1. 5.2 QUANTUM THEORY & ATOM

2. BOHR MODEL (1913): Proposed energy states for hydrogen atom:
Ground state: lowest energy state
n=1
2nd orbit n=2
3rd orbit n=3
etc.

3. Louis de Broglie (1924): Accounted for the fixed energy levels in Bohr’s model.
All moving particles have wave characteristics.
May be too small to be detected.
Ex. car
Equation:  = h mv
m=mass
v=velocity
Wave/particle duality of
matter.

4. EXCITED ELECTRON: Energy added from outside source causes electron to move to a higher level.
Ex. Move from n=1 to n=3
When electron returns to lower state, energy is released:
E = Ehigher orbit – Elower orbit
= Ephoton = h
Energy released gives off a photon corresponding to its frequency.
Shown as spectral lines
(Flame Test)

5. ELECTRON – Goes to higher energy level when absorb energy
ELECTRON – Goes to higher energy level when absorb energy. Releases energy when returns to ground state.

7. HEISENBURG UNCERTAINTY PRINCIPLE: Impossible to know precisely both velocity & position of a particle at the same time.
When take a measurement, the object is disturbed.
This effect is significant if the object is very small (ex. electron).
To find position of an electron, need photon of light.
Photon “bumps” into electron, changing its position.

8. SCHRODINGER WAVE EQUATION: Very complex
Further developed de Broglie’s wave-particle theory.

9. QUIZ 1) The ___ Model proposed energy states for the hydrogen atom.
2) De Broglie stated that all moving particles have ____ characteristics.
3) When energy is added, the atom becomes ___. Energy is ___ when the atom returns to the lower state.
4) Heisenburg’s ____ Principle states that it is impossible to know both the ___ and ___ of a small particle at the same time.
5) Schrodinger’s ____ equation further developed de Broglie’s wave particle theory.

10. ATOMIC ORBITAL: Describes the electron’s probable location
3-D region around nucleus (“fuzzy cloud”)
Quantum #’s – describe each electron in atom (each is unique)
1) Principle Quantum Numbers (n):
Called Principle Energy Levels
Indicate relative sizes & energies of atomic orbitals.
n = 1 (ground state)
n values of 1 – 7 (7 energy levels)

11. PRINCIPLE ENERGY LEVELS

12. PRINCIPAL ENERGY LEVELS SUBDIVIDED: Energy Sublevels
Labeled by shape: s p d f
s – spherical (1 pattern or orbital)
p – dumbbell (on x, y, z axis)
d – complex (5 patterns)
f - v. complex (7 patterns)
Orbitals represent orientation in space
or patterns

13. s, p, d, f ORBITALS
f orbitals:

14. ENERGY LEVELS & SUBLEVELS 3rd
n= 1 s s
2 s p s 2p
3 s p d s 3p 3d
s p d f s 4p 4d 4f
s=1 orbital f orbital
p=3 orbitals
d=5 orbitals
f=7 orbitals

15. # ORBITALS = n2 n = energy level
At: n= 1: 1 orbital (s)
n= 2: 4 orbitals (s, px,py,pz)
n= 3: 9 orbitals (s, px,py,pz, 5 “d”)
Each orbital holds 2 electrons:
(electron can spin clockwise or counterclockwise)
Thus # electrons at an orbital = 2 n2
Ex. At n=1, holds 2 electrons
n=2, holds 8 electrons

16. Orbitals fill from low to high energy:

17. ENERGY OVERLAP: “d” & “f” have high energies!

19. QUIZ
1) The lowest energy state for the electron is called the ____ state. An electron is ___ when it moves to a higher energy level.
2) The electron’s probable location is called the atomic ____. This is the 3-D region around the ___.
3) The 4 types are s, __, __, and __.
4) “s” orbitals have a ___ shape. “p” are ___.
5) Orbitals fill from ___ to ___ energy Overlap is seen for d and __ orbitals.
6) How many orbitals in n=4?
How many electrons in n=4?

20. THE END