1. The Periodic Table

2. During the nineteenth century, chemists began to categorize the elements according to similarities in their physical and chemical properties. The end result of these studies was our modern periodic table.

3. Dmitri Mendeleev
In 1869 he published a table of the elements organized by increasing atomic mass.

4. Lothar Meyer
At the same time, he published his own table of the elements organized by increasing atomic mass.

5. Henry Moseley
In 1913, through his work with X-rays, he determined the actual nuclear charge (atomic number) of the elements*. He rearranged the elements in order of increasing atomic number.
*“There is in the atom a fundamental quantity which increases by regular steps as we pass from each element to the next. This quantity can only be the charge on the central positive nucleus.”

6. Periodic Table Geography

7. Periodic Table Columns called Families/Groups
Family # indicates # valence (outer shell) electrons
Elements in same family have similar properties
Rows called Periods
Row # indicates # energy levels in atom

8. The horizontal rows of the periodic table are called PERIODS.

9. The elements in any group of the periodic table have similar physical and chemical properties!
The vertical columns of the periodic table are called GROUPS, or FAMILIES.

10. Metals/Nonmetals/Semiconductors
Metals: excellent conductors of heat & electricity; have luster, are ductile/malleable
Nonmetals: poor conductors of heat & electricity; are dull & brittle
Semiconductors(Metalloids): elements that under certain conditions conduct heat & electricity

11. Families of Elements Family 1: Alkali Metals
Family 2: Alkaline Earth Metals
Families 3 to 12: Transition Metals
Family 13: Boron Family
Family 14: Carbon Family
Family 15: Nitrogen Family
Family 16: Oxygen Family
Family 17: Halogens
Family 18: Noble Gases
Three general groups: metals, nonmetals, & semiconductors(metalloids)

12. Periodic Table

15. Periodic Law
When elements are arranged in order of increasing atomic number, there is a periodic pattern in their physical and chemical properties.

16. Alkali Metals

17. Alkaline Earth Metals

18. Transition Metals

19. Metals Alkali metals (Family 1) Alkaline Earth metals (Family 2)
Very reactive
Has 1 valence electron
When ionized has charge of 1+
Alkaline Earth metals (Family 2)
Reactive
Has 2 valence electrons
When ionized has charge of 2+
Transition metals (Families 3 to 12)
Somewhat reactive
Valence electron number varies
Ionized charge varies

20. InnerTransition Metals
These elements are also called the rare-earth elements.
InnerTransition Metals

21. Halogens

22. Noble Gases

23. Nonmetals
Include H, some elements from families 13 to 16, all elements from families 17 & 18. Zig-zag line divides metals from nonmetals.
Inert gases are unreactive; contain 8 valence electrons
Halogens are very reactive; contain 7 valence electrons; gain electrons becoming negatively charged
Elements in other families gain electrons to become negatively charged
These elements plentiful on Earth

24. Semiconductors (aka Metalloids)
Located along the zig-zag line
Includes:
Boron (B); Silicon (Si), Germanium (Ge), Arsenic (As), Antimony (Sb), Tellurium (Te), Polonium (Po)
Notice that Al is not considered a metalloid, it is considered a metal
Conduct heat & electricity under certain conditions
B is hard & added to steel to increase hardness; Sb is bluish-white and shin, Te is silvery-white & electrical conductivity increases with light exposure, Si important in solar cells & integrated circuits

25. Periodic Table Trends

26. #1. Atomic Size - Group trendsH
As we increase the atomic number (or go down a group). . .
each atom has another energy level,
so the atoms get bigger. Li Na K Rb

27. #1. Atomic Size - Period Trends
Going from left to right across a period, the size gets smaller.
Electrons are in the same energy level.
But, there is more nuclear charge.
Outermost electrons are pulled closer. Na Mg Al Si P S Cl Ar

28. #2. Trends in Ionization Energy
Ionization energy is the amount of energy required to completely remove an electron (from a gaseous atom).
Removing one electron makes a 1+ ion.
The energy required to remove only the first electron is called the first ionization energy.

29. Ionization Energy
The second ionization energy is the energy required to remove the second electron.
Always greater than first IE.
The third IE is the energy required to remove a third electron.
Greater than 1st or 2nd IE.

30. #3. Trends in Electronegativity
Electronegativity is the tendency for an atom to attract electrons to itself when it is chemically combined with another element.
They share the electron, but how equally do they share it?
An element with a big electronegativity means it pulls the electron towards itself strongly!

31. Electronegativity Group Trend
The further down a group, the farther the electron is away from the nucleus, plus the more electrons an atom has.
Thus, more willing to share.
Low electronegativity.

32. Electronegativity Period Trend
Metals are at the left of the table.
They let their electrons go easily
Thus, low electronegativity
At the right end are the nonmetals.
They want more electrons.
Try to take them away from others
High electronegativity.

33. The periodic table is the most important tool in the chemist’s toolbox!