1. The Periodic Table and Periodic Law
Chapter 6
The Periodic Table and
Periodic Law

2. 6.1 Development of the Modern Periodic Table
1. Antoine Lavoisier ( )
Compiled a list of all known elements
23 know elements at that time

3. 6.1 Development of the Modern Periodic Table
2. John Newlands (1837 – 1898)
Law of Octaves = every 8th element repeats a common set of properties
Not widely accepted due to missing elements and his use of musical terminology

4. 6.1 Development of the Modern Periodic Table
3. Dimitri Mendeleev (1834 – 1907)
1869 he published 1st periodic table by atomic mass and chemical properties
Predicted the properties of missing elements: scandium, gallium, and germanium

5. 6.1 Development of the Modern Periodic Table
4. Henry Moseley (1887 – 1915)
1913 equated number of protons with atomic number
Reordered the P.T. by atomic # fising some of the elements that didn’t fit their spots based on properties

6. 6.1 Development of the Modern Periodic Table
5. periodic law:
The repeating pattern of chemical and physical properties when elements are arranged by atomic number

7. 6.1 Development of the Modern Periodic Table
6. organization
Groups/families:
Columns of/on the P.T.

8. 6.1 Development of the Modern Periodic Table
Periods:
Rows of/on the P.T.

9. 6.1 Development of the Modern Periodic Table
Main group/representative elements:
Elements in groups where the number is followed with an “A”
Have a wide range of chemical and physical properties

10. TABLE: Metals: Location = to the left of the staircase except H
Properties = shiny, mostly solids, good conductors of heat/electricity, and malleable/ductile
Examples = copper (Cu), gold (Au), iron (Fe)

11. TABLE: Nonmetals: Location = to the right of the staircase plus H
Properties = brittle (when solid), mostly gases, poor conductors
Examples =helium (He), oxygen (O), Iodine (I)

12. TABLE: Semimetals/metalloids: Location = along the staircase except Al
Properties = properties of both metals and nonmetals
Examples = B, Si, Ge, As, Sb, Te, Po, and At

13. TABLE: Alkali metals: Location = 1A (except H)
Valence e- and charge = 1 ve- and +1
Properties = highly reactive; soft, gray solids

14. TABLE: Alkaline Earth metals: Location = 2A
Valence e- and charge = 2 ve- and +2
Properties = very reactive; soft, gray solids

15. TABLE: Transition metals: Location = “B” groups
Valence e- and charge = 2 ve- and +1,+2,+3 or +4
Properties = shiny, good conductors, can be polyvalent (means can have more than 1 possible charge)

16. TABLE: Halogens: Location = 7A Valence e- and charge = 7 ve- and -1
Properties = highly reactive; can be solids, liquids or gases

17. TABLE: Noble Gases: Location = 8A
Valence e- and charge = 8 ve- and no ion formation
Properties = extremely unreactive, gases, full outer energy level

18. TABLE: Rare Earth metals: Location = Bottom double rows
Valence e- and charge = 2 ve-
Properties = often used as phosphors (elements that emit light when struck by electrons)
Also know as the “Lanthanides” and “Actinides”

19. 6.2 Classification of the Elements
1. Valence Electrons = electrons in the highest principle energy level
Within a period: elements have the same # of energy levels as the period # where they are found
Within a group: elements have the same # of ve-s as their group (representative elements) and all transition and rare earth elements have 2 ve-s

20. 6.2 Classification of the Elements
2. s, p, d, f blocks
s block: group 1A, 2A, hydrogen and helium
p block: groups 3A-8A but not He
d block: transition elements; all have 2 ve-s because d’s are 1 energy level behind
f block: rare earth elements; all have 2 ve-s because f’s are 2 energy levels behind

21. 6.3 Periodic Trends Periodicity:
The repeating nature of the properties of the elements creating common groups (periodic law)

22. 6.3 Periodic Trends Atomic Radius
DEFINITION: relative size; distance from the center of the atom to the edge of the e- shell
PERIOD TREND:
GROUP TREND:

23. 6.3 Periodic Trends Ionic Radius
DEFINITION: relative size; distance from the center of the ion to the edge of the e- shell
PERIOD TREND:
GROUP TREND:

24. 6.3 Periodic Trends Ionization Energy
DEFINITION: the energy required to remove an electron from a gaseous atom
PERIOD TREND:
GROUP TREND:

25. 6.3 Periodic Trends Electronegativity
DEFINITION: the ability of an atom to attract electrons while in a chemical bond
PERIOD TREND:
GROUP TREND:

26. 6.3 Periodic Trends Lower Left Large (atomic/ionic radius)
Lower Left Low (ionization E./electroneg.)