1. Modern Atomic Model

2. Where are the electrons?
We know they are outside of the nucleus.
We say they are in an electron cloud.
But where?

3. Quantum-mechanical model
Takes into account the following:
Treats e- as waves within the atom.
e- inside of atoms have specific E and occupy 3-D regions about the nucleus called orbitals. [Orbitals are different than orbits.]
The size and shape of the orbitals depends on the E of the e- that occupy them.
All orbitals in an atom make up the e- cloud around the nucleus. The e- cloud gives the atom a size and shape.
The e- can’t be located exactly in the atom. There are areas of probability to find an e-.

4. The Quantum Mechanical Model
Has energy levels for electrons.
Contains orbitals of varying shapes and sizes
Limitation: It can only tell us the probability of finding an electron a certain distance from the nucleus.

5. Quantum Numbers
Principal Quantum Number (n) = the energy level of the electron.
Within each energy level complex math describes several shapes.
These are called atomic orbitals

6. Summary of atomic orbitals
# of shapes
Max electrons
Starts at energy level s 1 2 1 p 3 6 2 d 5 10 3 7 14 4 f

7. By Energy Level First Energy Level only s orbital only 2 electrons
Second Energy Level
s and p orbitals are available
2 in s, 6 in p
8 total electrons

8. By Energy Level Third energy level s, p, and d orbitals
2 in s, 6 in p, and 10 in d
18 total electrons
Fourth energy level
s,p,d, and f orbitals
2 in s, 6 in p, 10 in d, and 14 in f
32 total electrons

9. By Energy Level
Any thing past the fourth level - not all the orbitals will fill up.
You simply run out of electrons
The orbitals do not fill up in a neat order.

10. Filling order Lowest energy fill first. The energy levels overlap
The orbitals do not always fill up order of energy level.

11. Increasing energy 7s 2p 3p 4p 5p 6p 7p 3d 4d 5d 6d 6s 5s 4f 5f 4s 3s

12. Electron Configurations
The way electrons are arranged in atoms are guided by 3 rules:
Aufbau principle - electrons enter the lowest energy first
This causes difficulties because of the overlap of orbitals of different energies.
Hund’s Rule - When electrons occupy orbitals of equal energy they don’t pair up until all orbitals in that level are half full
Pauli Exclusion Principle - at most 2 electrons per orbital with different spins

13. Electron Configurations
Notations of e- in atoms:
Orbital diagram - unpaired e- represented by  or  , paired e- shown as .
The arrows are used to show that the two e- have opposite spin quantum #s.) (Use the periodic table to show how to do the following ex: H  1s, He 1s. [The E level & orbital should be written below the bar…] Do more ex.

14. Electron Configurations
Let’s determine the electron configuration for Phosphorus
Need to account for 15 electrons

15. Increasing energy 7p 6d 5f 7s 6p 5d 6s 4f 5p 4d 5s 4p 3d 4s 3p
The first two electrons go into the 1s orbital
Notice the opposite spins
only 13 more

16. Increasing energy 7p 6d 5f 7s 6p 5d 6s 4f 5p 4d 5s 4p 3d 4s 3p
The next electrons go into the 2s orbital
only 11 more

17. Increasing energy 7p 6d 5f 7s 6p 5d 6s 4f 5p 4d 5s 4p 3d 4s 3p
The next electrons go into the 2p orbital
only 5 more

18. Increasing energy 7p 6d 5f 7s 6p 5d 6s 4f 5p 4d 5s 4p 3d 4s 3p
The next electrons go into the 3s orbital
only 3 more

19. Increasing energy 7p 6d 5f 7s 6p 5d 6s 4f 5p 4d 5s 4p 3d 4s
The last three electrons go into the 3p orbitals.
They each go into separate shapes
3 unpaired electrons
1s2 2s2 2p6 3s2 3p3

20. Electron Configurations
Notations of e- in atoms:
Orbital diagram - unpaired e- represented by  or  , paired e- shown as .
Write orbital diagrams for elements 1-10.
Write orbital diagrams for Cr & Ce.
The arrows are used to show that the two e- have opposite spin quantum #s.) (Use the periodic table to show how to do the following ex: H  1s, He 1s. [The E level & orbital should be written below the bar…] Do more ex.

