1. Electrons and Quantum Mechanics
Unit 5

2. Electrons
Rutherford described the dense center of the atom called the nucleus.
But the Electrons spin around the outside of that nucleus.
Provide the chemical properties of the atoms.
Responsible for color and reactivity.

3. Energy Energy is transmitted from one place to another.
Light carries this energy.
Converted into heat.
Light is called Electromagnetic Radiation.

4. Electromagnetic Spectrum
Radio
Infrared
Visible Light
ROY G BIV
Ultraviolet
X Rays
Gamma Rays

5. Light Light travels as a wave. Wave Properties
Wavelength (λ) = distance between two waves (m)
Frequency (f) = number of peaks per second (Hz)
Speed of Light (c) = how fast light moves.

6. Light c= ƒλ Light Equation Speed of light is a constant = 3 x 108 m/s
Nothing travel faster than the speed of light!
Maybe?!?!?!?!?!?!?!?!?

7. Light The Dual Nature of Light
Light carries energy through space like a wave.
Light also behaves like a particle?!?
A beam of light is made of tiny packets of energy called PHOTONS!
Which travel in waves!?!

8. Light E = hƒ The Energy of a photon depends on its frequency.
So is the color of light!!!
E = hƒ
ELECTRONS are like photons!
Act as waves and particles.
Orbit the nucleus in a wave-like motion.

9. Blackbody Radiation
Rutherford could never explain why objects change colors when they are heated.
As the object heats, it must give off electrons of certain frequencies and energies.

10. Photoelectric Effect
Similarly, light on a metal object can knock off electrons.
Shine different colors on a metal.
Measure the number of electrons knocked off.
Found that no electrons were knocked off below a certain frequency.

11. The Bohr Model
Proposed the electrons orbit the nucleus with fixed energies.
Called Energy Levels
Much like the rungs of a ladder.
Quantum describes the amount of energy required to move an electron from one level to another.

12. The Bohr Model Ground State Excited State
Lowest possible energy of an electron.
Normal location
Excited State
If electron absorbs energy, it moves up an energy level (absorption)
If an electron gives off energy, it moves down an energy level (emission).

13. The Bohr Model

14. Atomic Spectra Hydrogen Atom Line Emission Spectrum
Expected continuous spectrum of light
But only specific frequencies were given off.
Red (656.6 nm)
Blue-green (486.1 nm)
Violet (434.1 nm)
Violet (419.2 nm)

15. Atomic Spectra Shine a light on an Atom Atomic Spectra
When atoms absorb energy, electrons move to higher energy levels.
When atoms release the energy, electrons return to the lower energy level.
Atomic Spectra
Frequencies of light emitted by a certain element.
No two elements have the same spectrum.

16. Flame Tests
Because no two atoms produce the same spectrum, elements can be identified by the colors they emit.
Spectral Analysis uses this properties to identify elements.

17. E = hf Quantum Mechanics Max Planck (1900) Albert Einstein (1905)
Founder of Quantum Mechanics
E = hf
Albert Einstein (1905)
Wave-Particle Duality
Electrons are small particles that move like waves.

18. mv/λ = h Quantum Mechanics Neils Bohr (1922) Louis de Brogelie (1923)
Electrons orbit in distinct energy levels.
Louis de Brogelie (1923)
Wave Mechanics says that ALL MATTER behaves like waves.
mv/λ = h

19. Quantum Mechanics Werner Heisenberg (1927) Erwin Schrödinger (1930)
Principle of Indeterminacy
You can’t know both the position and the velocity of an electron.
Erwin Schrödinger (1930)
Used wave mechanics to show the PROBABLE location of an electron.
Electrons exist in 3D clouds of probability!!!

21. Quantum Mechanical Model
Uses Schrodinger’s equation to predict the probable location of an electron.
Determines the energies an electron is allowed to have.
Determines how likely it is to find the electron in various locations around the nucleus.

22. Quantum Numbers Describes the location and behavior of an electron
Like an electron’s address
No two electrons can have the same quantum numbers.
Four Numbers

23. Quantum Numbers Principle (1st) Quantum Number (n) The Energy Level
Describes the size of the cloud and the distance of the cloud from the nucleus.
Shows the number of electrons
n = 1 = 2 electrons
n = 2 = 8 e-
n = 3 = 18 e-
n = 4 = 32 e-

25. Quantum Numbers 2nd Quantum Number (l) s = spherical p = peanut-shaped
Each energy level has sublevels.
The number of sublevels equals n.
Sublevels are called:
s = spherical
p = peanut-shaped
d = daisy-shaped
f = unknown?

27. Quantum Numbers s = 1 orbital p = 3 orbitals d = 5 orbitals
3rd Quantum Number (ml)
Divides sublevels into orbitals.
Tells the shape the electron moves in.
Number of orbitals = n2
Examples
s = 1 orbital
p = 3 orbitals
d = 5 orbitals
f = 7 orbitals

28. Quantum Numbers 4th Quantum Number (ms) Describes the electron’s spin.
Only two electrons fit in an orbital.
Their charges repel causing them to spin in opposite directions (+½ or –½)
Use up and down arrows.

30. Quantum Numbers Pauli Exclusion Principle Hund’s Rule
No two electrons can have the same set of 4 quantum numbers.
The electrons repel each other.
Hund’s Rule
Every orbital must get one electron before doubling up.

31. Quantum Numbers 1s 2s 2p 3s 3p 3d 4s 4p 4d 4f 5s 5p 5d 5f
Diagonal Rule
Electrons fill orbitals in predictable patterns
Some People Do Forget
Electrons dill the lowest energy level possible.
2s 2p
3s 3p 3d
4s 4p 4d 4f
5s 5p 5d 5f

32. Orbital Notation
Draw out the locations of each electron in an atom with arrows.

33. Electron Configuration
Write out the configurations of electrons using superscripts.
Examples:
H = 1s1
He = 1s2

34. Electron Configurations
Noble Gas Shorthand
Write the Noble Gas just before the element.
Add the remainder of the configuration.

35. Lewis Dot Diagrams
A way to show the number and position of the valence electrons.
Outermost energy level
Look at the column number to get this number.
Use the chemical symbol and number of valence electrons.
All four sides must have a dot before you double up.
p orbitals
s orbital X p1 s p2 p3