1. Electronic Structure and Light
Unit 5

2. The Wave Nature of Light
electrons, electrons, electrons
Knowing the arrangement, number of electrons, and energy of the electrons (electronic structure) in an atom is the key to understanding the physical and chemical properties of an element.
What does the study of light have to do with understanding electronic structure?

3. Electromagnetic Radiation
Radiant energy
Form of energy that has wave-like characteristics.
Speed in a vacuum is 3.00 x 108 m/s
Hi Mr. Lamberth :]
I hope you had fun with that girl you brought to the soccer game

4. Electromagnetic Spectrum

5. Electromagnetic Spectrum

6. Light c = speed of light = 3.00 x 108 m/s
Wavelength () - distance between troughs
Uses the units m(SI unit) or nm
Frequency (f) - the number of complete wavelengths that pass a given point each second.
Uses the units cycles per second; s -1 = /s (SI unit) or hertz (Hz) Hz = 1 s -1
Related by c = f 
Frequency and wavelength are _________ proportional.

8. Homework
HS: Complete Assignments 5-6 in your packet.

9. A revolution is born
Physicists wanted to understand the relationship between the temperature and the intensity and wavelengths of the emitted radiation.
Prevailing laws could not explain three phenomena:
Blackbody Radiation
Photoelectric Effect
Emisssion Spectrum

10. Max Planck ( )
Explained how radiation is emitted by hot objects (black-body radiation)
Assumed energy can be released (or absorbed) by atoms or molecules only in discrete quantities
Quantum (plural quanta) – discrete chunk of energy, a fixed amount.
E = hf
Planck’s constant, h = 6.63 x J s
He could not explain why, but it worked.

11. Albert Einstein ( )
Explained photoelectric effect (When electrons are ejected off of the surface of a metal with light with a certain minimum frequency shines on it.
Extended Planck’s theory
Photons - packets of energy that behave like particles; particles of light
If photons have enough energy when they strike the metal, they will pass that energy to the electrons causing the electrons to fly off!
Energy of photon = E = hf (radiant energy itself is quantized)

12. Is light composed of waves or particles?
What do you think?

13. Spectrum Continuous Spectrum - Continuous range of colors
Line (Emission) Spectrum – A spectrum containing radiation of specific wavelengths emitted by a substance.
Absorption Spectrum – A spectrum containing radiation of all wavelengths except those absorbed by a substance.
Who was the first to shine light through a prism and study it?

15. Hydrogen Spectra
Absorption spectrum
Line spectrum

16. Emission Spectrum of Elements

17. Bohr’s Model Three postulates:
Assumed that electrons move in circular orbits around the nucleus. These orbits correspond to certain definite amounts of energy.
An electron in a permitted orbit has a specific energy and is in an allowed energy state. It will not spiral into the nucleus.
Energy is emitted or absorbed by an electron as it changes from one energy state to another . This energy exists as a photon.
Ground State: Electrons are as close to the nucleus as they can be; they are in the lowest energy level
Excited State: Electrons are not as close to the nucleus as they can be; they are in a higher energy level

19. Louis de Broglie
Dual nature of the electron - suggests that if light can behave like a stream of particles then electrons may possess wave properties.
Soon after this was published it was experimentally demonstrated.

20. Werner Heisenberg
A wave extends into space, therefore its exact location can not be found.
Also because photons are used to detect electrons and the energy of each of these particles is similar, then any attempt to locate an electron with a photon will knock the elctron off its course.

21. Heisenberg Uncertainty Principle
It is impossible to know the momentum and position of a particle with certainty.
Because of this we know the electron does not orbit the nucleus in a well defined path (as Bohr thought)

22. Quantum Mechanical Model
The Schrodinger Equation incorporates both the wave and particle behavior of electrons.
The location of an electron cannot be described so simply.
Launched a field of physics called quantum mechanics.

23. Quantum Mechanical Model
Quantum mechanics mathematically defines the region where the electron might be at a given time.
Electron Density – the probability that an electron will be found in a particular region of an atom.

24. Electron Density

25. Orbitals
Electrons do not travel around the nucleus, instead they exist in regions called orbitals.
Region in the atom where there is a high probability of finding an electron.
Each has a characteristic shape and energy.

26. Quantum Numbers Principal quantum number (n): 1,2,3, etc.
Determines the size of the orbital, distance from the nucleus, and the energy level.
Azimuthal quantum number (l ): can have values from 0 to n-1.
Determines shape of the orbital. n
Value for l 1 2
0, 1 3
0, 1, 2 4
0, 1, 2, 3 5
0, 1, 2, 3, 4
Type of sublevel l s p 1 d 2 f 3

27. Quantum Numbers Magnetic quantum number (ml): values of l to -l
Orientation of the orbital (x, y, or z axis)
Subshell in orbital l
Possible values for ml 1
-1, 0, 1 2
-2, -1, 0, 1, 2 3
-3, -2, -1, 0, 1, 2, 3

28. Quantum Numbers Spin magnetic quantum number (ms): +1/2 or -1/2
Electrons move in opposite directions within subshell

29. Electron Configuration
H: 1s1
1 (the coefficient) – gives energy level
s – gives orbital shape
1s – gives subshell
Superscript – gives the number of electrons in that subshell

30. Electron Configuration
Pauli Exclusion Principle – No two electrons in an atom can have the same four quantum numbers. They will have different spins.

31. Electron Configuration
Hund’s Rule – the most stable arrangement of electrons in subshells is the one with the greatest number of parallel spins. Place electrons in each orbital one at a time, then go back and place two in them if neccesary.

32. Aufbau Principle
As protons are added one by one to the nucleus to build up the elements, electrons are similarly added to the atomic orbitals.
Electrons will fill lower energy subshells first then move on to higher energy subshells.
1s 2s 2p 3s 3p 4s 3d 4p etc…