21. Electron Configurations
Notations of e- in atoms:
Orbital diagram - unpaired e- represented by  or  , paired e- shown as .
e- configuration notation - no more lines & arrows, uses # of e- in a sublevel as a superscript over the sublevel designation
ex: H 1s1, He 1s2
Aufbau exceptions:
Expected: Actual:
Cr 1s2 2s2 2p6 3s2 3p6 4s2 3d4 1s2 2s2 2p6 3s2 3p6 4s1 3d5
Cu 1s2 2s2 2p6 3s2 3p6 4s2 3d9 1s2 2s2 2p6 3s2 3p6 4s1 3d10
1/2 full & filled d orbitals are more stable, so they take s e- to fill

22. Electron Configurations
Notations of e- in atoms:
Orbital diagram - unpaired e- represented by  or  , paired e- shown as .
e- configuration notation - no more lines & arrows, uses # of e- in a sublevel as a superscript over the sublevel designation
Write electron configuration notation for elements 1-10.
Write electron configuration notation for Cr & Ce.
ex: H 1s1, He 1s2
Aufbau exceptions:
Expected: Actual:
Cr 1s2 2s2 2p6 3s2 3p6 4s2 3d4 1s2 2s2 2p6 3s2 3p6 4s1 3d5
Cu 1s2 2s2 2p6 3s2 3p6 4s2 3d9 1s2 2s2 2p6 3s2 3p6 4s1 3d10
1/2 full & filled d orbitals are more stable, so they take s e- to fill

23. Electron Configurations
Notations of e- in atoms:
Shortcut for orbital diagrams & e- configuration notation
Uses noble gas (group 18 elements) “core”
ex: H 1s1, He 1s2
Aufbau exceptions:
Expected: Actual:
Cr 1s2 2s2 2p6 3s2 3p6 4s2 3d4 1s2 2s2 2p6 3s2 3p6 4s1 3d5
Cu 1s2 2s2 2p6 3s2 3p6 4s2 3d9 1s2 2s2 2p6 3s2 3p6 4s1 3d10
1/2 full & filled d orbitals are more stable, so they take s e- to fill

24. Noble Gas Shortcuts… 1. Find element on periodic table.
2. Move up 1 row on table and go to Noble Gas at end of that row
3. Put this noble gas symbol inside [ ].
4. Now, write out what is left over after the [ ].

25. Noble gas in row above is He [He] 2s1 is the same as 1s2 2s1
Examples
Li =1s2 2s1
Noble gas in row above is He
[He] 2s1 is the same as 1s2 2s1
Be = 1s2 2s2
[He] 2s2 is the same as 1s2 2s2
Na = 1s2 2s2 2p6 3s1
Noble gas in row above is Ne
[Ne] 3s1 is the same as 1s2 2s2 2p6 3s1

26. Electron Configurations
Try it…
Noble Gas Shortcut
Cr (use orbital diagrams)

27. Electron Configurations
Try it…
Noble Gas Shortcut
Ce (use e- configuration)

28. Periodic Table History

29. History of the Periodic Table
Dobereiner
Triads of elements with shared properties

30. History of the Periodic Table
Cannizzaro
method for measuring atomic masses and interpreting the results of measurements

31. History of the Periodic Table
Newlands
arranged elements by atomic masses, properties repeated after every 8 elements → law of octaves

32. History of the Periodic Table
Mendeleev & Meyer
arranged elements according to the increase in atomic mass

33. History of the Periodic Table
Mendeleev
left spaces for undiscovered elements & predicted properties of those elements
credited with discovering periodicity

34. History of the Periodic Table
2 Questions:
Why could most elements be arranged by increasing atomic mass, but a few could not?
What was the reason for chemical periodicity?

