1. “Electrons in the Atom”
AP Chemistry

2. Different elements give different colors.

3. The Nature of Light
Electromagnetic radiation, including light, is described using wave terms.
Crest
origin
Amplitude
Trough
Wavelength (l)

4. Frequency and Wavelength
Frequency (n) is the number of waves that pass a point in a given amount of time.
Hertz (Hz) are the units and 1 Hz is one wavelength/second.
Light has a speed of 3.00 x 108 m/s (c).
For light, c = l n, this also applies to all electromagnetic radiation.
l and n are inversely proportional.

5. The Electromagnetic Spectrum
Visible light is only a very small part of the electromagnetic spectrum.
The highest frequency g rays down to the lowest frequency radio waves range from 1022 Hz to 104 Hz. Visible light is in the 1015 Hz range, about in the middle.
RED
ORANGE
YELLOW
GREEN
BLUE
INDIGO
VIOLET

6. Wavelengths for some EMR

7. Light can be a particle also.
Max Planck explained a “problem” that scientists had with glowing hot objects by suggesting that light could be a particle.
He suggested that matter could lose or gain energy in small packets of energy called “quanta”. A quantum is the smallest amount of energy that an atom can lose or gain.

8. Planck’s Equation
By applying a constant (Planck’s constant) Planck found the relationship between the energy of a photon (a quantum of light) was:
E = hn , where h is Planck’s constant and E is the energy measured in Joules.
h = x 10-34J.s

9. Rainbows are “continuous spectra”.
The refraction and reflection of sunlight by water droplets produces light in the full or continuous spectrum.

10. Atomic Emission Spectra
Because energy must be released only in certain size packages, the light given off by exciting gases can be resolved into a series of lines with only certain colors of light. Each acts like a “fingerprint”.

11. Different gases produce different lights.

12. The Bohr Model
Niels Bohr attempted to “fix” Rutherford’s model by having the electrons move in energy levels around the nucleus. Each energy level could only hold a certain number of electrons.
The larger the energy level, the further from the nucleus it was.

13. Bohr Model of the Atom 1 2 3 4 e-
The numbers represent the principal quantum numbers.

14. Ground State Electrons
When an electron is in the lowest energy level possible it is said to be in the ground state.
When an electron doesn’t occupy the lowest energy level possible due to outside energy it is said to be in the excited state.

15. Changing Energy Levels
When ever an electron moves to a higher or lower energy level an energy change is required.
If the right amount of energy is added, the electron moves up. To move down, a certain amount of energy must be released.

16. The Quantum Mechanical Model
Louis De Broglie combined Planck’s equation E=hn and
Einstein’s E=mc2 to produce an equation that explains why the electrons can only occupy certain energy levels.
l = h/mn

17. Heisenberg
Werner Heisenberg proposed that we could never tell the position and momentum (speed) of an electron at the same time.
This uncertainty is the “Heisenberg Uncertainty Principle”.

18. Schrodinger
Erwin Schrodinger derived a formula that described the space that an electron of a certain energy would occupy.
Instead of an orbit, he described an “electron cloud” where you would have the greatest probability finding the electron.

19. Four Quantum numbers for electrons in an atom.
The principal quantum number (n) describes the size and energy of the electron orbital.
Sublevels (l) describe the shape of orbitals. The number of sublevels = n
The direction (m) describes orientation of the sublevels.
Spin (s) refers to how an electron spins.

21. Electron configuration
The lowest energy levels are filled first, this is the “Aufbau Principle”.
Once an energy level is full electrons can then fill the next highest energy level.
“Pauli Exclusion Principle” states that no 2 electrons in the same atom can have the same 4 quantum numbers.

22. Electron Sublevels s sublevels can have up to 2 e- and is spherical.
p sublevels can have up to 6 e- and looks like 3 “dumbbells” in the x, y and z axis
d sublevels can have up to 10 e- and the shape is more complicated
f sublevels can have up to 14 e-.

23. p and d orbitals

24. Electron orbitals Electron orbitals can contain 0, 1 or 2 e-.
s has 1 orbital, p has 3, d has 5, and f has 7.
We show orbitals as a box with electrons represented as arrows. Spin is represented by the direction of the arrows.

25. Hund’s Rule
Electrons will seek an unoccupied orbital within a sublevel before it will pair up with another electron in an orbital.
P sublevel
P sublevel NO
YES

27. Orbital Filling
The e- must fill the lowest energy level first and only 2 e- can be placed per orbital.
The first e- goes into the lowest principle energy level, n=1.
Level 1 only has 1 sublevel, s
An s sublevel can only have 1 orbital

28. Orbital filling continued
When the first energy level is filled, the e- can occupy the second energy level, n=2.
Level 2 can have 2 sublevels, s and p
After sublevel s is filled then p is filled according to Hund’s rule.

29. The Flow Chart of Science
Electron filling follows the following pattern:
1s 2sp 3sp 4s 3d 4p 5s 4d 5p 6s
4f 5d 6p 7s 5f 6d 7p … see p135.
Notice that each p is followed by the next higher level’s s.

30. Electron configuration
An easy way to represent the e- in an atom is through electron configuration.
2 p 5
The principle energy level is 2 and there are 5 e- in the p sublevel.

31. Example:
The electron configuration for the sulfur (Z= 16) atom is: 1s2 2s2p6 3s2p4
This says that the total number of e- is 16 and the highest energy level is n=3.

32. Valence Electrons
The chemical characteristics of atoms are based upon the number of e- that are in the outer energy levels. These are called valance electrons. They are only s and p sublevel e-.
For the sulfur atom, 1s2 2s2p6 3s2p4, there are 6 valence e-.