1. Electron Configurations

2. Recall Quantum Mechanics…
Describes the arrangement of electrons in atoms in terms of:
Main or principal energy levels (n)
Energy subshells
Orbitals (space occupied within the atom)

3. Generally symbolized by “n”, it denotes the shell (energy level) in which the electron is located.

4. Main Energy Levels (Shells)
A group of electrons in an atom all having the same principal quantum number (n)
n = 1, 2, 3, …
The first shell (n = 1) is lowest in energy, 2nd level next and so on 1<2<3<4
The number of electrons in each shell is limited to 2n2

5. Subshells Energy sublevels within energy level
All electrons in a subshell have the same energy
Designated s, p, d, f
Sublevel energy: s<p<d<f

6. Electron Locations
Main Energy Levels Sublevels n=4 4s, 4p, 4d, 4f n=3 3s, 3p, 3d n=2 2s, 2p n=1 1s

7. Electron Configuration
List of subshells containing electrons
Written in order of increasing energy
Superscripts give the number of electrons
Example: Electron configuration of neon
number of electrons
1s2 2s2 2p6
main shell subshell

8. Writing Electron Configurations
H 1s1 He 1s2 Li 1s2 2s1 C 1s2 2s2 2p2

9. Inner transition metals
Elements can be sorted into 4 different groupings based on their electron configurations:
Noble gases
Representative elements
Transition metals
Inner transition metals

11. Electron Configurations in Groups
Noble gases are the elements in Group 8A
Previously called “inert gases” because they rarely take part in a reaction; very stable = don’t react
Noble gases have an electron configuration that has the outer s and p sublevels completely full (thus making them unreactive)

12. Electron Configurations in Groups
Representative Elements are in Groups 1A through 7A
Display wide range of properties, thus a good “representative”
Some are metals, or nonmetals, or metalloids; some are solid, others are gases or liquids
Their outer s and p electron configurations are NOT filled

13. Electron Configurations in Groups
Transition metals are in the “B” columns of the periodic table
Electron configuration is now filling the “d” sublevel
A “transition” between the metal area and the nonmetal area
Examples are gold, copper, silver

14. Electron Configurations in Groups
Inner Transition Metals are located below the main body of the table, in two horizontal rows
Electron configuration is now filling the “f” sublevel
Formerly called “rare-earth” elements, but this is not true because some are very abundant

15. Understanding Electron Configuration

16. Energy Levels
Each energy level has a number called the PRINCIPAL QUANTUM NUMBER, n
Currently n can be 1 thru 7, because there are 7 periods on the periodic table

17. Energy Levels
n = 1
n = 2
n = 3
n = 4

18. Electron Configurations
A list of all the electrons in an atom (or ion)
Must go in order (Aufbau principle)
We need electron configurations so that we can determine the number of electrons in the outermost energy level. These are called valence electrons.
The number of valence electrons determines how many and what this atom (or ion) can bond to in order to make a molecule

19. Diagonal Rule
The diagonal rule is a memory device that helps you remember the order of the filling of the orbitals from lowest energy to highest energy
1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p6 6s2 4f14… etc.

20. Diagonal Rule 1 2 3 4 5 6 7 s s 2p s 3p 3d s 4p 4d 4f s 5p 5d 5f
Steps:
Write the energy levels top to bottom.
Write the orbitals in s, p, d, f order. Write the same number of orbitals as the energy level.
Draw diagonal lines from the top right to the bottom left.
To get the correct order, follow the arrows! 1 2 3 4 5 6 7 s
s 2p
s 3p 3d
s 4p 4d 4f
s 5p 5d 5f
s 6p 6d 6f
s 7p 7d 7f

21. d and f orbitals require LARGE amounts of energy
It’s better to skip a sublevel that requires a large amount of energy (d and f orbtials) for one in a higher level but lower energy
This is the reason for the diagonal rule! BE SURE TO FOLLOW THE ARROWS IN ORDER!

22. Orbitals and the Periodic Table
Orbitals grouped in s, p, d, and f orbitals
s orbitals
d orbitals
p orbitals
f orbitals

23. Shorthand Notation A way of abbreviating long electron configurations
Since we are only concerned about the outermost electrons, we can skip to places we know are completely full (noble gases…recall the “octet rule”), and then finish the configuration

24. Shorthand Notation
Atoms like to either empty or fill their outermost level. The optimum number of valence electrons is eight. This is called the octet rule.

25. Shorthand Notation
Step 1: Find the closest noble gas to the atom (or ion), WITHOUT GOING OVER the number of electrons in the atom (or ion). Write the noble gas in brackets [ ].
Step 2: Find where to resume by finding the next energy level.
Step 3: Resume the configuration until it’s finished.

26. Shorthand Notation [Ne] 3s2 3p5 Chlorine
Longhand is 1s2 2s2 2p6 3s2 3p5
You can abbreviate the first 10 electrons with a noble gas, Neon. [Ne] replaces 1s2 2s2 2p6
The next energy level after Neon is 3
So you start at level 3 on the diagonal rule (all levels start with s) and finish the configuration by adding 7 more electrons to bring the total to 17
[Ne] 3s2 3p5

27. Valence Electrons Electrons are divided between core and
B: 1s2 2s2 2p1
Core = [He] , valence = 2s2 2p1