1. Light Energy and Electron Configurations
Chapter 5.1
Light Energy and Electron Configurations

2. QUANTUM THEORY
Describes how electrons behave like both waves and particles
Is the modern theory of the atom
Properties of waves:
property
definition
symbol
SI units
velocity
distance traveled per second c
m/s
Amplitude
peak height above the midline A m
wavelength
peak-to-peak distance λ
Frequency
number of peaks passing by per second ν
s-1 (Hertz)

3. 3 ways to tell a particle from a wave:
wave behavior particle behavior
waves interfere particles collide
waves diffract particles effuse
waves are delocalized particles are localized

4. Electron Configurations
Chapter
Electron Configurations

5. The Bohr Model of the Atom
Ground State
the lowest energy state of an atom
all electrons are in their lowest possible orbitals.
when atoms are in the ground state, they do not emit light

6. The Bohr Model of the Atom
Excited State
state in which an atom or molecule picks up outside energy, causing an electron to move into a higher-energy orbital
when the electron falls back to a lower-energy orbit, the electron emits an amount of energy (quantum) -- light

7. The Bohr Model of the Atom
Electrons only move about in certain ENERGY LEVELS.
The smaller the electron’s orbit, the LOWER the atom’s energy state, or ENERGY LEVEL.
Bohr assigned a number, n, to each energy level and called this the PRINCIPAL QUANTUM NUMBER

9. The Heisenberg Uncertainty Principle
It is fundamentally impossible to know precisely both the VELOCITY and LOCATION of a particle at the same time.
By measuring one, the other is affected.

10. ATOMIC ORBITALS
The quantum model of the atom shows that electrons do not follow orbits around the nucleus, but rather have areas of high probability where they may be found. These areas are called ENERGY LEVELS.
We use QUANTUM NUMBERS to describe the location of an electron.

11. ATOMIC ORBITALS
Each energy level is broken up into sub-levels that have different shapes.
These sublevels are defined by their shapes as S, P, D, or F.
Each sub-level contains a SPECIFIC NUMBER OF ORBITALS

12. WHAT DO THE ORBITALS LOOK LIKE?

13. ATOMIC ORBITALS
Each orbital can hold 2 electrons that are spinning in OPPOSITE DIRECTIONS.
Each sub-level has a different number of orbitals :
Sub-Level
# of Orbitals s 1 p 3 d 5 f 7

14. Reading the Periodic Table

15. WRITING ELECTRON CONFIGURATIONS
The arrangement of an atom’s electrons is called the ELECTRON CONFIGURATION.
Three rules define how electrons can be arranged in an atom’s orbitals:

16. AUFBAU PRINCIPLE
Electrons must occupy the orbital with the lowest energy first

17. Increasing energy 6s 5s 4s 3s 2s 1s 6p 5p 5d 4p 4d 4f 3p 3d 2p

19. Pauli Exclusion Principle
Orbitals can only have 2 electrons max
No two electrons in the same atom can be in the same orbital, spinning the same way.
To occupy the same orbital, two electrons must have opposite spins; that is, the electron spins must be paired.

20. Hund’s Rule
Every orbital in a subshell is singly occupied with one electron before any one orbital is doubly occupied
all electrons in single occupied orbitals have the same spin.

22. NOTATION: Complete electron configurations are written using:
Numbers to represent ENERGY LEVEL (n)
Letters to represent SUBLEVELS (l)
Superscript numbers to represent INDIVIDUAL ELECTRONS
The order in which the energy levels, sub-levels, and orbitals will be filled is represented on your PERIODIC TABLE

23. 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   

24. Arrangement of Electrons in Atoms
Electrons in atoms are arranged as
SHELLS (n)
SUBSHELLS (l)
ORBITALS (ml)

25. Writing Atomic Electron Configurations
Two ways of writing configs. One is called the spdf notation. 1 s
value of n
value of l
no. of
electrons
spdf notation
for H, atomic number = 1

26. Another way of writing electron configurations is called ORBITAL BOX NOTATION.
This includes additional information showing each orbital and each electron including the spin!

