1. Electrons In Atoms
Where are they?

2. Development of Atomic Models
Plum Pudding Model (1897)
J.J. Thomson
Electrons scattered in a “sea” of positive charges

3. Development of Atomic Models
Rutherford’s Model (1911)
Discovered nucleus (disproves Plum Pudding)
Electrons orbit nucleus like planets around the sun
Cannot explain many of the properties of atoms

4. Development of Atomic Models
Bohr Model (1913)
Electrons move around nucleus in circular orbits at specific allowed distances
These distances relate to allowable energy levels
Energy levels – fixed energies an e- can have
Quantum of energy – energy needed to move an electron from one E level to another

5. Development of Atomic Models
More Bohr Model
Electrons can gain or lose energy
Ground state – lowest energy level available
Excited state – higher energy level
Absorb energy (gain E)
Go from lower to higher E levels
Emit energy (lose E)
drop from higher to lower E levels
Give off E in the form of radiation (quanta of light)

6. Development of Atomic Models
More Bohr Model
Energy levels get closer together as they get farther from the nucleus
Problem: Works well with the hydrogen atom but not much else

7. Development of Atomic Models
Quantum Mechanical Model (1926)
Modern description of electrons in atoms
Cloud model or Quantum Theory
Schrodinger – developed mathematical equation to predict atomic behavior
Electrons NOT in exact path
Heisenberg Uncertainty Principle
Impossible to know both location and energy of an electron
Can measure one or the other – NOT both
Exact motion of electron unknown

8. Development of Atomic Models
More Quantum Mechanical Model
Determines allowed energies an electron can have and how likely it is to find the electron in various locations around the nucleus
These locations are called principal energy levels
Within these energy levels are sublevels
Sublevels are subdivided into atomic orbitals

9. Development of Atomic Models
More Quantum Theory
Atomic Orbitals
Region of space in which there is a high probability of finding an electron
4 types of orbitals – s, p, d, and f
Different orbitals have different shapes
Each orbital can hold up to 2 electrons
2 electrons in same orbital must have opposite spins (+1/2 and -1/2)

10. s and p orbitals

11. d sublevels

12. f orbitals

13. Quantum Numbers 4 numbers used to describe electron location
Principal Energy Level (Principal Quantum Level)
n = 1, 2, 3…
Energy Sublevel Number
specifies s, p, d, or f sublevel
l = 0 to n-1
l = 0 s sublevel
l = 1 p sublevel
l = 2 d sublevel
l = 3 f sublevel

14. Quantum Numbers ms = +1/2 or -1/2 Orbital quantum number (m)
m = -l to +l
Specifies which orbital within a sublevel the electron is located
Within sublevels, orbitals differ only in spatial orientation, not energy
Spin quantum number (ms)
ms = +1/2 or -1/2
1st electron in orbital has + spin

15. Energy Levels, Sublevels, and Orbitals
Principal Energy
Level
Number of
Sublevels
Types of sublevels 1
1s (1 orbital) 2
2s (1 orbital), 2p (3 orbitals) 3
3s (1 orbital, 3p (3 orbitals)
3d (5 orbitals) 4
4s (1 orbital), 4p (3 orbitals)
4d (5 orbitals), 4f (7 orbitals)
*** Remember: Each orbital can hold 2 electrons ***

16. Orbitals and Electrons
Energy level, n
Sublevel
(Orbitals)
Max # Electrons / sublevel
Max # Electrons / energy level 1 s 2
s p 8 3
s p d 18 4
s p d f 32
s sublevel – 1 orbital, 2 electrons p sublevel – 3 orbitals, 6 electrons
d sublevel – 5 orbitals, 10 electrons f sublevel – 7 orbitals, 14 electrons
Maximum # electrons / energy level = 2n2
where n = energy level

17. Electron Configuration
The way in which electrons are arranged in various orbitals around the nucleus of an atom
Aufbau Principle
Electrons occupy the
orbitals of lowest energy first

18. Electron Configuration
Pauli Exclusion Principle
An atomic orbital may describe at most two electrons
Opposite spins
Boxes represent orbitals and arrows represent electrons
3s sublevel with 1 electron 
4s sublevel with 2 electrons 

19. Electron Configuration
Hund’s Rule
Electrons occupy orbitals of the same energy in a way that makes the number of electrons with the same spin as large as possible
Orbitals of equal energy each get 1 electron before any pair up
2p sublevel

20. Electron Configuration
Diagonal rule
s holds 2
p holds 6
d holds 10
f holds 14

21. Atomic Structure Practice
3a) 1s22s22p3 Nitrogen ( 7 electrons)
b) 1s22s2 Beryllium (4 electrons)
c) 1s22s22p63s23p3 Phosphorus (15 electrons)
d) 1s22s22p63s23p63d54s2 Manganese (25 electrons)
e) Potassium (19) f) Zirconium (40)
g) Promethium (61) h) Selenium (34)

22. Atomic Structure Practice
4a) Cu0 (29 e-) 1s22s22p63s23p64s23d9
Cu+ (28 e-) 1s22s22p63s23p64s23d8
Cu2+ (27 e-) 1s22s22p63s23p64s23d7
4b) Al0 (13 e-) 1s22s22p63s23p1
Al3+ (10 e-) 1s22s22p6

23. Electron Configuration
Short Cut Method
Rare Gas Configuration, Noble Gas Configuration, or Inert Gas Configuration (Either name OK)
Relate back to the previous rare gas
Put that element in [ ]
Start at s sublevel using whatever period the element is in
Nickel 1s22s22p63s23p64s23d8 or [Ar] 4s23d8