35. History of the Periodic Table
Mosely
shooting electrons at various metals to produce X-rays
frequencies of the X-rays were unique to the metals
assigned a whole number to each element → atomic numbers

36. History of the Periodic Table
Mosely
shooting electrons at various metals to produce X-rays
arrange elements by atomic numbers to get families with similar properties

37. Periodic Table Information

38. Types of Elements

39. Periodic Table Information
Types of elements:
Metal – lustrous, good conductors, most are solids, malleable, ductile
Nonmetal - poor conductors, no luster, neither malleable nor ductile, most are gases, wide variety of other physical properties
Metalloid - properties of metals & non-metals

40. Parts of the Periodic Table

41. Periodic Table Information
horizontal rows of elements
Group (family)
vertical columns of elements

42. Periodic Table Information
Group (family)
alkali metals – group 1 (except hydrogen)
alkaline earth metals – group 2
transition metals – all of the d-block elements
inner transition metals – all of the f-block elements
metalloids – elements that touch the “staircase”
halogens – group 17
noble gases – group 18 (elements with filled outer shells of electrons)

43. Sodium + Water video Potassium + Water video

44. Calcium + Water video

45. Chlorine + Sodium video

46. Back to Electrons… (electron configurations)

47. Electron Configuration Notation (using shortcut)

48. More Electron Configurations
Practice writing shortcut e- configurations.
Element # 15
Element # 8
Element # 34

49. Group Electron Configuration

50. More Electron Configurations
Group e- configurations – all elements in family have similar “endings”

51. More Electron Configurations
Write the shortcut e- configuration for:
Li, Na, K
all end with s1, so group configuration is s1

52. More Electron Configurations
Write the shortcut e- configuration for:
C, Si, Ge
all end with p2, so group configuration is p2

53. More Electron Configurations
Practice identifying elements using group configurations:
Element in period 3 with group configuration p4
Element in period 6 with group configuration s2
Element in period 2 with group configuration p6
Element in period 7 with group configuration s1

54. Periodic Trends

55. Periodic Trends
Electron configurations are able to cause periodic variations in elemental properties...

56. Periodic Trends
Valence electrons - electrons that may be lost/gained/shared when chemical compounds are formed

57. Periodic Trends Valence electrons
Period  As we go across a period, the number of valence electrons increases.

58. Periodic Trends Valence electrons
Group  As we go down a group, we find that the number of valence electrons stays constant.

59. Periodic Trends
Size of the atomic radius – one-half the distance between the nuclei of identical atoms joined in a molecule

60. Periodic Trends Size of the atomic radius –
Period  As we go across a period, the general trend is for the atomic radii to decrease.
This trend is due to the increasing positive charge in the nucleus – stronger ability to pull in the electrons.

61. Periodic Trends Size of the atomic radius –
Group  As we go down a group, there is a general increase in atomic radii. This happens because we are seeing more and more energy levels being added.

62. Periodic Trends
Ionization energy – energy required to overcome nuclear attraction and remove an electron from a gaseous element
Fluorine is the most electronegative element and is assigned a value of 4. All other values are calculated relative to this value.

63. Periodic Trends Ionization energy
Period  As we go across a period, the general trend is for the ionization energy to increase.
This means that the elements on the right side of the Periodic table “want” electrons more than the elements on the left.

64. Periodic Trends Ionization energy
Group  As we go down a group, there is a general decrease in ionization energy.
This is also due to size considerations. Larger atoms have more “shielded” nuclei than smaller elements.

65. Periodic Trends
Electronegativity - measure of the power of an atom in a chemical compound to attract electrons
Fluorine is the most electronegative element and is assigned a value of 4. All other values are calculated relative to this value.

66. Periodic Trends Electronegativity
Period  As we go across a period, the general trend is for the electronegativity to increase.
This means that the elements on the right side of the Periodic table “want” electrons more than the elements on the left.

67. Periodic Trends Electronegativity
Group  As we go down a group, there is a general decrease in electronegativity.
This is also due to size considerations. Larger atoms have more “shielded” nuclei than smaller elements.

68. Alkali Metals + Water (trends in reactivity of an elemental family)
Also found at
Video from which is ok, but not as exciting as it could be… -
Best video so far -