1. THE PERIODIC TABLE
HISTORY
&
STRUCTURE

2. Historical Background
Classification of the elements began in 1817 with J.W. Dobereimer.
Grouped the elements in groups of 3 by atomic number.
Problem: did not leave room or system for the addition of new elements.

3. And then…..
1866: John Newlands (musician) grouped the elements by atomic mass in groups of 8
Problem: Still no room for new elements, and atomic mass was a problem when you consider isotopes.

4. But then 1869: Dmitri Mendeleev (Russia) and Julius Meyer (Germany)
Grouped elements in sequence from left to right and top to bottom by increasing atomic mass and reaction with chlorine and oxygen
Gave us the first significant organization of elements in that new elements could be added by their reactivity

5. But then:
In 1905, Henry Mosely discovered that the elements could be grouped according to their energy relationships in order of increasing atomic number
THIS LED TO……….

6. THE PERIODIC LAW:
The physical and chemical properties of elements are the periodic functions of their atomic number

7. Elements are arranged in GROUPS & PERIODS
Vertical columns of elements= GROUPS
similar chemical properties
same number of electrons in their outermost shell (valence shell)

8. Horizontal Rows of elements = PERIODS
No similar chemical properties
Members have the same number of electron shells

9. Elements may be categorized as:
METALS
Chemical Properties
Form + charged ions by losing electrons
Have 3 or less electrons in their valence shells
Physical Properties
Good conductors of heat and electricity
High density
Luster
Malleable and ductile

10. NONMETALS Chemical Properties Physical Properties
Form – ions by gaining electrons (except 8A)
Have 5-8 electrons in the valence shell
Physical Properties
Poor conductors of heat and electricity
Solids tend to be brittle and non-ductile
Low density

11. METALLOIDS Have characteristics of metals and nonmetals
Traditionally used as semiconductors

12. Some groups have special names and characteristics
Group 1A: The Alkali Metals
1 electron in valence shell
Form +1 ions
Highly reactive with oxygen and water
Mixed in compounds to form alkaline mixtures

13. Group 2A: The Alkali Earth Metals
2 electrons in valence shell
Form +2 ions
Reactive
Form alkaline solutions in water
Fire resistant (do not melt in fire)

14. Group 6A: The Chalcogens
6 electrons in valence shell
All found in copper ores (named from Greek word chalcos, meaning ore)
Show increasing metallic properties as atomic number increases.
Form -2 ions

15. Group 7A: The Halogens
Very active because of need to fill valence shell
Form - 1 ions
7 electrons in valence shell
Tend to form salts with elements from 1A and 2A

16. Group 8A: The Noble Gases
8 electrons in valence shell “filled”
Inert gases (stable and unreactive)

17. Arrangement of electrons determines chemical properties
Model of arrangement is called
Noble Gas Shell Model
Shell: region of space about the nucleus where electrons reside
Each shell represents collection of probability clouds
Shells are 3-D spheres
Each shell holds max number of electrons filling from the inside to the outside
1st shell: 2 electrons
2 nd shell: 8 electrons
3 rd shell: 18 electrons
Electrons fill shells in one direction until ½ filled, then pair in the opposite direction

18. Effective Nuclear Charge: Z*
Effective nuclear (Z*) charge is the essentially the positive charge from the nucleus that an outer electron “sees.”
The greater the Z*, the more attracted an electron is to the nucleus.
Outer shells have diminished nuclear charges
Effective nuclear charge diminishes with distance from the nucleus
Larger the atom, lower the effective nuclear charge
The less filled the valence shell, the lower the effective nuclear charge
Effective nuclear charge increases from the lower left hand corner to the upper right hand corner of the periodic table

20. Higher the Z* = Higher Ionization Energy
Relates to an atom’s ability to lose electrons
Represents the amount of energy needed to remove an electron from the valence shell
Higher the Z* = Higher Ionization Energy
Higher Ionization Energy = greater tendency to gain electrons
Lower Ionization Energy = greater tendency to lose electrons

21. Electron Affinity Relates to the ease by which atoms gain electrons
Atoms on the upper right side of the periodic table have a greater electron affinity
Exception: Noble Gases have no electron affinity