1. Unit 5: Electron Configuration, Periodic Table
Chapter 3 & 4

2. Key Concepts Orbitals, Energy Levels, Energy Sublevels
Electron Configuration
Noble Gas Shortcut
s, p, d, f Blocks
Early Periodic Table, Modern Periodic Table
Periodic Law
Periods, Groups
Valence Electrons, Octet/Duet Rules, Ions
Alkali Metals, Alkaline Earth Metals, Halogens, Noble Gases
Metals, Nonmetals, Transition Metals
Periodic Table Trends

3. Orbitals
Atomic Orbitals: where electrons are likely to be found in an atom
Energy Levels are the size of orbitals
Energy Sublevels are the shapes of orbitals

4. Energy Levels (size of orbital)
Lower energy levels are smaller and closer to the nucleus
Higher energy levels are bigger and farther away from the nucleus IR
Visible
Light
Electrons in higher energy levels have greater potential energy.
Dropping to n=1 (UV) releases more energy than dropping to n=2 (visible light) releases more energy than dropping to n=3 (IR), because of the difference in energies/distance.
[x-ray and gamma ray are radiation particles that leave the nucleus, not electrons in energy levels] UV

5. Energy Sublevels (shape of orbitals)
p, d, and f orbitals have multiple orientations in space
s orbital
p orbital
d orbital
f orbital
1 orientation
3 orientations
5 orientations
7 orientations

6. Energy Sublevels (shape of orbitals)
Orbitals Website
Orbitals Website:

7. Orbitals
Energy Levels: 1, 2, 3, 4…
Energy Sublevels: s, p, d, f

8. Orbitals 2 electrons fit in each orbital! Energy Level
(size of orbitals)
Energy Sublevel
(shape of orbitals)
Number of Orbitals
Number of Electrons 1 s
1 total
2 e- 2
s, p
1, 3 = 4 total
8 e- 3
s, p, d
1, 3, 5 = 9 total
18 e- 4
s, p, d, f
1, 3, 5, 7 = 16 total
32 e-
2 electrons fit in each orbital!

9. Electron Configuration
Electron Configuration: the energy levels and sublevels that electrons exist in
Ex. The electron configuration of neon is written
Ne: 1s2 2s2 2p6
e- always fill up orbitals in the same order!
(lower energy level  higher energy level)

10. Electron Configuration
Write out the electron configuration of neon:
How many electrons does the element have?
Ne has 10 e-
Fill the orbitals with electrons
Ne: 1s2 2s2 2p6
2 e- in orbital 1s, 2 e- in orbital 2s, & 6 e- in orbital 2p = 10 e- total

12. Electron Configuration
Write the electron configuration for the following elements
O K+
Mg S2-
O (8 e-): 1s2 2s2 2p4
Mg (12 e-): 1s2 2s2 2p6 3s2
K+ (18 e-): 1s2 2s2 2p6 3s2 3p6
S2- (18 e-): 1s2 2s2 2p6 3s2 3p6 (Should be 3p4, but because 2 added e- it is 3p6)

13. Noble Gas Shortcut
Elements gain or lose e- until they’re isoelectronic (have the same e- configuration) with a noble gas
Noble Gas Shortcut: put the closest noble gas before an element in [brackets] to shorten its electron configuration
Ex. What is the noble gas shortcut of phosphorus?
P: [Ne] 3s2 3p3

14. Periodic Table

15. Periodic Table Rows = Energy Levels Blocks = Energy Sublevels 3 2 1 45 6 7 1 3 2 4 5 6 3 4 5 6 7 4 5
Rows = Energy Levels
Blocks = Energy Sublevels

16. Periodic Table 7p
Rows = Energy Levels
Blocks = Energy Sublevels

17. Early Periodic Table Mendeleev: invented the first periodic table
Grouped the known elements together according to their properties
Put elements in order of their atomic mass
Mendeleev left gaps for undiscovered elements. For example, he knew all the elements in group 4 except for the one between Silicon and Tin, and he even made predictions about its properties. When Germanium was discovered, his predictions were confirmed.

