1. Electron Configuration

3. The Electromagnetic Spectrum
Know what is in the red boxes
High frequency
Short wavelength
High energy
lower frequency
longer wavelength
lower energy

5. Jumping Electrons
normally electrons exist in the ground state, meaning they are as close to the nucleus as possible
when an electron is excited by adding energy to an atom, the electron will absorb energy and "jump" to a higher energy level
heating a chemical with a
Bunsen burner is enough
energy to do this

6. Emission line spectrum
energy is applied to a specific element
this “excites” the element
visible light, infrared, and UV can be emitted when electrons fall back down to the ground state
unique for every element and are used to identify atoms (much like fingerprints are used to identify people)

7. Give off energy when falls back down to ground energy level
More on emission line spectrum
Give off energy when falls back down to ground energy level

8. this is a repetitive slide- just couldn’t bear to delete it
an electron in the atom gains (absorbs) energy from heating
electron jumps up an energy level.
electron is now unstable (unwelcome) in this level and is “kicked out”
when the electron loses the energy and come back to the original level, light is emitted

11. Scandium 3-D video (2:31)
3-D Graphic Examples of Atomic Orbitals

12. Quantum Numbers (however, actual numbers are often not used)
each electron in an atom is described by four different quantum numbers
think of the 4 quantum numbers as the address of an electron… country > state > city > street
electrons fill low energy orbitals before they fill higher energy ones

13. Principle quantum number (n)
Quick intro, more later.
Principle quantum number (n)
describes the SIZE of the orbital or ENERGY LEVEL (shell) of the atom.
Angular quantum number (l)
a SUB-LEVEL (shell) that describes the type or SHAPE of the orbital
Magnetic quantum number (m)
the NUMBER of orbitals
describes an orbital's ORIENTATION in space
Spin quantum number (s)
describes the SPIN or direction (clockwise or counter-clockwise) in which an electron spins

14. Principle Quantum # (n)
LEVEL/SIZE 1 2 3 4
Angular Quantum # (l)
ORBITAL SHAPE or SUBLEVEL s
s p
s p d
s p d f
Magnetic Quantum # (m)
AXIS/
ORIENTATION
or ORBITALS
1 orbital
4 total
orbitals
9 total orbitals
16 total orbitals
Spin Quantum # (s)
DIRECTION OF ELECTRON SPIN
2 e-
8 e-
18 e-
32 e-

15. 4f 4d 4p 4s
14 (7)
10 (5)
6 (3)
2 (1) 32
= level and sub-level
= max. # of electrons
= # of electrons
= number of orbitals 3d 3p 3s
10 (5)
6 (3)
2 (1) 18 2p 2s
6 (3)
2 (1) 8 1s
2 (1) 2

16. Principle Quantum Number (n) or Energy Level
values 1-7 used to specify the level the electron is in
describes how far away from the nucleus the electron level is
the lower the number, the closer the level is to the atom's nucleus and less energy
maximum # of electrons that can fit in an energy level is given by formula 2n2 (n = energy level)

19. Angular Quantum Number (l) or Sub-Levels
determines the shape of the sub-level
number of sub-levels equal the level number
ex. the second level has two sub-levels, the third has three and so on…
they are numbered but are also given letters referring to the sub-level type
l=0 refers to the s sub-level
l=1 refers to the p sub-level
l=2 refers to the d sub-level
l=3 refers to the f sub-level
just know this

22. Magnetic quantum number (m) or Orbitals
Electron Orbitals YouTube 1:37
the third of a set of quantum numbers
tells us how many sub-levels there are of a particular type and their orientation in space of a particular sub-level
only two electrons can fit in an orbital
= electron

23. S sub-level has only 1 orbital only holds two electrons

24. P sub-level has 3 orbitals holds up to six electrons

26. D sub-level has 5 orbitals holds up to 10 electrons

28. F sub-level has 7 orbitals holds up to 14 electrons

31. Spin quantum number (s)
the fourth of a set of quantum numbers
number specifying the direction of the spin of an electron around its own axis.
only two electrons of opposite spin may occupy an orbit
the only possible values of a spin quantum number are +1/2 or -1/2.

