1. Quantum Theory & the Atom

2. Historical Development of quantum theory
At the beginning of the 20th century light’s wave properties were well accepted but they theory was not able to account for everything.

3. The Problem of Black Body Radiation
Why does a piece of iron glow at different colours when heated?
First infra-red, then red, then yellow then white hot. What about the other colours in the spectrum?

4. Max Planck
In 1900 Max Planck proposed a solution. If the energy was dispersed in defined packets of energy (he named them quanta, quantum singular) this could account for the specific colours that represented specific frequencies/wavelengths of energy.
He defined these quanta as having the energy
E = h
E = energy
h = constant (now named Planck’s constant)
 = frequency of the light energy
This theory was not well accepted when published

5. Einstein to the rescue!
In 1905 Einstein solved a problem called the photoelectric effect using Planck’s idea of quanta
The photoelectric effect is the phenomenon in which different colours of light have different effectiveness in expelling electrons from a metal’s surface.
Einstein proposed that light consists of tiny packets of energy called photons whose effectiveness depends on their frequency.
Ie violet light (higher frequency photons) is more powerful that red light with a lower frequency photons
Einstein’s use of Planck’s idea popularized it

6. Dual Nature of Light
This led to a greater understanding of light as having properties of both waves and particles
Demonstrated in 1923 by Arthur Compton

7. Line Spectra Revisited
The Bohr model of the atom was informed by the line spectrum of hydrogen
The single electron in hydrogen must be in quantized orbits with different energies

8. Bohr Model Each orbit is a specific energy level
An electron that is excited or gains energy jumps to a higher energy level, when it drops a level it releases a photon with a specific energy (IE colour)
Quantum number (n) – the number given to the orbits
Lowest energy n=1 (closest to nucleus)
Ground state – element at its lowest energy level
Bohr’s model worked well for hydrogen but not for elements with more than one electron

9. Matter Waves
1924 Louis De Broglie proposed wave behavior for particles breaking down the barrier between matter and energy
Electrons move as waves
Thus redefining the idea of an electron orbit

10. Heisenberg’s Uncertainty Principle
1927 Werner Heisenberg proposes that it is impossible to know both the position and momentum of an electron simultaneously
Measurement of an electron’s position or movement alters its path in an unpredictable manner
This leads to the idea that one can only talk about probabilities for the location and movement of electrons
Atomic “Orbits” are calculated as probability
clouds not defined orbits

11. Quantum Mechanical Model
The energy of electrons is quantized
Electrons exhibit wavelike behavior
It is impossible to know the exact position and momentum of an electron at any given moment

12. Orbitals
Since it is inaccurate to speak of electron paths with precise orbits the term orbital is used instead
An Orbital is the region of probability around a nucleus for finding an electron with a particular energy
Calculated orbitals have the appearance of a cloud

13. Orbital shapes Orbitals are designated by the letters s, p, d, f
s orbitals are spherical
p orbitals are dumbbell shaped
The others are more complex

14. Orbitals & Energy
Principal Energy levels – main energy levels found in quantum model
n=1,2,3,…
Determined by principal quantum number (like Bohr’s model)
Unlike Bohr’s model they are further divided into sublevels
The number of sublevels is determined by n for that energy level

15. Notice that there is overlap between different energy levels

16. Orbitals and Energy Levels
Notice that each orbital type has a unique population in each energy level

17. Electron Spin
In addition to energy electrons have a property called spin
They can spin in either of two directions (clockwise or counterclockwise)
These spins result in magnetic poles pointing in opposite directions
Pauli Exclusion Principle (1925) – each orbital can hold a maximum of two electrons and they must have opposite spins
Therefore each level has a maximum number of electrons