ELECTRONIC CONFIGURATION
It refers to how electrons are arranged around the nucleus.
Bohr’s Model
Based on the Bohr’s model of the atom, we can describe how many levels of energy around the nucleus are occupied by electrons and how many electrons are in each level. (see table below)
| SHELL | NUMBER OF ELECTRONS |
| 1 | 2 |
| 2 | 8 |
| 3 | 18 |
Before placing an electron in a level, all inner levels must be filled out first.
Quantum Model
The QUANTUM MODEL OF THE ATOM is much more specific than Bohr’s model.
It includes not only the LEVELS OF ENERGY (shells) but also SUB-LEVELS, and also explains the behavior of the electron as a wave instead of a particle.
Each level is made of sublevels. Different sublevels have a fixed maximum number of electrons that they can hold.
Each sublevel contains a certain number of orbitals. Orbitals are zones around the nucleus of the atoms whare is possible to find an electron. Each orbital can accommodate up to 2 electrons, only if they have opposite spin.
s-orbitals
p-orbitals
d-orbitals
f-orbitals
As we increase the levels of energy, more and more sublevels are added.
Also, as we go up in levels, the difference of energy between them gets smaller and smaller. It reaches a point where the sublevels begin to overlap.
Levels of Energy
Sublevels of energy
Why do we need electron configurations?
When atoms come into , it is the outermost electrons of these atoms, or valence shell, that will interact first. An atom is least stable when its valence shell is not full.
The valence electrons are responsible for an element’s chemical behavior. Elements that have the same number of valence electrons often have similar chemical properties.
They can also predict stability. An atom is most stable when all its orbitals are full. These configurations occur in the noble gases. The noble gases are very stable elements that do not react easily with any other elements.
Electron configurations can help us predict about the ways in which certain elements will react, and the chemical compounds or molecules that different elements will form.
Three Important Rules:
In order to fill out the electron configuration we need to follow the rules below:
AUFBAU PRINCIPLE: (about filling up the electrons in an atom) Electrons will fill the lower energy levels first, so a higher level cannot have an electron until the levels below are not completed.
PAULI EXCLUSION PRINCIPLE: no two electrons in the same atom can have the same set of quantum numbers; that is, cannot simultaneously occupy the same energy (quantum) state of an atom.
HUND’S RULE: Every orbital in a sublevel is singly occupied before any orbital is doubly occupied. All of the electrons in singly occupied orbitals have the same spin.
Electron configurations of atoms
It is a list showing how many electrons are in each shell, subshell in an atom or ion. subshell notation: list subshells of increasing energy, with number of electrons in each subshell as a superscript
- examples
- 1s22s22p5 means “2 electrons in the 1s subshell, 2 electrons in the 2s subshell, and 5 electrons in the 2p subshell”
- 1s22s22p63s23p3 is an electron configuration with 15 electrons total; 2 electrons have n=1 (in the 1s subshell); 8 electrons have n=2 (2 in the 2s subshell, and 6 in the 2p subshell); and 5 electrons have n=3 (2 in the 3s subshell, and 3 in the 3p subshell).
- ground state configurations fill the lowest energy orbitals first electron configurations of the first 11 elements, in subshell notation. Notice how configurations can be built by adding one electron at a time.
| Atom | Z | Ground state electronic configuration |
| H | 1 | 1s1 |
| He | 2 | 1s2 |
| Li | 3 | 1s22s1 |
| Be | 4 | 1s22s2 |
| B | 5 | 1s22s22p1 |
| C | 6 | 1s22s22p2 |
| N | 7 | 1s22s22p3 |
| O | 8 | 1s22s22p4 |
| F | 9 | 1s22s22p5 |
| Ne | 10 | 1s22s22p6 |
| Na | 11 | 1s22s22p63s1 |
We can use the periodic table to fill out the electron configuration beginning at the top and from left to right.
Writing electron configurations
Start with hydrogen, and build the configuration one electron at a time. Fill subshells in order by counting across periods, from hydrogen up to the element of interest.
watch out for d & f block elements; orbital interactions cause exceptions to the Aufbau principle
Half-filled and completely filled d and f subshells have extra stability. Know these exceptions to the Aufbau principle in the 4th period.
Examples: Give the ground state electronic configurations for:
- Al
- Fe
- Ba
- Hg
Some elements do not follow the Aufbau rule and this is because they are more stable with the “d” subshell half filled or totally filled.
| exception | configuration predicted by the Aufbau principle | true ground state configuration |
| Cr | 1s22s22p63s23p64s23d4 | 1s22s22p63s23p64s13d5 |
| Cu | 1s22s22p63s23p64s23d9 | 1s22s22p63s23p64s13d10 |
Electron configurations including spin
- Unpaired electrons give atoms (and molecules) special magnetic and chemical properties
- To show unpaired electrons we use orbital box diagrams
Examples of ground state electron configurations in the orbital box notation that shows electron spins.
- Since only the valence electrons are chemically active, the noble gas core under the valence shell is chemically inert.
- We write the notation for electron configurations by replacing the core with a noble gas symbol in square brackets:
Element full electron configuration core valence electrons Noble gas notation O 1s22s22p4 He 2s22p4 [He] 2s22p4 N 1s22s22p3 He 2s22p3 [He] 2s22p3 Cl 1s22s22p63s23p5 Ne 3s23p5 [Ne] 3s23p5 Mg 1s22s22p63s2 Ne 3s2 [Ne] 3s2 Al 1s22s22p63s23p1 Ne 3s23p1 [Ne] 3s23p1
Electron configuration of ions.
Most elements are more stable by losing or gaining electrons to fulfill the outermost shell with a Noble gas structure (chemically inert)
Examples:
Ca = 1s22s22p63s23p64s2
Ca2+ =1s22s22p63s23p6
Cl = 1s22s22p63s23p5
Cl– = 1s22s22p63s23p6
Elements that belong to the “d” block, lose the “s” electrons first.
Zn = 1s22s22p63s23p64s23d10 = [Ar]4s23d10
Zn2+ = 1s22s22p63s23p63d10 = [Ar]3d10