Chapter 7-Atomic Structure & Periodicity
Electromagnetic Radiation-a form of energy that exhibits wavelike behavior as it travels through
space.
c = v
c - speed of light, 3x 10 8 m/s
Wavelength (  ) - distance between two peaks on a wave. Measured in nm, cm, m.
1m = 1x 10 9 nm

Frequency
(v) - # of waves that pass a given point in a specific time. Measured in Hz.
E = hv

J
 E - energy of a wave or photon, 34
H - Planck’s constant, 6.626x 10 Js
ROYGBIV - red, orange, yellow, green, blue, indigo, violet
Visible light spectrum - ranges from 400nm to 700 nm, violet being 400nm and red being 700nm.

Bohr Model - e  moves
 to excited state (energy is absorbed), then e falls back to ground state
(energy emitted), only explains hydrogen atom.
1
1
E  (2.178x1018 )( 2  2 )
n final n initial


Quantum numbers are used to completely describe orbitals. They specify the properties of atomic
orbitals and the properties of the electrons in the orbitals.
Principal quantum # (n) - size and energy of orbital.

Angular momentum quantum # ( l ) - shape of an atomic orbital, it can have any # from 0 to (n-1).
Magnetic quantum # (m) - position of orbital in space, it can have any # from - l to l including 0.
Electron spin quantum # (ms) - electron spin, +1/2 or -1/2.
n
m
ms
# of
# of
l  shape
e ‘s orbitals

1
0
s
0
-1/2,1/2
2
1
0
s
0
-1/2,1/2
2
1
2
1
p
-1,0,1
-1/2,1/2
6
3
3
0
s
0
-1/2,1/2
2
1
1
p
-1,0,1
-1/2,1/2
6
3
2
d
-2,-1,0,1,2
-1/2,1/2
10
5
4
0
s
0
-1/2,1/2
2
1
1
p
-1,0,1
-1/2,1/2
6
3
2
d
-2,-1,0,1,2
-1/2,1/2
10
5
3
f
-3,-2,-1,0,1,2,3 -1/2,1/2
14
7
Heisenberg Uncertainty Principle - there is a limitation to how precise the position and momentum
of a particle can be known at a given time.
Electron Configuration - there are 3 ways to write electron configuration.
Electron configuration -1s22s22p6
Noble gas electron configuration - noble gas from row above and valence electrons [Ar]4s23d8
Orbital notation Aufbau’s principle - an e- occupies the lowest possible energy orbital (ground state).
Pauli exclusion principle - no 2 e-‘s can have the same 4 quantum #’s.
Hund’s rule - orbital’s of equal energy are each occupied by the 1st e- before any orbital is occupied
by a 2nd e-, and all e-‘s in singly occupied orbital’s must have the same spin.
Valence electrons - e-‘s in the outermost shell of an atom.
Core electrons - remaining e-‘s in the atom.
Paramagnetic - an element that has at least one unpaired e -.
Diamagnetic - an element that doesn’t have any unpaired e-‘s.
Groups = columns
Periods = rows
Atomic radius - half of the distance between nuclei of identical atoms in a covalent bond.
Increasing positive charge as you move from left to right on the periodic table which makes
electrons get pulled in closer to the nucleus making a smaller atomic radius. More energy
levels are occupied as you move down the periodic table, which makes it bigger.
Ionization energy - energy needed to remove one electron from an atom. Increasing positive
charge as you move across the periodic table, this makes it harder to remove an electron
from an atom, which requires higher ionization energy as you move across the periodic
table. As you move down the periodic table the valence electrons are further from the
nucleus, which makes it easier to remove an electron.
Electron affinity - the energy change that occurs with the addition of an electron to a gaseous atom.
There is an increasing positive charge as you move across the periodic table, which
means a higher electron affinity. Atomic radii and shielding increase as you move up and
down the periodic table, which means there is a lower electron affinity as you move down
the periodic table.
Electronegativity - the measure of the ability of an atom to attract electrons. Elements on the left
side of the periodic table attract electrons weakly because they have few valence electrons
and fewer protons. Elements at the bottom of the periodic table have more filled energy
levels so they attract electrons less strongly.