33. Principle Quantum # (n)
LEVEL/SIZE 1 2 3 4
Angular Quantum # (l)
ORBITAL SHAPE or SUBLEVEL s
s p
s p d
s p d f
Magnetic Quantum # (m)
AXIS/
ORIENTATION
or ORBITALS
1 orbital
4 total
orbitals
9 total orbitals
16 total orbitals
Spin Quantum # (s)
DIRECTION OF ELECTRON SPIN
2 e-
8 e-
18 e-
32 e-

34. Principle energy level (n) Type of sublevel
Table 3-6b Orbitals and Electron Capacity of the First Four Principle Energy Levels
Principle energy level (n)
Type of sublevel
Number of orbitals per type
Number of orbitals per level(n2)
Maximum number of electrons (2n2) 1 s 2 4 8 p 3 9 18 d 5 16 32 f 7

36. “Rules” for Writing Electron Configurations
a method of writing where electrons are found in various orbitals around the nuclei of atoms.
three rules in order to determine this:
Aufbau principle
Pauli exclusion principle
Hund’s rule

37. Aufbau Principle
electrons occupy the orbitals of the lowest energy first
each written represents an atomic orbital (such as or or or ….)
electrons in the same sublevel/shell have equal energy ( same energy as )
principle energy levels/shells (1,2,3,4..) can overlap one another
ex: 4s orbital has less energy than a 3d orbital

38. Pauli Exclusion Principle
Hamster video 1:00
only two electrons in an orbital
must have opposite spins
represents one electron
represents two electrons in an orbital
actually incorrect as well, see next slide

39. Hund’s Rules
every orbital in a subshell must have one electron before any one orbital has two electrons
all electrons in singly occupied orbitals have the same spin.

40. Writing Orbital Diagrams

41. Energy

45. Orbitals grouped in s, p, d, and f orbitals (sharp, proximal, diffuse, and fundamental)
s orbitals
d orbitals
p orbitals
f orbitals

49. Boron
Atomic # 5

50. Boron ion (3+)
Atomic # 5

51. Neon
Atomic # 10

52. Bromine
Atomic # 35

53. Bromine ion (1-) Atomic # 35

54. ?

56. Orbital diagrams Electron Configurations
1s2 2s2 2p6 3s1
1s2 2s2 2p6 3s2
1s2 2s2 2p6 3s2 3p1
1s2 2s2 2p6 3s2 3p2
1s2 2s2 2p6 3s2 3p3
1s2 2s2 2p6 3s2 3p4
1s2 2s2 2p6 3s2 3p5
1s2 2s2 2p6 3s2 3p6 Na Mg Al Si P S Cl Ar
1s s p s p
Orbital diagrams Electron Configurations

57. Electron Configurations
2p4
Number of electrons in the sublevel
Energy Level
Sublevel
1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p6 6s2 4f14… etc.

58. Writing Electron Configurations
To write out the electron configuration of an atom:
use the principal quantum number/energy level (1,2,3, or 4…)
use the letter term for each sub-level (s,p,d, or f);
don’t worry about orientation such as x,y,z axis but you do have to be able to draw these for IB
use a superscript number indicates how many electrons are present in each sub-level
hydrogen =1s1.
Lithium =1s22s1.
don’t write anything for spin

64. Sometimes levels are switched in order to keep the level together.
Order of Electrons
Sometimes levels are switched in order to keep the level together.
I hate when they do that! 4s requires less energy and I think it should be before 3d.
1s2 2s2 2p6 3s2 3p6 3d10 4s2 4p6 4d10 4f14…
electron configuration video (3:24)

65. exceptions (don’t need to know this, just be aware that there are exceptions)
orbitals “like” to be empty, half filled, or full
therefore, an electron leaves the 4s (leaving it half full) and goes to the 3d in order to make it full Cr
we would predict:
1s2 2s2 2p6 3s2 3p6 4s2 3d4
but it is actually:
1s2 2s2 2p6 3s2 3p6 4s13d5 Cu
1s2 2s2 2p6 3s2 3p6 4s2 3d9
1s2 2s2 2p6 3s2 3p6  4s1 3d10

66. Noble Gas Shortcut
same