1. Electrons in Atoms
Ch. 5 Notes

3. Light & Quantized Energy
Nuclear Model was incomplete
How are electrons arranged?
Why aren’t e- pulled into the nucleus?
Why do elements behave differently?
Early 1900s
Certain elements emit visible light when heated in a flame

4. Electromagnetic Radiation
A form of energy that exhibits wavelike behavior
Waves can be described…
Wavelength (l)
Frequency (n)
Amplitude

5. Wavelength – The shortest distance between equivalent points on a wave
Frequency – The # of waves that pass a given point in one second
Amplitude – Height from the origin to the crest/trough

6. Wave Diagram

7. All electromagnetic radiation travels at the speed of light (c).
c = 3.00 x 108 m/s
The velocity of a wave is equal to the wavelength times frequency
c = ln

8. Though the speed of electromagnetic radiation is always the same both the frequency and the wavelength can vary.
The electromagnetic spectrum encompasses all forms of electromagnetic radiation.
Only difference in portions are wavelength and frequency

9. Considering light as a wave does not explain everything
Different colors of light correspond to different frequencies and wavelengths.

10. Quantum – The minimum amount of energy that can be gained or lost.
1900 Max Planck trying to figure out why certain types of light are emitted from heated objects.
Discovered matter can only gain or lose energy in small specific amounts (quanta).
Quantum – The minimum amount of energy that can be gained or lost.

11. Equantum = hn
h – Planck’s constant
h = x 10-34
Photoelectric effect – photo e- are emitted from the surface of a metal when light with a certain frequency shines on its surface.
Wave model could not explain

12. 1905 Albert Einstein
Electromagnetic radiation has both wavelike and particlelike natures.
Photon – Particle of electromagnetic radiation with no mass that carries one quantum of energy
Ephoton = hn

13. Atomic emission Spectrum – the set of frequencies of electromagnetic radiation emitted by atoms of an element
Each element has a unique spectrum

14. Electron Configurations
The arrangement of electron’s in an atom
Tend to assume lowest possible energy (ground state)
Rules that tell how e- are arranged
Aufbau Principle
Pauli Exclusion Principle
Hund’s Rule

15. All orbitals of an energy sublevel are of equal energy
Aufbau Principle
Each electron must occupy lowest energy level orbital available
Fig p. 133
All orbitals of an energy sublevel are of equal energy
Energy sublevels within a principal energy level have different energies.
s < p < d < f
Orbitals from one sublevel and principal level can over lap with those from another principal level

16. . Pauli Exclusion Principle (Wolfgang Pauli) Hund’s Rule
Max of 2 electrons can occupy one orbital
Iff they have opposite spins
Hund’s Rule
Single electrons w/ the same spin must occupy each equal-energy orbital before opposite spin electrons can occupy the same orbitals

17. Electron-dot structure (G.N. Lewis)
Valence electrons
Electrons in the outermost orbitals
Highest principal energy level
Chiefest in determining chemical properties of an element
Electron-dot structure (G.N. Lewis)
Element’s symbol surrounded by dots
Symbol = nucleus and inner electrons
Dots = valence electrons