Electron Dot and Orbital Notation
AUFBAU Principle – e-s first occupy the lowest energy level available
valence electrons – e-s in the highest energy levels, they are furthest from the nucleus
-
These are the e-s that form bonds and deal with reactivity
-
Shows only the valence e-s, the e-s in the out most energy level
Only up to 8 e-s
8 e-s = STABLE
Electron Dot Notation:
Ex:
H
He
Li
Be
Mg
Lithium shows 1e- because it only has 1e- in the outer most level  1s2 2s1
Li
C
 this is wrong!
*you can have dots paired or single, as long as there are 2 per side & no more
Ex1:
Ex2:
C
Ti = 22e-
OR
C
NOT
1s2 2s2 2p6 3s2 3p6 4s2 3d2
** 3d ARE NOT VALENCE
Valence eSo, titanium’s electron dot configuration is:
Ti
** ‘d’s’ are NEVER valence e-s, they ALWAYS fill AFTER a higher energy level
** the same applies for ‘ f ‘
C
Orbital Notation:
N ( 7 e-s ) ___
1s
___
2s
___ ___ ___
2p
F ( 9 e-s ) ___
1s
___
2s
___ ___ ___
2p
Hund’s Rule – orbitals of equal energy are each occupied by one e- before any orbital is occupied by a
second e-, and all e-s in singly occupied orbitals must have the same spin.
Energy Level
Sublevel
Orbitals per
sublevel
1
2
s
s
p
s
p
d
s
p
d
f
1
1
3
1
3
5
1
3
5
7
3
4
Total Orbitals
per energy
level
1
4
e-s per
sublevel
e-s per energy
level
2
2
6
2
6
10
2
6
10
18
2
8
9
16
18
32
Excited state – e-s in a higher than normal energy state
Ex: Nitrogen has 7 e-s, in ground state its configuration is  1s2 2s2 2p3
In an excited state, its configuration is 1s2 2s2 3s1 OR 1s2 2s2 2p2 3p1
Ions – atoms that have lost or gained electrons
Anions – are negatively ( - )charged, they have gained electrons
Cation – are positively ( + ) charged, they have lost electrons
**BIG IDEA**
-atoms will gain or lose e-s to become more stable
Na: 1s2 2s2 2p6 3s1 vs. Na+: 1s2 2s2 2p6
 no 3s1, Na loses this valence e- to become more stable
-Alkali metals, like Na, want to LOSE their 1 valence e- to become more stable.
Cl: 1s2 2s2 2p6 3s2 3p5 vs. Cl : 1s2 2s2 2p6 3s2 3p6
we go from 3p5 to 3p6 because Cl wants to fill its
last 3p orbital in order to become more stable
-Halogens, like Cl, want to GAIN a valence e- to become stable.
** filled and half-filled sublevels are more stable than partially filled sublevels.
Ex: Chromium = 24 e-  1s2 2s2 2p6 3s2 3p6 4s2 3d4
However, when filling orbitals, something funny happens!
___
1s
___
2s
___ ___ ___
2p
___
3s
___ ___ ___
3p
___
4s
___ ___ ___ ___ ___
3d
Although this SHOULD BE the orbital notation for Cr, we have a problem! Our 3d sublevel has 1 orbital
without an electron in it. Sublevels prefer to be either half full or completely full.
In this instance, Cr will take an e- from the 4s sublevel in order to make its 3d sublevel half full.
___
1s
___
2s
___ ___ ___
2p
___
3s
___ ___ ___
3p
___
4s
___ ___ ___ ___ ___
3d
NOW, our 4s orbital is half full as well as our 3d orbitals. Cr is stable, and is happy!
The same is true for Copper, by the following reasoning:
Copper has 29eCu: 1s2 2s2 2p6 3s2 3p6 4s2 3d9
___
1s
___
2s
___ ___ ___
2p
___
3s
___ ___ ___
3p
___
4s
___ ___ ___ ___ ___
3d
Again our 3d is not completely full, the last orbital would like to have 1 more electron! So it takes that
electron from the 4s orbital.
___
1s
___
2s
___ ___ ___
2p
___
3s
___ ___ ___
3p
___
4s
___ ___ ___ ___ ___
3d
NOW, our 3d is completely full and our 4s is half filled. Cu is more stable and happy!
**Take Home Message**
-orbitals (as well as sublevels) are stable when either FULL or HALF FULL as in the case with chromium
and copper.