1. The Periodic Table

2. Early Periodic Table Simplest arrangement was by atomic weight
Most significant relationships though have nothing to do with weight but how they react with each other
Arrangement by reactivity shows that there is a relationship among atomic weights but its periodic (repeats every 7th element)

3. Periodic Law
When elements are arranged by atomic number (# of p+) there is a repetition of chemical and physical properties

4. Representative Elements
Modern Periodic Table
Organized in columns called groups or families
Rows are called periods
Group A – representative elements (1A-7A)
Group B - transition elements (1B-8B)
Representative Elements
(Group A)
Representative Elements
(Group A)
Transition Elements
(Group B)

5. Classification of Elements
*based on physical properties

6. Metals
Location: middle & left hand side of the periodic table (under “stairs”)
Physical Properties
shiny, smooth, solids (except mercury)
Good conductors of heat and electricity
High densities, melting/boiling points
Malleable – bent or pounded into sheets
Ductile – drawn into wires

8. Nonmetals
Location: top right corner of the periodic table (above “stairs”)
Physical Properties
Gases or brittle, dull looking solids
Poor conductors of heat and electricity
Usually have lower densities, melting point, and boiling point than metals
sulfur

10. Metalloids (Semimetals)
Location: along the “stairs” (except Al)
Physical properties similar to both metals and nonmetals
They are metallic-looking brittle solids
Relatively good electrical conductivity

11. Electronic Stability
Stability based on the electrons (location and # of electrons; i.e. electron configuration)
Determines chemical properties of an element and therefore how elements react (“behave”) with each other and their environment
Elements are grouped by their chemical reactivity (“behavior”) and therefore elements in the same group/family have similar chemical properties

12. Electron Configurations for Alkali Metals
[He]2s1
[Ne]3s1
[Ar]4s1
[Kr]5s1
Lithium 1s22s1 Sodium 1s22s22p63s1 Potassium 1s22s22p63s23p64s1 Rubidium 1s22s22p63s23p64s23d104p65s1

13. Electron Configurations for Alkaline Earth Metals
Beryllium s22s2
Magnesium 1s22s22p63s2
Calcium s22s22p63s23p64s2
Strontium 1s22s22p63s23p64s23d104p65s2
[He]2s2
[Ne]3s2
[Ar]4s2
[Kr]5s2

14. Electron Configuration for Halogens
F – 1s22s22p5 Cl – 1s22s22p63s23p5 Br - 1s22s22p63s23p64s23d104p5
[He]2s22p5
[Ne]3s23p5
[Ar]4s23d104p5

15. Electron Configuration for Noble Gases
He – 1s2 Ne – 1s22s22p6 Ar – 1s22s22p63s23p6 Kr - 1s22s22p63s23p64s23d104p6

16. Ground State Electron Configurations of the Elements
ns2np6
Ground State Electron Configurations of the Elements
ns1
ns2np1
ns2np2
ns2np3
ns2np4
ns2np5
ns2

17. 1A: 𝑛𝑠 1
5A: 𝑛𝑠 2 𝑛𝑝 3
2A: 𝑛𝑠 2
6A: 𝑛𝑠 2 𝑛𝑝 4
3A: 𝑛𝑠 2 𝑛𝑝 1
7A: 𝑛𝑠 2 𝑛𝑝 5
4A: 𝑛𝑠 2 𝑛𝑝 2
8A: 𝑛𝑠 2 𝑛𝑝 6

18. Valence Electrons
Electrons in the outermost s and p orbitals (highest n shell)
These electrons participate in chemical reactions and determine electronic stability
Represented on the periodic table as the group/family numbers (representative elements & noble gases only)

19. Lewis Dot Symbols
Show the chemical symbol for an element surrounded by the number of valance electrons as dots

20. Octet Rule
Atoms gain, lose, or share electrons to acquire a full set of eight valence electrons (to be chemically stable like a noble gas)
Eight is great!!!