Electrons and Electron Configuration
Electrons:
Energy Levels (orbitals, shells, clouds)
 On the outer rings surrounding the nucleus in
energy levels (clouds, shells, or orbitals)
 Each orbital represents an energy level in which
an electron can inhabit
o Energy levels are available even if they are
empty (unoccupied by electrons)
 When energy or heat is applied in the correct
amount an electron will jump to a higher energy
level becoming Excited!
o It will NEVER be in between orbital levels
o Greater the distance jumped = more energy
needed
o In an excited state
Atoms wish to have as little energy as possible at all
time!
 ASAP – electrons will go back down to the lowest
available energy level, releasing all of the
absorbed energy
o Sometimes it results in visible light!
o Red and Orange = low energy
o Blue and Violet = high energy
 Atoms can be identified by the color of light they
produce when exited!
Electrons want to be as close to the nucleus as possible
because there are protons in the nucleus BUT there is
a limit to how many e-’s can fit on an energy level.
Electron Configuration
Rules
 Electrons “fill up” energy levels from lowest to
highest
o Lowest is level 1
o When level 1 is full, the next e- must occupy
level 2
o When level 2 is full, e-‘s go to level 3
o Continues until all electrons are placed
Aufbau Principle
 Energy levels contain sublevels
 # of sublevels # of main energy levels
o Level 1 has ONE sublevel (s)
o Level 2 has TWO sublevels (p)
o Level 3 has THREE sublevels (d)
o Level 4 has FOUR sublevels (f)
s sublevel
 Every level has an s sublevel
 The name of the sublevel is written next to the # of
the level
o 1s, 2s, 3s, etc.
 Able to fit 2 electrons ONLY
p sublevel
 Starts at level 2 and higher
o 2p, 3p, 4p, etc.
 Level 1 DOES NOT have a p sublevel
 Able to fit 6 electrons ONLY
d sublevel
 ONLY on levels 3 and higher
o 3d, 4d, 5d, etc.
 Able to fit 10 electrons ONLY
f sublevel
 ONLY on level 4
o 4f, 5f, 6f, etc.
 Able to fit 14 electrons ONLY
Unfortunately electrons do not fill in logical manner, it
is seemingly random
1s2  2s2  2p6  3s2  3p6  4s2  3d10  4p6  5s2  4d10 
5p6 6s2  4f14  5d10  6p6  7s2  5f14  6d10  7p6  8s2
Diagonal Rule
Two ways of showing the location of electrons
Orbital Diagram
8 electrons =
Oxygen
Electron configuration – Short-hand method




Oxygen (8 electrons) – 1s2 2s2 2p4 = [He]2p4
Hydrogen (1 electron) - 1s1
Carbon (6 electrons) - 1s2 2s2 2p2 = [He]2p2
Titanium (22 electrons) - 1s2 2s2 2p6 3s2 3p6 4s2 3d2
= [Kr] 4s23d2
More Rules
 Identify the total number of electrons in the
element or ion
 Always begin with 1s
 Represent the actual # of electrons in each
sublevel with a superscript number
 Follow the “diagonal rule” to determine filling
order
Pauli Exclusion Principle
 The most an orbital can hold is two electrons
 They must spin in opposite directions
 Spin is the quantum mechanical property of
electrons and may be thought of as clockwise or
counter clockwise
 Use arrows to indicate the spin of an electron
Hund’s Rule
 The electrons will fill each orbital up before
sharing with another electron.
 Spin in the same direction with the same energy
Electron Configuration can help determine the
element
Which neutral atom has the following configuration:
1s22s22p63s23p64s23d104p5
 Simply add all the superscripts to find the total
number of electrons
 2+2+6+2+6+2+10+5 = 35 electrons
 Electrons = Protons
 35 protons is Bromine
Paramagnetic Vs Diamagnetic
Paramagnetic = an element with an unpaired electron
 Quantum number 1 because it as an uneven spin
either clockwise or counter clockwise
Diamagnetic = an element with every electron paired
 Quantum number 0 because the net spin is
canceled out.
Practice!
What is the electron configuration for Silicon?
What is the electron configuration for Nickel?
What is the electron configuration for ____?