1. Quantum Theory & Periodicity

2. Atomic Models Rutherford Nucleus with electrons revolving in orbits
Couldn’t explain why electrons didn’t crash into nucleus

3. Bohr’s Model Electrons can only be certain distance from nucleus
Each distance corresponds to certain quantity of energy
Close to nucleus – lowest energy level
Difference in energy between 2 energy levels = quantum of energy
Cannot be in between levels
Does not give off energy while in a given level

4. Present Day Electrons located in orbitals
Regions with high probability of finding electrons
Electrons clouds

5. Electrons & Light
Bohr
Electrons can move from low to high energy by absorbing energy
Electrons are unstable at high energy level  move to lower energy level by releasing energy
Released as light (with specific wavelength)
Each move from a level will release light of a different wavelength
Ground state – at lowest possible energy
Excited state – higher energy
Pg picture of hydrogen emitting light

6. Electron Configuration
Form of notation which shows how the electrons are distributed among various orbitals and energy levels
1s1 = hydrogen
1 = energy level
s = sublevel
1 = number of electrons in that sublevel

7. So if n = 3  sublevels would be s, p, and d (3 total)
n = energy level
Indicates how many sublevels there are
n = 1  1 sublevel
n = 2  2 sublevels
n = 3  3 sublevels
n = 4  4 sublevels
Sublevels
1st sublevel = s
2nd = p
3rd = d
4th = f
So if n = 3  sublevels would be s, p, and d (3 total)

8. What is the order of the sublevels?

9. Orbitals Each type of sublevel holds different number of orbitals
Orbitals can hold 2 electrons
Pauli Exclusion Principle
Sublevel
# of orbitals
Max # of electrons s 1 2 p 3 6 d 5 10 f 7 14

10. Aufbau Principle
Electrons fill orbitals with the lowest energy first

11. Orbital Notation (or diagram)
Follow Hund’s rule – orbitals are filled with one electron of the same spin before pairing it with the second electron of a different spin

12. Orbitals & Electron Capacity of First Four Energy Levels
Principle Energy Level
Type of sublevel
Number of orbitals per type
Number of orbitals per level (n2)
Max # of electrons (2n2) 1 s 2 4 8 p 3 9 18 d 5 16 32 f 7

14. Exceptions Cr Cu [Ar] 4s2 3d4  [Ar] 4s1 3d5
Half filled and full d orbitals are more stable

15. Areas of Periodic Table
Main group elements
Group 1, 2, 13-18
Regular and consistent with same number of valence electrons
Group 1- alkali metals
Group 2- alkaline earth metals
Group 17- halogens
Group 18- noble gases

16. Metals Alkali Alkaline earth
React with water to make alkaline solutions (base)
K + H20  KOH
Stored in oil
Not found as pure elements in nature
Alkaline earth
Found as compounds in nature also
More stable  takes more energy to lose 2 electrons

17. Metals Transition Lanthanides & Actinides Different configurations
Placed due to lack of space
Named after 1st element in row
Actinides are radioactive

18. Halogens Most reactive nonmetals  only needs to gain 1 electron
React with metal to make salts
F, Cl  g
Br  l
I, At  s

19. Other groups Noble gases Hydrogen
Full set of electrons in outer energy level
Stable
Inert- but able to get Xe to react
Hydrogen
In class by itself
1 proton, 1 electron
Reacts with many elements

20. Terms to know Electron shielding Nuclear charge
Inner electrons pulled into nucleus due to attractive forces protect the electrons farther out
Only a factor when going down a group
Nuclear charge
Protons in nucleus attract the electrons (pull them in)
Only a factor when going across a period

21. Atomic radius Size of atom Hard to determine Going down a group
Another principal energy level is filled
Electron shielding protect outer electrons from being pulled in
So as you move down, radius increases

22. Atomic radius (cont.) Going across a period
Electrons being added to same energy level
Pulled in due to attractive forces with protons (nuclear charge increases)
So going across, radius decreases

24. Ionic radius (size of ion)
Metals - the ion of a metal is smaller than a neutral atom of that metal
forms an ion as it loses electrons from the valence shell, typically emptying the shell of all electrons size of atom is smaller.

25. Non-metals - the ion of a non-metal is larger than a neutral atom of that non-metal
forms an ion as it gains electrons in its valence shell  By filling the valence shell it gets larger because it is experiencing a high electron-electron repulsion in the valence shell.

27. Ionization Energy Energy required to remove an electron
Moving down a group
Element contains more electrons
More energy levels
Electron shielding protects them  takes less energy to remove
So as you move down, IE decreases

28. Ionization energy (cont.)
Moving across a period
Number of protons and electrons increase
Electrons added to the same energy level
Nuclear charge increases (pull electrons in)
More energy to remove electrons
So as you move across, IE increases

29. Successive Ionization Energies - more energy is required to remove additional electrons
IE1 < IE2 < IE3

31. Electronegativity Ability of atom in compound to attract electrons
Number that is derived from various measurable properties
Scale from 0 to 4, with a value of 4 representing an element with the greatest attraction for a free electron
Higher number, bigger pull on electrons towards itself

32. Electronegativity (cont.)
Going down a group
Distance to valence electrons increases
Cannot attract them
Smaller electronegativity
Going across a period
Electrons pulled in closely
Higher electronegativity