1. Electrons and Energy
Chapter 5

2. Rutherford’s Failure
Rutherford’s model
could not describe
behavior of atoms
Think About It:
Why do certain
atoms react
while others
don’t? Why do
some fireworks
glow red and
others green?

3. Bohr’s Atom
Hypothesis: each
electron exists in a
certain “orbit” or
“energy level”
Electrons can never
exist in between
Higher energy levels
are further from the
nucleus

4. Quanta
A quantum (pl. quanta) is the amount of energy
required for an electron to move up an energy
level

5. Quantum Mechanical Model
Also called “electron
cloud model”
Mathematically-based

6. Quantum Mechanical Model
Determines allowed
electron energies and
how likely it is to find
an electron in various
locations around the
nucleus
Dense cloud = high
probability

7. Atomic Orbitals
Region of space in
which there is a high
probability of finding
an electron
Why It’s Important:
The exact location
of each electron
controls the atom’s
properties!

8. “Address” of Electron
Each electron will be found
On a PRINCIPAL ENERGY LEVEL
In one of several SUBLEVELS
On an ORBITAL

9. Principal Quantum Numbers
Each principal energy
level is assigned a
principal quantum
energy number, n
n can be 1, 2, 3, or 4,
etc…

10. Sublevels
Each principal energy level
has specific sublevels
where an electron can be
found
# of sublevels corresponds
to principal quantum
number
 Level 1 has 1 sublevel
 Level 2 has 2 sublevels…

11. Sublevels
Sublevels are given
letter designations
(according to shape)
s
p
d
f

12. Why s, p,
d, and f?
The orbital names (s,
p, d, f, g, h,...) are
derived from the
characteristics of their
spectroscopic lines:
sharp, principal, diffuse
and fundamental, the
rest being named in
alphabetical order. For
mnemonic reasons,
some call them
spherical & peripheral.

13. Atomic Orbitals
s sublevel – 1 orbitals
p sublevel – 3
d sublevel – 5
f sublevel – 7

14. Atomic Orbitals
Each orbital can
hold one pair of
electrons
Each member of
the pair has an
opposite spin

15. How many totalHow many total
electrons in each?electrons in each?

16. 1st
Principal Energy Level
• 1 sublevel (s)
• 2 electrons in s orbital
• 2 total

17. 2nd
Principal Energy Level
• 2 sublevels (s and p)
• 2 electrons in s orbital
• 6 electrons in p orbitals
• 8 total (10 combined with 1st
)

18. 3rd
Principal Energy Level
• 3 sublevels (s, p and d)
• 2 electrons in s orbital
• 6 electrons in p orbitals
• 10 electrons in d orbitals
• 18 total (28 combined with 1st
and 2nd
)

19. 4th
Principal Energy Level
• 4 sublevels (s, p, d and f)
• 2 electrons in s orbital
• 6 electrons in p orbitals
• 10 electrons in d orbitals
• 14 electrons in f orbitals
• 32 total (60 combined with 1st
, 2nd
, and 3rd
)

20. Electron Arrangement
Blue Book: 13.2
Red Book: 5.2

21. Electron Configuration
Way that
electrons are
arranged in
various orbitals
Arrangement
governed by
three rules

22. Aufbau Principle
Electrons first
occupy
sublevels/orbitals of
lowest energy
Sublevels from some
principal energy
levels overlap others

23. Aufbau Diagrams
Display the order in which
sublevels/orbitals are filled
(to achieve the most
stable configuration)

24. Aufbau Principle

26. Pauli Exclusion Principle
Each orbital holds, at most, two electrons.
Two occupy the same orbital, two
electrons must have opposite spins.
This is written or

27. Hund’s Rule
In each sublevel, one
electron enters each
orbital until each has
one with the same spin
direction
Not until all orbitals
have one electron can
any have two

28. Orbital Filling Diagrams
Show
which
orbitals are
filled by
electrons

29. Writing Electron
Configurations

30. Electron Configurations
Sublevels are filled in the
following order:
1s, 2s, 2p, 3s, 3p,
4s, 3d, 4p, 5s

31. Electron Configurations
Contain 3 items:
Number represents energy level
Letter represents sublevel
Superscript represents number of electrons
1s2

32. Electron Configurations
Additional segments needed for each
sublevel
1s2
2s2

33. Electron Configurations
1s2
2s2
2p6

34. Electron Configurations
of Specific Elements
In a neutral element, the number of electrons is
equal to the number of protons.

35. Practice Problems
Please write orbital
diagrams for the
following elements:
Phosphorus
Iron
Please write electron
configurations for the
following elements:
Sodium
Iron
Platinum

36. Electron Configurations
Hydrogen
Oxygen
Neon
Silicon
Calcium
Superscripts add up
to the total number
of electrons!