27. Draw the orbital notation for chlorine
Spdf notation:
1s2 2s2 2p6 3s2 3p5
Orbital Diagram: 1s 2s 2p 3s 3p

28. Draw the orbital notation for aluminum
Spdf notation:
1s2 2s2 2p6 3s2 3p1
Orbital Diagram: 1s 2s 2p 3s 3p

29. Practice Problems: 1s22s22p3
Write the electron configuration for nitrogen.
1s22s22p3

30. Practice Problems: 1s22s22p63s23p64s2
Write the electron configuration for calcium
1s22s22p63s23p64s2

31. Practice Problems: 1s22s22p63s23p64s23d104p65s24d10 5p66s24f145d106p6
Write the electron configuration for Radon.
1s22s22p63s23p64s23d104p65s24d10
5p66s24f145d106p6

32. Practice Problems: 1s22s22p63s23p64s23d104p65s24d10
Write the electron configuration for Americium.
1s22s22p63s23p64s23d104p65s24d10
5p66s24f145d106p67s25f7

33. Kernel (Core) Notation
I wish there was some faster way...

34. What is the kernel (core) of an atom?
The kernel, or core, of an atom includes all of the nucleus of the atom and any electrons not on the outermost energy level.
For electron configurations, you can use brackets and the elemental symbol of the previous noble gas to represent the electron configuration for the core or kernel of the atom.
[Ne] = 1s22s22p6

35. What is the core notation for:
Chlorine
[Ne] 3s23p5
Uranium
[Rn] 7s25f4

36. Orbitals fill in an order
Lowest energy  higher energy.
Adding e-’s changes energy of orbital. Full orbitals are best situation.
half filled orbitals  lower energy, next best
more stable.

37. Write the electron configurations for these elements:
Titanium - 22 electrons
1s22s22p63s23p64s23d2
Vanadium - 23 electrons
1s22s22p63s23p64s23d3
Chromium - 24 electrons
1s22s22p63s23p64s23d4 (expected)
But this is not what happens!!

38. Chromium is actually: 1s22s22p63s23p64s13d5 Why?
This gives us two half filled orbitals (the others are all still full)
Half full is slightly lower in energy.
The same principal applies to copper.

39. Copper’s electron configuration
Copper has 29 electrons so we expect: 1s22s22p63s23p64s23d9
But the actual configuration is:
1s22s22p63s23p64s13d10
This change gives one more filled orbital and one that is half filled.
Remember these exceptions: d4, d9

40. Irregular configurations of Cr and Cu
Chromium steals a 4s electron to make its 3d sublevel HALF FULL
Copper steals a 4s electron to FILL its 3d sublevel

41. Exceptional Electron Configurations
Some actual electron configurations differ from those assigned using the aufbau principle because although half-filled sublevels are not as stable as filled sublevels, they are more stable than other configurations.

42. Electron Configurations for IONS
Isoelectronic with noble gases
“iso-” = SAME “electronic”= ELECTRON CONFIGURATIONS
When making positive ions… they can only lose valence electrons!
When making negative ions… they can only gain valence elctrons!

43. What is the electron configuration for a sodium ion?
Na+1
Number of Electrons = 10!!
1s22s22p6

44. Excited State Electron configurations
Atoms that are excited have electrons that have been elevated to higher energy levels or subshells due to an input of energy.
They will violate the Aufbau principle!!
You can tell which element it is by counting the superscripts to see how many total electrons there are 

45. Valence Electrons Determine the chemical properties of an element
bonding
Number of electrons found in the atom’s outermost orbitals
Only the s and p electrons of the highest energy level!
Corresponds with the element’s group number
Transition metals may have multiple valence electrons

46. Electron-Dot Structures
Consists of an element’s symbol surrounded by dots
The dots represent the valence electrons

47. Reading the Periodic Table