1661 - Robert Boyle

Defined an element as a substance that could not be broken down into a simpler substance by a chemical reaction.

1829 -Johann Wolfgang Döbereiner

Grouped known elements into families of triads of elements with similar properties.

1865 - John Newlands - "Law of octaves"

Found that similar elements were separated by intervals of eight when listed in order of increasing atomic mass.

1869 - Dmitri Ivanovitch Mendeléev - Created the first accepted version of the periodic table.

	Grouped elements on the basis of similar chemical properties.  Left blank spaces open to add new elements where he predicted they would occur.  Accepted minor inversions when placing the elements in order of increasing atomic mass.  Predicted properties for undiscovered elements, allowing for his theories to be tested.

                A version of Mendeléev's periodic table published in the journal Annalen der Chemie in 1871.

The Modern Periodic Table

1944 - Glenn T. Seaborg - Created the modern version of the periodic table.

Actinide Hypothesis and Periodic Table Reconstruction             "In 1944, I formulated the “actinide concept” of heavy element electronic structure. This concept predicted that the fourteen actinides, including the first eleven transuranium elements, would form a transition series analogous to the rare-earth series of lanthanide elements and therefore show how the transuranium elements fit into the  periodic table."              "I was warned at the time that it was professional suicide to promote this idea, which has since been called one of the most significant changes in the periodic table since Mendeleev’s 19th century design. Luckily, I stuck  to my guns and have seen the actinide concept become the foundation for many significant discoveries in heavy element research." Glenn Seaborg was the only individual to have an element of the periodic table named after him while still living.  He autographed the periodic table in Room 100 during a visit to TAMU in 1998.

See links to periodic table sites below!

Period - A horizontal row in the periodic table.

The energy levels of the s and p orbitals are numbered by the row in which they are located.
e.g.  The 2s orbital is in the second row (Li and Be) and the 3p orbitals are in the third row (Al, Si, P, S, Cl, Ar)
The d orbitals are placed one row below their energy level.
e.g.  The 3d orbitals are in the fourth row

Group - A vertical column, or family, in the periodic table.

Numbered in two ways:

Group 1 (IA) - Alkali Metals (excluding H)

Li (lithium), Na (sodium), K (potassium), Rb (rubidium), Cs (cesium), and Fr (francium)
All form hydroxides (e.g.  NaOH)
All are active metals
Activity increases as you move down the column
React violently when they come into contact with water
All have one valence electron

	Li:		[He] 2s1		Rb:		[Kr] 5s1
	Na:		[Ne] 3s1		Cs:		[Xe] 6s1
	K:		[Ar] 4s1		Fr:		[Rn] 7s1

All lose one electron to form cations with a charge of +1

e.g.  Li⁺, Na⁺, and K⁺

Group 2 (IIA) - Alkaline Earth Metals
Be (beryllium), Mg (magnesium), Ca (calcium), Sr (strontium), Ba (barium), and Ra (radium)
All but Be and Mg are active metals
Activity increases as you move down the column
Ca, Sr, and Ba react violently when they come into contact with water
All have two valence electrons

	Be:		[He] 2s2		Sr:		[Kr] 5s2
	Mg:		[Ne] 3s2		Ba:		[Xe] 6s2
	Ca:		[Ar] 4s2		Ra:		[Rn] 7s2

