
d
3
f
•ml is the magnetic quantum number
•0,  1,  2, … ,  l
•Number of values for it in a given subshell is 2l + 1
•Ms is the electron spin magnetic quantum #. It’s only ½ or -1/2
Orbitals are the places where electrons like to kick it.
The lowest energy orbital is s, which has 1 subshell.
Next is p, with 3 subshells, then d with 5 subshells,
followed by f with 7 subshells.
There are subshells after these, but they don’t really
matter for the AP test because there are no known
elements to fill them
Here’s the easiest way to figure out the electron
configuration :
All you have to do is alter the Periodic Table by
putting the Lanthanides in the proper places as
shown by the picture.
Then you find your
element. That’s
basically it.
Quantum numbers for electron configurations
After you’ve figured out what the electron configuration is by
using the period table, you just fill in the little boxes until you’ve
filled everything up to and including the configuration you
found.
1s
2s
2p
3s
3p
4s
3d
4p
n is the number preceding the orbital for the selected atom
l is still like it was in the table
ml goes from the lowest number to the highest, left to right.
Ms has the +1/2 pointing up and the -1/2 pointing down
The Pauli exclusion principle says that no 2 electrons can have
the same set of all 4 quantum numbers
There are 3 major periodic trends

1. Radius
2. Electronegativity
3. Ionization Energy



Radius
•As one goes from right to left, the number of protons decreases, so there are
less intramolecular attractions causing the orbitals to expand
•As one goes from top to bottom, there are new orbitals which take up more
space and increase the radius
•Cations are smaller than the original atom because it loses electrons which
causes the orbitals to empty or constrict
•Anions are larger than the original atom because there is less force of proton
per electron, so orbital can spread out farther
Electronegativity
Ionization Energy
This theory mainly deals with
hybridization.
Here’s a helpful table for
hybridization if you can see it
It’s also on your summary cheat
sheet.
Valence bond theory says that there are 2 types of
bonds: Sigma (σ) bonds and Pi (Π) bonds
Sigma bonds occur when hybridized orbitals
directly overlap.
Pi bonds form from the unhybridized orbitals left
after hybridization.
In order for a pi bond to form, there must first be a
sigma bond, because sigma bonds are much
stronger.
Probably the most important part about molecular
bond theory is that there are two kinds of orbitals,
bonding orbitals and antibonding orbitals.
•In these orbitals, the antibonding orbitals are of much
higher energy than the bonding orbitals.
•A group of atoms is “stabilized” and form chemical
bonds when they are in the lower energy, bonding
orbitals
The last principle of molecular
bond theory is that electrons
are assigned to orbitals of
successively higher energy.
This means that if you have a
chart like that to the right, you
start in the bottom and fill
electrons in the spaces going
up. Also, it’s important to
remember that you fill in this
chart with all of a molecule’s
electrons, not just the valence
electrons.
This nifty chart also helps
illustrate how a filled bonding
orbital and a filled antibonding
orbital cancel each other out. If
both orbitals are filled with
electrons, there is no net bond
for this reason.
Lastly, bond order =
(1/2)(# of electrons in bonding orbitals - # of
electrons in antibonding orbitals)