This presentation covers all of Chapter 5 on the topic of Electron Configurations.

1. Chapter 5 Notes
By Kendon Smith
Columbia Central High School, Brooklyn, MI
Electron Configurations:
Electrons move about the nucleus
in a highly organized system!

2. Evolution of Atomic Models
1. Dalton Model (1803)
– John Dalton pictured atoms as tiny and
indestructible with no internal structure –
tiny, solid spheres!
DALTON MODEL

3. Evolution of Atomic Models
2. Thomson Model (1897)
– J. J. Thomson discovered the electron with
his Cathode ray tube experiment.
– Thomson pictured atoms as spheres of
positive charge embedded with
negatively charged electrons.
– Also called the Plum Pudding model,
electrons are stuck in atoms “like raisins in
plum pudding.”

4. Evolution of Atomic Models
2. Thomson Model (1897)
Electrons (-)
Positive Matrix (+)

5. Evolution of Atomic Models
3. Rutherford Model (1911)
– Ernest Rutherford discovered a solid core
called the nucleus with his gold foil
experiment.
– The nucleus is tiny, dense, and positively
charged!
– Electrons move around the nucleus in what
is mostly empty space!
* Rutherford’s atomic model could not explain

6. Evolution of Atomic Models
3. Rutherford Model (1911)
Electron Cloud (-) e-
e-
e-
Nucleus (+) +
e-
e-
e-

7. Evolution of Atomic Models
4. Bohr Model (1913)
- Neils Bohr proposed that electrons move
around the nucleus in circular paths, or
orbits, which are located at fixed distances
from the nucleus.
- Each electron orbits have fixed energies, so
these possible energies are called energy
levels.
- Electrons can jump from one energy level
to another by gaining or losing the right

8. Evolution of Atomic Models
4. Bohr Model (1913)
Electrons (-)
Nucleus (+) +

9. Evolution of Atomic Models
- A quantum of energy is the amount of energy
required to move an electron from one
energy level to another energy level.
- In general, the higher energy levels are
farther from the nucleus!
- The amount of energy an electron loses or
gains is not always the same.
- Higher energy levels are closer together,
requiring less energy to move electrons.

10. Ex. The Stair Step Analogy More Energetic
Electrons
Stairs = Energy Levels
6
Less Energetic 5
4
Increasing
Electrons
Energy!
3
2
1
*An electron needs less energy to jump from
Level 4 to level 5, than from level 1 to level 2

11. Evolution of Atomic Models
5. Quantum Mechanical Model
- The quantum mechanical model uses mathematical
equations to describe the behavior of electrons
within the electron cloud.
- Does not define an exact path the electron
takes around the nucleus, but rather,
electron location is described as a fuzzy
cloud where it spends most of its time.
- It calculates the probability of finding an
electron in a certain position.

12. Ex 1. The Dart Board Ex 2. Propeller Blades
Where is the fan blade
What are the chances my next when it’s moving?
dart will hit a particular ring?

13. II. Electron Configurations
A. Electron configurations are the ways in which
electrons are arranged in various orbitals
around the nucleus.
- Electrons move about the electron cloud in a
highly ORGANIZED system.
Example: What are the levels of organization
in your address? How does a letter
find you when it is mailed?

14. II. Electron Configurations
B. Levels of organization
1. Principle Quantum Number (Energy Level)
a. Indicates the main ENERGY LEVEL
surrounding a nucleus.
b. Symbol = n
Ex. Hydrogen – n = 1
Lithium – n = 2
3
Sodium – n =

15. 2. Orbital Quantum Number (Orbitals)
a. Also called: SUBLEVELS or SUBSHELLS
b. Symbols: s, p, d, f
c. Indicates the SHAPE of an orbital.
- s orbital = SPHERE
- p orbital = PEANUT
- d orbital = DOUBLE PEANUT
- f orbital = FAR TOO COMPLEX

16. s orbital = sphere

17. p orbital = peanut

18. d orbital = double peanut

19. f orbital = “far too complex”

20. d. The n energy level has n subshells
- 1st energy level has 1 subshells
- 2nd energy level has 2 subshells
- 3rd energy level has 3 subshells
- 4th energy level has 4 subshells
Principle
Quantum Number Types of Orbitals
1 1s
2 2s 2p
3 3s 3p 3d
4 4s 4p 4d 4f

21. C. Spin Quantum Number
a. Has only TWO possible values.
b. Indicates direction of ELECTRON SPIN.
c. Electrons spin on an imaginary axis, much like
the earth, thus generating a MAGNETIC field.
d. Symbols = ↑ ↓
e. Due to magnetism, electrons with OPPOSITE
spins pair up in each orbital, or subshell.
f. Each orbital is made up of an ELECTRON
PAIR.

