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 and the light is viewed through a spectroscope
a continuous spectrum is NOT observed, but a series of very bright lines of specific colors with black spaces in-between instead
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 normal energy level
More on emission line spectrum
Give off energy when falls back down to normal energy level

8. the process
electrons surround the nucleus in specific orbitals or energy levels
when electrons are excited (heat/electricity) they can move to a higher energy level
when they move back down they emit energy in the form of electromagnetic radiation
because electrons can only exist in certain energy levels, only certain transitions can occur
the color of the light emitted depends on the frequency of the emitted photon

9. 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

10. The Atomic Emission Spectrum of Hydrogen
the emission spectrum of hydrogen is the simplest emission spectrum because there is only one electron
it is not uniform, but concentrated into bright lines, indicating the existence of only certain allowed electron energy levels
visible light, infrared, and UV are emitted when electrons fall back down level
McGraw Hill animation link

12. convergence up here (levels are close together)

15. after a short time, this electron will spontaneously "fall" back to a lower energy level, giving off a quantum of light energy called a photon
the key to Bohr's theory was the fact that the electron could only "jump" and "fall" to precise energy levels, thus emitting a limited spectrum of light.
quantum is the amount of energy required to move an electron from one energy level to another

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

18. 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
the first three of these quantum numbers (n, l, and m) represent the three dimensions in which an electron could be found
the fourth quantum number (s) refers to a certain magnetic quality called spin

19. 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

20. 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-

21. 4f 4d 4p 4s
14 (7)
10 (5)
6 (3)
= level and sub-level
= max. # of electrons
= # of electrons
= number of orbitals
2 (1) 32 3d 3p 3s
10 (5)
6 (3)
2 (1) 18 2p 2s
6 (3)
2 (1) 8 1s
2 (1) 2

22. 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

25. 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
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

28. 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

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

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

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

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

37. 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.

39. Principle Quantum # (n)
SHELL/SIZE 1 2 3 4
Angular Quantum # (l)
ORBITAL SHAPE or SUBSHELL 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-

40. 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

42. “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

43. 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

44. 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

45. 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.

46. Writing Orbital Diagrams

47. Energy

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

55. Boron
Atomic # 5

56. Boron ion (3+)
Atomic # 5

57. Neon
Atomic # 10

58. Bromine
Atomic # 35

59. Bromine ion (1-) Atomic # 35

60. ?

62. 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

63. 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.

64. 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

70. 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 4s2 3d10 4p6 5s2 4d10 5p6 6s2 4f14 5d10 6p6 7s2 5f14 6d10 7p6
Weird electron configuration video (3:24)

71. 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

72. Noble Gas Shortcut
same