1. Chapter 6 Modern Atomic Theory

2. Review… Dalton Thomson Rutherford
Model doesn’t explain how the negative electron can stay in orbit and not be attracted to the positive proton

3. Electromagnetic Radiation
Light travels in
Light is a form of
Form of energy that exhibits

4. Electromagnetic Radiation
All waves can be described in 3 ways:
Amplitude –
Wavelength (l):
Frequency (n):

5. Electromagnetic Radiation
Speed of light in air: Electromagnetic radiation moves through a vacuum at speed of
Since light moves at constant speed there is a relationship between wavelength and frequency:
Wavelength and frequency are inversely proportional
C = speed of light
L = wavelength
V = frequency

6. Electromagnetic Spectrum
Visible spectrum – white light (what we can see)
Wavelengths of visible spectrum are 350 nm to 700 nm

7. Photoelectric Effect The emission of Photon =
Albert Einstein (1905) used Planck’s equation to explain this phenomenon;
proposed that light consists of
Photon =

8. Photoelectric Effect
He (Einstein) explained that the photoelectric effect would not occur if the frequency and therefore
Analogy:
70 cents placed in soda machine: no soda
30 cents more and you will get your soda

9. Niels Henrik David Bohr
Physicist
Worked with Rutherford
1912
Studying line spectra
of hydrogen

10. Niels Henrik David Bohr
1913 – proposed new atomic structure
Electrons exist in
Electrons

11. The Bohr Atom Nucleus with Electrons move in
When an electron moves from one state to another the energy lost or gained is in
Each line in a spectrum is produced when an electron moves from

12. The Bohr Atom
Model didn’t seem to work with atoms with more than one electron
Did not explain chemical behavior of the atoms
Safeco field example

13. Now…
Light can be described as
What does this mean for the atom???

14. Line Spectrum Elements in gaseous states give off colored light
High temperature or high voltage
Always the same
Each element is unique
Spectra
Also known as the “atomic emission spectrum of an element”
Visible light or electromagnetic radiation is emitted when an atom passes from a state of higher potential energy to a state of lower potential energy

15. Line Spectrum
Ground state
Excited state

16. Line Spectrum
Electron
Color of light emitted depends on

17. Line Spectrum Each band of color is produced by light of a different
Each particular wavelength has a definite
Each line must therefore be produced by emission of photons with

18. Line Spectrum Whenever an excited electron
The energy of this photon is equal to the difference

19. Wave Matters… Louis de Broglie (1924)
Proposed that electrons might have a
Used observations of normal wave activity
Safeco field example

20. Problems… Wave theory does not explain Heated iron gives off heat
1st red glow yellow glow white glow
How elements such as barium and strontium give rise to green and red colors when heated

21. Energy is released in Beginnings… Max Planck (1858-1947)
Proposed that there is a fundamental restriction on the amounts of energy that an object emits or absorbs, and he called each of these pieces of energy
Energy is released in

22. Beginnings
A quantum is a finite quantity of energy that can be gained or lost by an atom
This constant, h, is the same for all electromagnetic radiation

23. En = (-RH)(1/n2) Bohr’s Equation Where RH = 2.18 x 10-18J
And n = principal quantum number, 1 to infinity

24. Jumping electrons…
If an electron moves from one energy level to another, the change in energy can be determined by the following equation:
E = Ef – Ei = hν
Or simply: E = hv
Where h=6.626 x J s

25. Then… by substitution…( E RH ( 1 1 - =
ν =
ni2
nf2 h h

26. h λ = mν Finally… Matter waves All moving particles
Some is apparent, some not.
De Broglie’s equation h λ = mν

27. Smart guy… Erwin Schrodinger (1926)
Used mathematical understanding of wave behavior – devised an equation that treated electrons moving around nuclei as waves
Quantum Theory
Safeco field example

28. Uncertainty principle
Heisenberg:

29. Quantum Theory
Describes mathematically the wave properties of electrons and other very small particles
Applies to all elements (not just H)

30. Quantum Numbers Numbers that specify the Principle Quantum Number:
Symbolized by n,

31. Energy Levels of Electrons
Principle energy levels
Designated by letter n
Corresponds to the
Each level divided into sublevels
1st energy level has
2nd energy level has
Etc.

32. Orbitals Electrons don’t Orbital: region in space where
Each orbital sublevel can hold
Safeco field example

33. Orbitals Each sublevel (orbital) has a specific shape
Safeco field example

34. Quantum Numbers Orbital Quantum Number:
Indicates the shape of an orbital
(subshell or sublevels)
s, p, d, f
Principal Quantum # Orbital Quantum # 1 2 3 4

35. Quantum Numbers Magnetic Quantum Number: Indicates the
Orbital position with respect to

36. Orbitron For a full view of the different orbital shapes, visit

37. Orbitals Pauli exclusion principle: Electrons can only spin Shown with
Safeco field example

38. Rules for Orbital Filling
Pauli’s Exclusion Rule
No two electrons have
Hund’s Rule
Electrons will remain
1s s 2p s 3p

39. Rules for Orbital Filling
Diagonal Rule
The order of filling once the d & f sublevels are being filled
Due to energy levels

40. Rules for Orbital Filling
Safeco field example

41. Application of Quantum Numbers
Several ways of writing the address or location of an electron
Lowest energy levels are filled first
Electron Configuration:
12C:
32S:

42. Application of Quantum Numbers
Orbital filling electron diagram: using Hund’s rule and the diagonal rule write out the location of all electrons
See examples on whiteboard