22. C. Rules governing electron configurations
1. Aufbau principle
a. An electron occupies the LOWEST ENERGY
orbital that is available to receive it.

23. C. Rules governing electron configurations
2. Hund’s rule
a. Orbitals of equal energy are each occupied
by one electron before any one orbital is
occupied by a SECOND ELECTRON.
- For example, you would put one electron in
each p orbital, then come back and add the
second electron.
b. In orbitals that have only one electron, they
must have the SAME SPIN.

24. C. Rules governing electron configurations
3. Pauli Exclusion Principle
a. An atomic orbital may describe at most two
electrons.
b. Electrons that occupy the same orbital must
have opposite spins.

25. Electron States
1. Electrons in their lowest energy level are
said to be in the ground state.
2. By adding energy to an atom, an electron
can jump into a higher energy level. This
is called the excited state.
3. Almost immediately, the electrons fall
back to their original ground state, giving
off the added energy as visible light or
radiation.

26. Electron States
Atom becomes excited… Atom goes back to ground state.
E3 E3
E2 E2
E1 E1
+ +
Nucleus Nucleus
Energy LIGHT!

27. Some basic rules for electron configurations.
1. Atoms are in their ground state.
2. Atoms are neutral, having the same number of
protons and electrons.
3. Lowest energy orbitals fill first.
4. Each orbital can hold a pair of electrons with
opposite spin.
5. Completely fill all orbitals of the same energy
before starting to fill the next level.

28. Orbital types and quantities:
1. There is one s orbital in each energy level.
- holds a total of 2 electrons.
2. There are three p orbitals in each energy level.
- holds a total of 6 electrons.
3. There are five d orbitals in each energy level.
- holds a total of 10 electrons.
4. There are seven f orbitals in each energy level.
- holds a total of 14 electrons.

29. Energy levels and their orbitals.
- The n energy level contains n types of orbitals.
Energy Level Orbitals
1 1s
2 2s 2p
3 3s 3p 3d
4 4s 4p 4d 4f
5 5s 5p 5d 5f
6 6s 6p 6d 6f
7 7s 7p 7d 7f

30. Question: In what order do we fill orbitals?
Answer: Lowest energy orbitals first!
The order:
7p
Mr. Smith says, 6d
“You MUST 7s 5f
6p
memorize this 5d
order!!!” 6s 4f
5p
4d
4p 5s You say,
4s 3d “How in the world
3s 3p am I supposed to
2p
1s 2s memorize this?”

31. Use this chart to memorize the order to fill the orbitals!
Simply draw
s p d f
diagonal arrows
starting with the 1 1s
1s orbital.
2 2s 2p
Follow the arrows 3 3s 3p 3d
to fill orbitals in
each energy level. 4 4s 4p 4d 4f
5 5s 5p 5d 5f
Jump from the
head of one arrow 6 6s 6p 6d 6f
to the tail of the 7 7s 7p 7d 7f
next.

32. 3 Types of notations for configurations.
1. Orbital notation:
a. An unoccupied orbital is represented by:
b. An orbital occupied by a single electron: ↑
c. An orbital occupied by an electron pair: ↑↓
d. The lines are labeled with the proper principal
quantum number and subshell letter.
Ex. Sulfur ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑ ↑
1s 2s 2p 3s 3p

33. 2. Electron-Configuration Notation:
a. Uses no lines or arrows.
b. The number of electrons is shown as superscripts.
Ex. Sulfur
Orbital ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑ ↑
Notation 1s 2s 2p 3s 3p
Electron
Configuration 1s2 2s2 2p6 3s2 3p4
Notation

34. 3. Electron-Dot Notation:
a. Electron-dot notation shows only the electrons in
the highest energy level or valence shell.
b. Valence electrons – only those electrons located in
the outermost energy level.
c. Valence electrons are represented by dots placed in
pairs around the element symbol.
d. Most atoms can have 8 electrons in their valence
shell before the begin filling the next level.
Exceptions: Hydrogen and Helium – the first
energy level only has an s orbital so it fills up
with only 2 electrons!

35. Example: Sulfur
Orbital ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑ ↑
Notation 1s 2s 2p 3s 3p
Electron Configuration 1s2 2s2 2p6 3s2 3p4
Notation
6 valence
Electron Dot Notation electrons in
S the 3rd
energy level
*Note: Place single electrons
before you begin pairing them.

36. More Examples:
Hydrogen: ↑
1s1 H
1 e- 1s
Lithium: ↑↓ ↑
1s2 2s1 Li
3 e- 1s 2s
Sodium:
↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑
11 e- 1s 2s 2p 3s
1s2 2s2 2p6 3s1 Na

37. The End