18. Modern Periodic Table
Elements are now organized by their atomic number (# of protons), not their atomic mass

19. Periodic Law
Periodic Law: when elements are arranged by their atomic numbers, elements with similar properties appear at regular intervals
Elements in the same columns have similar properties!
Periodic Law: elements with similar properties appear “periodically”

20. Period vs. Group Period: horizontal row on the periodic table
Group: vertical column on the periodic table

21. Groups s Block p Block d Block f Block Alkali Metals: Group 1
Alkaline Earth Metals: Group 2
p Block
Other metals, nonmetals, & metalloids: Groups 3-6
Halogens: Group 7
Noble Gases: Group 8
d Block
Transition Metals
f Block
Rare Earth Elements

22. Groups

23. Valence Electrons
Valence Electrons: electrons in the outermost shell/orbital of an atom
Valence electrons of an atom participate in chemical reactions with other atoms
Elements with the same number of valence electrons tend to react in the same way C H

24. Octet Rule
Octet Rule: atoms gain or lose e- to attain an e- configuration of the nearest noble gas
Noble gases have complete outer e- shells, which make them stable!
Duet Rule is for H & He only
Octet Rule is for all other elements

25. Ions +1 +2 +3 ±4 -3 -2 -1
TM form ions with various charges
When an atom/molecule gains a negative electron, it becomes a negative ion or an anion.
When an atom/molecule loses a negative electron, it becomes a positive ion or cation.
When an atom or molecule gains a negative electron, it becomes a negative ion
When an atom or molecule loses a negative electron, it becomes a positive ion

26. Alkali Metals Alkali Metals: Highly reactive, reacts with water
Loses 1 valence e- to get to stable e- configuration
Stored in oil so they don’t react with the moisture in the air

27. Alkaline Earth Metals Alkaline Earth Metals:
Highly reactive, reacts with oxygen
Loses 2 valence e- to get to stable e- configuration
Slightly less reactive than alkali metals because it takes more energy to lose 2 e- than 1 e-

28. Halogens Halogens: Highly reactive, reacts with metals making salts
Gains 1 valence e- to get to stable e- configuration

29. Noble Gases Noble Gases: stable, not reactive
Full set of e- in their outermost energy level

30. Metals vs. Nonmetals

31. Metals vs. Nonmetals Metals Nonmetals
Left side of periodic table
Shiny
Malleable/ductile
Conductive
Solid at room temp
Forms ions with positive charge
Right side of periodic table
Dull
Brittle
Not conductive
Gas at room temp
Forms ions with negative charges
Metalloids: along staircase between metals & nonmetals
Both metallic and non-metallic properties

32. Transition Metals
Same properties as the other metals with a few additional properties
Forms ions with various positive charges
Forms colorful compounds
Good catalysts
Catalysts accelerate a chemical reaction without being affected by it.
They aren’t actually part of the chemical reaction, they just speed it up.
Catalysts:
-Manganese dioxide
-Iron oxide
-Copper oxide
Catalyst

33. Transition Metals Fe(SO4) Fe(NO3)3
Work out the charges of iron in both compounds. Fe2- is green and Fe3- is orange.
Fe(SO4)
Fe(NO3)3

34. Transition Metals + Potassium dichromate Sodium sulfite
The color change indicates a change in charge.
Potassium
dichromate
Sodium
sulfite +

35. Periodic Table Trends Radius: size of atom
Electronegativity: attraction of e- in a bond
Ionization Energy: energy required to remove e-
When atoms are smaller, their e- are closer

36. Periodic Table Trends Radius: size of atom
Electronegativity: attraction of e- in a bond
Ionization Energy: energy required to remove e-
Radius Decreases
Electronegativity Increases
Ionization Energy Increases
When atoms are smaller, their e- are closer

37. Periodic Table Trends
When atoms are smaller, their e- are closer to the nucleus so it takes more energy to remove them.
When atoms are bigger, their e- are farther from the nucleus so it takes less energy to remove them.
When atoms are smaller, their e- are closer

38. Periodic Table Trends
Ex. Compare barium and sulfur’s radius, electronegativity, and ionization energy.
Ex. What is the largest element on the periodic table? What is the smallest element on the periodic table?
Radius Decreases
Electronegativity Increases
Ionization Energy Increases