1. Electron Configuration
SC1. Obtain, evaluate, and communicate information about the use of the modern atomic theory and periodic law to explain the characteristics of atoms and elements.
g. Develop and use models, including electron configuration of atoms and ions, to predict an element’s chemical properties.

2. Electron Cloud Bohr’s model made the assumption that atoms are flat
Electron location can be described with 1 coordinate, or quantum number
n for energy level
Schrödinger’s model expanded the atom into 3-dimensional space – the electron cloud
This required more quantum numbers to describe the location of electrons
l for subshell, ml for orbital, ms for spin

3. Electron Cloud
Probable locations of electrons around the nucleus (based on their energy) can be predicted within the cloud in regions called subshells
There are 4 kinds of subshells in use
s (l=0), p (l=1), d (l=2), f (l=3)

4. Electron Cloud Variations on subshells are called orbitals
s only has 1 orbital
ml =0
p has 3 orbitals
ml = -1, 0, 1
d has 5 orbitals
ml = -2, -1, 0, 1, 2
f has 7 orbitals
ml = -3, -2, -1, 0, 1, 2, 3

5. Shell-Subshell-Orbital

6. Quantum Numbers
Quantum numbers are like the address for a particular electron in an atom.
No two electrons, within an atom, can have the same set of 4 quantum numbers
Carbon
electron #1: n=1, l=0, ml=0, ms=+½
electron #2: n=1, l=0, ml=0, ms=-½
electron #3: n=2, l=0, ml=0, ms=+½
electron #4: n=2, l=0, ml=0, ms=-½
electron #5: n=2, l=1, ml=-1, ms=+½
electron #6: n=2, l=1, ml=0, ms=+½

7. Assigning Quantum Numbers
Three rules:
electrons fill subshells from lowest to highest energy (Aufbau)
electrons fill orbitals singly before any orbital gets a second electron (Hund)
no two electrons can fill one orbital with the same spin (Pauli)

8. Aufbau Principle 7s 7p 7d 7f 6s 6p 6d 6f 5s 5p 5d 5f 4s 4p 4d 4f

10. Blocks and Subshells
We can use the periodic table to predict which subshell is being filled by a particular element

11. Hund’s Rule
Electrons fill orbitals INDIVIDUALLY before PAIRING UP!

12. Pauli Exclusion Principle
Electrons PAIR UP with OPPOSITE spins
ms= +½ or -½
2 electrons per orbital
Subshell maxima:
s – 2
p – 6
d – 10
f – 14

13. Orbital Notation
The orbital notation of an atom is a method of drawing the location of electrons.
Aufbau is used to determine the order of filling, and arrows (either ↑ or ↓) are used to represent electrons.

15. Electron Configuration Notation
The electron configuration notation of an atom is a shorthand method of writing the location of electrons.
The energy level number is written first, followed by the subshell letter, and finally a superscript representing the number of electrons present.

17. Electron Configuration Notation
There are a few exceptions to the “rules”
If the e-config ends in d4, d9, f6 or f13
Chromium:
Copper:

18. Noble Gas Notation
This is a shorthand version of electron configuration notation
You simply replace core electrons with a bracketed noble gas symbol
Li = [He] 2s1
Si = [Ne] 3s23p2
Br = [Ar] 4s24p5
W = [Xe] 6s24f145d4

20. Valence Electrons
When an atom undergoes a chemical reaction, only the outermost electrons are involved.
These electrons, called the valence electrons, are of the highest energy and are farthest away from the nucleus.
The valence electrons are typically the s and p electrons in the highest energy level represented.

21. Predicting Valence Electrons
When using the IUPAC designations for group numbers for the representative elements, the ones digit corresponds to the number of valence electrons.
The transition metals, which mostly have 2 (4s and 5s), will sometimes use d orbital electrons as valence electrons

22. Octet Rule
Atoms which have a full set of eight valence electrons have increased stability
The exceptions are elements 1-5. Why?
The only atoms that have achieved this alone are called the “noble gases.”
So how do the other atoms achieve this state?
By sharing, gaining, or losing valence electrons!

23. Electron Dot Formulas
An electron dot formula of an element shows the symbol of the element surrounded by its valence electrons.
We use one dot for each valence electron.
Consider phosphorus, P, which has 5 valence electrons. Here is the method for writing the electron dot formula.

24. Bonding
When atoms share valence electrons, it is called covalent bonding.
Nonmetals typically share with other nonmetals
All atoms involved in the bond share enough valence electrons to achieve the stable octet.
There are combinations that involve more than an octet (SF6 or PCl5)

25. Bonding
When atoms lose or gain valence electrons, it is called ionic bonding
Metals lose valence electrons and form cations, and nonmetals gain valence electrons and form anions
The electrostatic attract between positive and negative cause the ionic bond to form

26. Predicting Ionic Charge
Because there is a pattern of valence electrons, there is a pattern of ionic charges:
Group 1
Group 2
Group 13
Group 14
Group 15
Group 16
Group 17

27. Ion Electron Configurations
When we write the electron configuration of a positive ion, we remove one electron for each positive charge:
Na → Na+
1s2 2s2 2p6 3s → 1s2 2s2 2p6
When we write the electron configuration of a negative ion, we add one electron for each negative charge:
O → O2-
1s2 2s2 2p4 → 1s2 2s2 2p6