1. The whole range is called a continuous spectrum
EM – ElectroMagnetic
Spectrum
The whole range is called a continuous spectrum
Low energy
High energy
Radio
Waves
Radiowaves
Microwaves
Micro
Waves
Infrared .
Infra
Red
Ultra-
Violet
Ultra-violet
X-Rays X
rays
Gamma
Rays
GammaRays
Low Frequency
High Frequency
Longer Wavelength
Shorter Wavelength
Visible Light
continuous spectrum
R O Y G B I V

2. Speed of light: Frequency and wavelength
c = 3.00 x 108 m/s
c = l n
l and n are inversely related
As one goes up the other
goes down.
Different frequencies (wavelengths) of light have different colors.

3. E = h The energy of a photon is proportional to its frequency.
Light is Radiant Energy made of photons
Photons are discrete packets of radiant energy
The energy of a photon is proportional to its frequency.
E = h
E: energy (J, joules)
h: Planck’s constant (6.63  J·s)
: frequency (Hz)

4. Bohr Model of Hydrogen Examined the spectrum of H atom
Hypothesized the electrons are in “energy levels” around the nucleus
Calculated the radii of the first few orbits

5. Spectroscopic analysis of the hydrogen spectrum…
…produces a “bright line” spectrum

6. Electron transitions involve jumps of definite amounts of energy.

7. The Quantum Mechanical Model
The nucleus is found inside a blurry “electron cloud”
A area where there is a high probability of finding an electron.

8. *No more than 2 electrons in each orbital*
Electron Configuration
Principal Quantum Number (n) = the energy level of the electron. (max # e- = 2n2 )
Sublevels- like theater seats arranged in sections (s,p,d,f) (# of sublevels = n )
Within each energy level, the complex math of Schrodinger’s equation describes several shapes.
These are called atomic orbitals - regions where there is a high probability of finding an electron
*No more than 2 electrons in each orbital*

9. Summary # of orbitals Max electrons Starts at energy level sublevel s1 2 1 p 3 6 2 d 5 10 3 7 14 4 f

10. Relative sizes of the spherical 1s, 2s, and 3s orbitals of hydrogen.

11. Diagram of principal energy levels 1 and 2.

12. The shapes and labels of the five 3d orbitals.

13. Electron Configurations
The way electrons are arranged in atoms.
Three main “guiding” principles:
Aufbau principle- electrons enter the lowest energy first.

14. Hund’s Rule- When electrons occupy orbitals of equal energy (same sublevel ) they don’t pair up until they have to.
- one electron enters each orbital
until all orbitals contain one electron
with spins parallel.

15. Pauli Exclusion Principle- at most 2 electrons per orbital
Why only 2 electrons per orbital?
like charges repel each other, but…
electrons spin on axis,
clockwise or counterclockwise,
creating magnetic polarization ( ↑↓ )

16. Writing Electron Configurations
(Letter - number designation)
# of e- H
1s1
Sublevel
Principle Energy Level

17. Orbitals being filled for elements in various parts of the periodic table.

18. Orbital Diagrams ↑ ↓ ↑ ↓ ↑ ↑ ↑ N 1s 2s 2p 2 2 3
Show each orbital as a circle, box, or line.
Show each electron as an arrow, ↑ or ↓ spin.
Label electron configuration (letter –number). ↑ ↓ ↑ ↓ ↑ ↑ ↑ N 7 1s 2s 2p 2 2 3

19. P Increasing energy 15 7p 6d 5f 7s 6p 5d 6s 4f 5p 4d 5s 4p 3d 4s 3p 3s

20. Do Problems pgs 330 -332 # 38, 40, 60, 77 Remember…
When we write electron configurations, we are writing the lowest energy (ground state) configuration.
Do Problems pgs
# 38, 40, 60, 77