Periodicity

The Periodic Table

I.
History

A. Mendeleev:	Designed First Periodic Table in 1869
	Based the table on increasing atomic mass
	Made an eight column table that left blanks for elements that weren’t discovered yet (ex. Ge)
	His table had exceptions based on chemical properties
	-Tellurium switched with Iodine
	-Cobalt switched with Nickel
	-Argon switched with Potassium
B. Mosley:	Designed the Modern Periodic Table
	Research into protons in 1913 gave him the basis for MODERN PERIODIC LAW
C. Modern Periodic Law:	The properties of the elements are a periodic function of their atomic numbers (number of protons)

II.
Vocabulary

Word	Definition
Period	Row
Family	Group or Column
Group 1 (except H)	Alkali Metals
Group 2	Alkaline Earth Metals
Group 3-11	Transition Metals
Group 17	Halogens
Group 18	Inert or Noble Gases
Elements 57-71	Lanthanide Series
Elements 89-103	Actinide Series
Luster	mirror like shine *
Conductivity	ability to transfer e- well *
Malleability	ability to be hammered into sheets *
Ductility	ability to be drawn into wire *
Brittle	Crumbles, easily broken #

* Metals # nonmetals

III.
Trends

A. Metallic Character: Degree to which it acts like a metal. Chemically this is defined as giving up electrons to form a positive ion.
	· As you go across the table the value DECREASES
You can see this when we compare Na, a metal located on the left of Period 3, to Ar which is a non-metal on the right of Period 3. As we progress from left to right across the table we cross the "staircase" that divides metal and non-metals on the table.
Why?
As we move across a period, we increase the number of protons in the nucleus of the atom. This increased nuclear charge makes it harder to remove electrons from the atom.
	·As you go down the table the value INCREASES
It is easiest to find this trend when we look at group 14 or 15. At the top of the table we have C and N. Both of these elements are nonmetals. As we progress down the table we cross the "staircase" that indicates the presence of metalloids then we get to the metals Pb and Bi at the bottom of these columns.
Why?
As we move down a group we add occupied principal energy levels to the atom. Each added level acts to shield the valence e- from the pull of the nucleus. This makes it easier for electrons to be removed.

B. Ionization Energy: Energy required to remove an electron
	· As you go across the table the value INCREASES
We can get this information directly from Reference Table S. Ionization Energies are listed on this table and we just simply look up a few sample elements to see the trends that form.
Why?
The reason remains the same as before. An increase in the number of protons in the nucleus of the atom increases the nuclear charge and makes it harder to remove electrons from the atom.
	· As you go down the table the value DECREASES
Again we can use Table S to get this information.
Why?
As seen in the diagram above, the added principal energy levels create a barrier that shields the valence e- from the pull of the nucleus. This makes it easier for electrons to be removed.

C. Electronegativity: Attraction for a bonded pair of electrons
	· As you go across the table the value INCREASES
Electronegativity follows the same trend as ionization energy and is listed on reference table S. Like before, we look up selected elements in order to find the pattern.
Why?
Once again there is an increase in the number of protons in the nucleus of the atom. Thus increasing the nuclear charge. This makes it easier for an element to attract a bonded pair of electrons.
	· As you go down the table the value DECREASES
Why?
Again we can see in the diagram above, the added principal energy levels create a barrier that shields the valence e- from the pull of the nucleus. In this case that reduces an elements ability to attract a shared pair of electrons.
Group 18 is the exception to this trend because it has a stable valence octet and will not bond. If no bond forms then you can't have an attraction for a bonded pair of electrons.

D. Atomic Size: based on radius not on mass
	· As you go across the table the value DECREASES
Reference Table S lists this value directly in the form of atomic radii. We can look up a few sample elements to see the trends that form across and down the periodic table.
Why?
As we progress across the table the nuclear pull grows because of the increase in the number of protons. This attraction pulls the outer electrons in closer making it smaller in size.
	· As you go down the table the value INCREASES
Why?
The atom grows in size because with each step down the table we add a principal energy level causeing the atom to get larger. In addition, we see the shielding (diagram above) of the nuclear attraction that occurs due to the addition of principal energy levels.

E. Ionic Radius: a comparison between the radius of an atom and the radius of it’s ion
	· Ionic radius in Metals is SMALLER than atom
Why?
Simply stated metals form positive ions by losing electron(s). When they lose electrons they get smaller.
	·Ionic radius in Nonmetal is LARGER than atom
Why?
Nonmetals form negative ions by gaining electron(s). When they gain electrons the atom gets larger.

F. Reactivity: for a metal thats giving up e- and for a non-metal that's gaining electrons
	· In metals: As you go down the table INREASES
Why?
Again we can use the diagram above to see that an increase in the number of principal energy levels creates a barrier that shields the valence e- from the pull of the nucleus. This makes it easier for electrons to be given up.
	· In nonmetals: As you go down the table DECREASES
Why?
Eelectrons must be gained by a nonmetal in a reaction. The shielding created by the addition of principal energy levels will make it more difficult for the nucleus to grab electrons from other atoms.

IV.
Chemistry of a Group

Group 1:	Alkali Metals (except H)
	Have 1 valence e-
	Form +1 ion
	Most Metallic
	Lowest Ionization energy and electronegativity
	Highly Reactive
	Not found freely in Nature (only in compounds)

Group 2:	Alkaline Earth Metals
	Have 2 valence e-
	Form +2 ion
	Very Metallic
	Low Ionization energy and electronegativity
	Reactive
	Not found freely in Nature (only in compounds)

Groups 3-11:	Transition Metals
	Multiple Oxidation States
	more than one type of ion and more than one compound
	Colored solutions
	Copper = blue	Nickel = green
	Cobalt = Red	Iron = orange/yellow

Group 17:	Halogens
	Have 7 valence e-
	Form -1 ion
	Most Reactive Non-Metals
	High Ionization energy and electronegativity
	Highly Reactive
	Not found freely in Nature (only in compounds)

Group 18:	Inert or Noble Gases
	Have 8 valence e-
	Form no ion
	Most Stable
	Highest Ionization energy and no electronegativity
	Found as monatomic gas in Nature (not in compounds)

Metalloids or Semimetals (7)

Have both metallic and nonmetallic properties
Boron	(B)	Silicon	(Si)	Germanium	(Ge)
Arsenic	(As)	Antimony	(Sb)	Tellurium	(Te)
		Astatine	(At)

Last Revised May 9, 2007