1. Unit 2
Atomic
Structure

2. Electromagnetic Spectrum
This is the spectrum of all radiation resulting from fluctuations of electric currents and vibrations of charged particles.
This radiation comes in many forms such as visible light, x-rays, gamma rays, UV light, and radio waves.
The spectra looks like a rainbow (ROY G BIV) that varies in wavelength and frequency.

3. Wavelength is the distance between a point in a wave and the corresponding point in the next wave (measured in nanometres)
Frequency is the number of wave vibrations per unit time (measured in hertz)
The red part of the spectrum has the longest wavelength while the violet part has the shortest wavelength
But since wavelength and frequency are inversely related, the red part of the spectrum have a lower frequency while the violet part has the highest frequency.

4. Some spectra examples:

5. The line spectra of elements are important in nature especially for astronauts trying to determine the elements of an unknown light source (stars, nebula, etc…)
Also, line spectra are important in the manufacturing of fireworks
Barium is yellow/green, Strontium is bright red, Calcium is orange/red, Sodium is bright yellow, Potassium is light purple, Copper is blue, and Lithium is light red.

6. Historical Development of the Quantum Mechanical Model
The quantum mechanical model is a model of the atom that shows the probable location of electrons.
Many scientists have worked on this theory.
Those who have made significant contributions in the past include Niels Bohr, Johannes Rydberg, Louis De Broglie, Erwin Schödinger, Max Planck, and Werner Heisenberg
All of these scientists have worked towards our present model of the atom.
You must create a timeline for the research done by the above mentioned scientists.

7. Electron Configurations
Electrons exist on orbits.
The lowest orbit level is called the ground state
As the electrons absorb more quantum energy, they get more excited and can jump to a higher orbit or energy level
There are different principal energy levels each with their own sublevels
Typically the principal levels correspond to the row number (except the transition elements)
The sublevels include s, p, d, and f.

8. In the s sublevel, only 2 electrons fit.
In the p sublevel, only 6 electrons fit.
In the d sublevel, only 10 electrons fit.
In the f sublevel, only 14 electrons fit.
There are rules for building electron configurations of neutral atoms in the ground state:
1) each electron is placed in the lowest energy level and sublevel
2) no more than two electrons can be paired up (in drawing)
3) when drawing, each sublevel must contain one electron before a second electron can be placed in that orbital

9. The electrons are placed in the sublevels according to the following diagram:
7s 7p
6s 6p 6d 6f
5s 5p 5d 5f
4s 4p 4d 4f
3s 3p 3d
2s 2p 1s

10. For Boron (5 electrons) the diagram would be:
Pictorially, we can draw an electron configuration using the following outline: 1s 2s 3s 2p 3p
For Boron (5 electrons) the diagram would be: 1s 2s 3s 2p 3p

11. We can shorthand electron orbits without drawing
For our purposes, groups 1 and 2 go in the s-sublevel, groups 13 through 18 go in the p-sublevel, and groups 3 through 12 go in the d-sublevel
For example, fluorine (9 electrons) has a configuration of 1s2 2s2 2p5
ex: What’s the configuration of Aluminum?
Aluminum has 13 electrons 2 2 6 2 1
The orbits go 1s 2s 2p 3s 3p

12. Abbreviated e- Configurations
You must pay attention to the noble gases
You write down the noble gas that comes right before the element in the question.
Then, continue the configuration from that noble gas until you’ve reached the desired element.
ex: Silicon
(14 e-)
1s2 2s2 2p6 3s2 3p2
[Ne] 3s2 3p2

13. ex 2: Iodine
1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p5
[Kr] 5s2 4d10 5p5

14. Electronegativity
Recall from Chem 30S that electronegativity is the ability of an atom in a molecule to attract electrons to itself.
You will be responsible for identifying four types of electronegative bonds:
1) non-polar covalent
2) moderately polar covalent
3) very polar covalent
4) ionic

15. To classify the bonds, you must take the difference between the electronegativity for the two elements in the bond.
If the difference is:
1) 0.0 – 0.4 = non-polar covalent
2) 0.4 – 1.0 = moderately polar covalent
3) 1.0 – 2.0 = very polar covalent
4) > 2.0 = ionic
ex: What bond type is O:Br?
O = 3.5, Br = 2.8
so 3.5 – 2.8 = 0.7
Therefore, it’s moderately polar covalent!

16. Periodic and Group Trends
There are three areas trends that we will look at: atomic radii, ionization energy, and ionic radii.
1) Atomic Radii
a) Periodic Trends: the atomic radii generally decreases as you move across the period. This is because each additional valence electron get pulled closer to the nucleus making the atomic radius smaller

17. b) Group Trends: the atomic radii increases as you move down a group
b) Group Trends: the atomic radii increases as you move down a group. This is because as you go down, you have more orbits which shield the valence electrons from the pull of the nucleus therefore the atomic radius gets bigger.
ex: Which has a bigger radii, Carbon or Tin?
Since they are in the same group, Tin is bigger because it is further down!

18. Visually we can see this trend in atomic radii with the following:

19. 2) Ionic Radii
a) Periodic Trends: positive and negative ions get smaller as you move across the period
b) Group Trends: both positive and negative ions get bigger as you mover down the group.
ex: Which is bigger, P-3 or Cl-?
Since negative ions get smaller as you move across, P-3 is the bigger ion!

20. Visually we can see this by:

21. 3) Ionization Energy
a) Periodic Trends: ionization energy increases as you move across a period. This is because the more electrons an atom has, the stronger the hold on the electrons (towards a full shell).
b) Group Trends: ionization energy decreases as you move down a group. This is because each successive step pushes the valence electrons further from the nucleus allowing less energy to remove them (because the attraction is lower).

22. ex: Which has higher ionization energy, Li or K?
Since Li is higher, it has a higher ionization energy.
Visually we can see this trend: