1. Electrons and Periodicity
Adapted from

2. Rutherford Model of the Atom (1911)
What is wrong with this model of the atom?

3. Bohr Model of the Atom (1913)
Electrons are found in circular paths around the nucleus
The paths are called Energy Levels (n)
Each level has a specific energy
The lowest level is called Ground State
Electrons can jump to higher levels
by absorbing energy (Excited State)
Electrons can return to a lower
level by releasing energy (light)
Bohr’s model only works for
atoms of Hydrogen

4. Wave-Particle Duality
Photons of light and most subatomic particles (i.e. electrons) behave as both PARTICLES and WAVES.

5. Heisenberg Uncertainty Principle (1927)
Problem of defining nature of electrons in atoms solved by W. Heisenberg.
Cannot simultaneously define the position and momentum (= m•v) of an electron.
We can define electron energy exactly but accept limitation that we do not know exact position.
W. Heisenberg

6. Heisenberg Uncertainty Principle (2015)

7. Electron Cloud Model (1926)
Proposed by Erwin Schrödinger
Applied idea of electrons behaving as a wave
Electrons are no longer in defined paths/orbits
Uses Electron Clouds (Orbitals) – defined areas where an electron is likely to be found
Based on probability

8. Electromagnetic Spectrum
In increasing energy, ROY G BIV

9. Wave Behavior Visible light Ultraviolet radiation wavelength Amplitude
Node

10. Visible Spectrum What colors have --long wavelengths?
Long wavelength --> small frequency
Short wavelength --> high frequency
What colors have
--long wavelengths?
--high frequencies?
--short wavelengths?
--small frequencies?
increasing frequency
increasing wavelength

11. Wave Behavior
Wavelength – distance between 2 crests of a wave
Abbreviated with Greek letter lambda (λ)
Units are meters (m)
Frequency – speed at which crests pass a given point
Abbreviated with Greek letter nu (ν)
Units are 1/s, s-1 or Hertz
Wavelength and Frequency are related by the equation:
 •  = c
c =speed of light = 3.00 x 108 m/sec
Inverse relationship (one goes up, the other goes down)

12. E = h x ν Wave Behavior Energy and frequency are directly related
Amount of energy released by an excited electron corresponds to a specific frequency of light
E = h x ν
(E) Energy measured in Joules (J)
(h) Planck’s constant = x J s
(ν) Frequency measured in 1/s or hertz

13. Wave Behavior
What is the frequency of a wave that has a wavelength of 2.56 x 10-7 m?
c = λ x ν
2.8 x 108 m/s = (2.56 x 10-7 m) x ν
2.8 x 108 m/s = ν
2.56 x 10-7 m
1.1 x /s = ν
What is the Energy of this wave?
E = h x ν
E = (6.626 x Js) x (1.1 x /s)
E = 7.2 x J

14. Learning Check
What is the frequency of EMR having a wavelength of 5.55 x 10-7 m?
What is the Energy of this same wave?

15. Continuous Spectrum of White Light
If you separate white light with a prism you will see ALL colors of the rainbow
Continuous because there are no gaps between colors

16. Creating an Atomic Spectrum
Discharge tube – tube containing atoms of a gas connected to an electricity source
When electricity is turned in the gas will emit light

17. Atomic Spectrum of Excited Hydrogen Gas
If the light from a discharge tube is separated by a prism
you will only see very specific lines of color separated by
black regions
Called atomic emission spectra or line spectra

18. Atomic Emission Spectra of Excited Atoms
When electricity is supplied to the gas atoms the electrons absorb Energy and become EXCITED
Electrons eventually have to release this energy in the form of light
Excited atoms emit light of only certain wavelengths (lines of color)
The wavelengths of emitted light depend on the element.

19. Line Spectra of Other Elements

20. Arrangement of Electrons in Atoms
Electrons in atoms are arranged as
LEVELS (n)
SUBLEVELS (l)
ORBITALS (ml)

21. Quantum Numbers
Every electron has a unique set of 4 quantum numbers that defines it’s position in the atom
Think of it as the electron’s address!

22. Quantum Numbers

23. Learning Check Write a set of quantum numbers for the 4f orbital
Which orbital is represented by the following sets of Quantum numbers
n = 3, l = 0, ml = 0
n = 2, l = 1, ml = 1
n = 4, l = 2, ml = -1
n = 3, l = 3, ml = 2
n = 3, l = 1, ml = 2

24. Energy Levels
Each energy level has a number called the PRINCIPAL QUANTUM NUMBER (n)
Currently n can be 1 thru 7, because there are 7 periods on the periodic table

25. Energy Levels
n = 1
n = 2
n = 3
n = 4

26. Types of Orbitals Orbital - most probable area to find an electron
Orbitals can take on different shapes
So far, we have 4 shapes. They are named s, p, d, and f.
No more than 2 e- assigned to an orbital – one spins clockwise, one spins counterclockwise

27. Types of Orbitals

28. S Orbital (sharp) The s orbital is the simplest orbital
The shape is a sphere
An electron can move anywhere within the sphere
L = 0
Every principal level (n) has an s orbital

29. Relative sizes of the spherical 1s, 2s, and 3s orbitals of hydrogen.

30. p Orbitals (proximal)
A p orbital has a dumb-bell shape
There are three different directions for p-orbitals
These are called x, y, and z
The electron can move anywhere in the nodes
L = 1
Every level above n = 1 has a p orbital
There is a PLANAR NODE thru the nucleus, which is an area of zero probability of finding an electron

31. p Orbitals
The three p orbitals lie 90o apart in space

32. d Orbitals (diffuse) A d orbital has a clover shape
There are 5 different directions for a d orbital
L= 2
Every level above n=2 has a d orbital

33. The shapes and labels of the five 3d orbitals.

34. f Orbitals (fundamental)
f orbitals have a flower shape
There are 7 directions for d orbitals
L= 3
Every level above n = 3 has an f orbital

35. Electron Configurations
The arrangement of electrons in an atom into atomic orbital
Electron configurations are filled according to three rules (Aufbau, Pauli, Hund)
We need electron configurations so that we can determine the number of electrons in the outermost energy level. These are called valence electrons.
The number of valence electrons determines the number and types of bonds and atom can make
1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p6 6s2 4f14… etc.

36. 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.

37. Electron Configuration Rules
Aufbau Principle – electrons are filled starting from the lowest energy levels first
Pauli Exclusion Principle – no more than 2 electrons can fill into orbitals at one time and they must have opposite spin (clock-wise/counter-clockwise)
Hund Rule – In p, d and f sublevels, electrons will have to fill one clockwise electron into all available orbitals before filling in counter-clockwise electrons
Energy Level (n)  Sublevel (l)  Orbital (ml)

38. Aufbau Principle
Aufbau Principle states that electrons fill from the lowest possible energy to the highest energy
The Aufbau Diagram is a memory device that helps you remember the order of the filling of the orbitals from lowest energy to highest energy

39. Aufbau Diagram Why are d and f orbitals always in lower energy levels?
d and f orbitals require LARGE amounts of energy to fill
It’s easier to fill an s orbital before a d or f
This is the reason for the diagram
BE SURE TO FOLLOW THE ARROWS IN ORDER!

40. Pauli Exclusion Principle
How many total electrons can be in a sublevel?
Remember: A maximum of two electrons can be placed in an orbital.
s orbitals
p orbitals
d orbitals
f orbitals
Number of
orbitals
Number of electrons

41. Lithium
Group 1A
Atomic number = 3
1s22s1 ---> 3 total electrons

42. Carbon Group 4A Atomic number = 6 1s2 2s2 2p2 --->
6 total electrons
Here we see HUND’S RULE.

43. Let’s Try It!
Write the electron configuration for the following elements: H He B N F Ne Na

44. Orbitals and the Periodic Table
Orbitals grouped in s, p, d, and f orbitals (sharp, proximal, diffuse, and fundamental)
The d orbitals are n-1 for Rows 4-7
The f orbitals are n-2 for Rows 6-7
s orbitals
d orbitals
p orbitals
f orbitals

45. Shorthand Notation A way of abbreviating long electron configurations
Use Noble gases as placeholders for completely filled levels
Only write out the configuration for the Valence (outermost) Level
He Ne Ar Kr Xe Rn

46. Shorthand Notation
Step 1: Find the closest noble gas to the atom (or ion) that has a lower atomic number.
Step 2: Write the symbol for the Noble Gas in brackets [ ]
Step 3: Find the energy level below your Noble Gas
Step 4: Write the configuration for that energy level and stop when you reach your target element

47. Shorthand Notation [Ne] 3s2 3p5 Chlorine
Longhand is 1s2 2s2 2p6 3s2 3p5
You can abbreviate the first 10 electrons with a noble gas, Neon. [Ne] replaces 1s2 2s2 2p6
The next energy level after Neon is 3
Write out the configuration for level 3 until you reach Cl
[Ne] 3s2 3p5

48. Practice Shorthand Notation
Write the shorthand notation for each of the following atoms: Br K Sb Zn Sn Np

49. Valence Electrons
Electrons are divided between core and valence electrons
B 1s2 2s2 2p1
Core = [He] , valence = 2s2 2p1
Br [Ar] 3d10 4s2 4p5
Core = [Ar] 3d10 , valence = 4s2 4p5

50. Electron Configurations for Ions
Remember…
Negative ions have gained electrons
Positive ions have lost electrons
Electrons that are lost or gained should be added/removed from the highest energy level (not the highest orbital in energy!)

51. Electron Configurations for Ions
Tin (Sn)
Atom: [Kr] 5s2 4d10 5p2
Sn+4 ion: [Kr] 4d10
Sn+2 ion: [Kr] 5s2 4d10
Note that the electrons came out of the highest energy level, not the highest energy orbital!

52. Electron Configurations for Ions
Bromine
Atom: [Ar] 4s2 3d10 4p5
Br- ion: [Ar] 4s2 3d10 4p6
Note that the electrons went into the highest energy level, not the highest energy orbital!

53. Ion Configurations
To form anions from elements, add 1 or more e- from the highest sublevel.
P [Ne] 3s2 3p3 + 3e- ---> P3- [Ne] 3s2 3p6 or [Ar]

54. Try Some Ions! Write the longhand notation for these: F- Li+ Mg+2
Write the shorthand notation for these:
O-2
Ba+2
Al+3

55. Exceptions to the Aufbau Principle
(HONORS only)
Exceptions to the Aufbau Principle
Remember d and f orbitals require LARGE amounts of energy
If we can’t fill these sublevels, then the next best thing is to be HALF full (one electron in each orbital in the sublevel)
There are many exceptions, but the most common ones are the elements whose configurations end in
d4 – Cr, Mo, W
d9 – Cu, Ag, Au

56. Exceptions to the Aufbau Principle
(HONORS only)
Exceptions to the Aufbau Principle
d4 is one electron short of being HALF full
In order to become more stable (require less energy), one of the closest s electrons will actually go into the d, making it d5 instead of d4.
Example: Cr (theoretical) [Ar] 4s2 3d4
Cr (actual) [Ar] 4s1 3d5.
Procedure: Find the closest s orbital. Steal one electron from it, and add it to the d.

57. Exceptions to the Aufbau Principle
(HONORS only)
Exceptions to the Aufbau Principle
OK, so this helps the d, but what about the poor s orbital that loses an electron?
Remember, half full is good… and when an s loses 1, it too becomes half full!
So… having the s half full and the d half full is usually lower in energy than having the s full and the d to have one empty orbital.

58. Exceptions to the Aufbau Principle
(HONORS only)
Exceptions to the Aufbau Principle
d9 is one electron short of being full
Just like d4, one of the closest s electrons will go into the d, this time making it d10 instead of d9.
Example: Au (theoretical) [Xe] 6s2 4f14 5d Au (actual) [Xe] 6s1 4f14 5d10.
Procedure: Same as before! Find the closest s orbital. Steal one electron from it, and add it to the d.

59. Write the shorthand notation for: Cu W Au
(HONORS only)
Try These!
Write the shorthand notation for: Cu W Au

60. General Periodic Trends
Atomic and ionic size
Ionization energy
Electronegativity
Higher effective nuclear charge
Electrons held more tightly
Larger orbitals.
Electrons held less
tightly.

61. Atomic Size Size goes UP on going down a group.
Because electrons are added further from the nucleus, there is less attraction. This is due to additional energy levels and the shielding effect. Each additional energy level “shields” the electrons from being pulled in toward the nucleus.
Size goes DOWN on going across a period.

62. Atomic Size
Size decreases across a period owing to increase in the positive charge from the protons. Each added electron feels a greater and greater + charge because the protons are pulling in the same direction, where the electrons are scattered.
Large
Small

63. Which is Bigger?
Na or K ?
Na or Mg ?
Al or I ?

64. Ion Sizes Does the size go
up or down when losing an electron to form a cation?

65. Ion Sizes CATIONS are SMALLER than the atoms from which they come.Li +
, 78 pm
2e and 3 p
Forming a cation.
Li,152 pm
3e and 3p
CATIONS are SMALLER than the atoms from which they come.
The electron/proton attraction has gone UP and so size DECREASES.

66. Ion Sizes
Does the size go up or down when gaining an electron to form an anion?

67. Ion Sizes Forming an anion.-
, 133 pm
10 e and 9 p
F, 71 pm
9e and 9p
Forming an anion.
ANIONS are LARGER than the atoms from which they come.
The electron/proton attraction has gone DOWN and so size INCREASES.
Trends in ion sizes are the same as atom sizes.

68. Trends in Ion Sizes
Figure 8.13

69. Which is Bigger?
Cl or Cl- ?
K+ or K ?
Ca or Ca+2 ?
I- or Br- ?

70. Ionization Energy
IE = energy required to remove an electron from an atom (in the gas phase).
Mg (g) kJ ---> Mg+ (g) + e-
This is called the FIRST ionization energy because we removed only the OUTERMOST electron
Mg+ (g) kJ ---> Mg2+ (g) + e-
This is the SECOND IE.

71. Trends in Ionization Energy
IE increases across a period because the positive charge increases.
Metals lose electrons more easily than nonmetals.
Nonmetals lose electrons with difficulty (they like to GAIN electrons).

72. Trends in Ionization Energy
IE increases UP a group
Because size increases (Shielding Effect)

73. Which has a higher 1st ionization energy?
Mg or Ca ?
Al or S ?
Cs or Ba ?

74. Electronegativity, 
 is a measure of the ability of an atom in a molecule to attract electrons to itself.
Concept proposed by
Linus Pauling

75. Periodic Trends: Electronegativity
In a group: Atoms with fewer energy levels can attract electrons better (less shielding). So, electronegativity increases UP a group of elements.
In a period: More protons, while the energy levels are the same, means atoms can better attract electrons. So, electronegativity increases RIGHT in a period of elements.

76. Electronegativity

77. Which is more electronegative?
F or Cl ?
Na or K ?
Sn or I ?