 
Tend to lose two electrons to form cations with a charge of +2

e.g.  Be²⁺, Mg²⁺, and Ca²⁺

Alkaline earth metals are less active than their adjacent alkali metals

e.g.  Be is less active than Li; Mg is less active than Na

Groups 3-12 (IIIB - VIIIB, IB & IIB) - Transition Metals
All have valence electrons in the d orbitals
Form compounds in which they have various oxidation numbers
Ag (silver) generally forms Ag⁺ (+1 oxidation state)
Zn (zinc) generally forms Zn²⁺ (+2 oxidation state)
Brass is an alloy of copper and zinc
Group 13 (IIIA)
B (boron) - only element in the group that is not a metal; has semimetal and nonmetal characteristics.
Al (aluminum) - fairly active metal, third most abundant in the earth's crust.
Loses three electrons to form Al³⁺
Forms compounds in which it has an oxidation state of +3
Other metals - Ga (gallium), In (indium), and Tl (thallium) - very scarce active metals
Group 14 (IVA)
C (carbon) - nonmetal
Elemental forms of carbon include:
Graphite (crystalline) - Strong bonds between atoms within planes resulting in extremely high melting and boiling points.  Weaker bonds connecting the planes which account for the soft texture of graphite.
Diamond (crystalline) - Hardest naturally occurring substance with extremely high melting and boiling points.  Atoms arranged in a tetrahedral array with strong C-C bonds.
Charcoal - Results from heating wood without oxygen present
Coke (amorphous) - More structured than other amorphous forms of carbon; made from coal.
Carbon Black (amorphous) - Formed by burning natural gas or other carbon compounds in a limited amount of air
Has strong C-C single bonds, C=C double bonds, and CC triple bonds.
Forms covalent bonds with other elements
Can form double and triple bonds with other nonmetals
S (silicon) and Ge (germanium) - semimetals
Silicon is the second most abundant element in the earth's crust
Sn (tin) and Pb (lead) - less reactive metals
Both form compounds in which their oxidation states are +2 or +4
Bronze is an alloy of tin and copper
Group 15 (VA)
N (nitrogen) and P (phosphorus) - nonmetals
Nitrogen is found in its elemental form at room temperature as a diatomic gas (N₂).
Nitrogen makes up approximately 80% of earth's atmosphere by volume.
In compounds, the oxidation of nitrogen can range from -3 to +5.
Haber process - mixing N₂ and H₂ gases at 200 to 300 atm and 400C to 600C over a finely divided iron catalyst to produce NH₃
Pure elemental phosphorus is white phosphorus (P₄).  It is highly reactive and combusts with air at room temperature but is unreactive with water.
Red phosphorus is formed when white phosphorus is heated and is much less reactive than white phosphorus.
Phosphorus can expand its valence (outermost) shell to hold more than eight electrons (can store extra electrons in the 3d orbitals).
NN triple bonds are much stronger than PP.
P-P single bonds are stronger than N-N single bonds.
As (arsenic) and Sb (antimony) - semimetals
Bi (bismuth) - metal
Group 16 (VIA)
O (oxygen), S (sulfur), and Se (selenium) - nonmetals
Oxygen is the most abundant element on earth, making up approximately 45% of the earth's crust (by weight), 85% of the oceans (by weight) and 20% of the atmosphere (by volume).
Oxygen is generally diatomic (O₂) in its elemental form, but ozone (O₃) is an allotrope of O₂.
At concentrations above 1 ppm, ozone is toxic.
Ozone can absorb ultraviolet (UV) radiation from the sun, and serves as a filter in the atmosphere.
Oxygen is a very strong oxidizing agent, weaker only to fluorine.
Oxygen generally takes on a -2 oxidation number in compounds.
In peroxide (O₂^2-), oxygen has a -1 oxidation number, e.g. H₂O₂, hydrogen peroxide.
Elemental sulfur is a yellow solid at room temperature with a cyclical molecular structure (S₈).
O=O double bonds are much stronger than S=S double bonds.
S-S single bonds are stronger than O-O bonds.
Sulfur can expand its valence (outermost) shell, to hold more than eight electrons (can store extra electrons in the 3d orbitals).
In compounds, sulfur can have oxidation numbers ranging from -2 to +6.
The prefix thio- is given to compounds in which an S atom replaces an O atom.
Te (tellurium) and Po (polonium) - semimetals
Group 17 (VIIA) - Halogens
F (fluorine), Cl (chlorine), Br (bromine), I (iodine), and At (astatine)
All are nonmetals except for At which is a semimetal
All are diatomic in their elemental form
All gain one electron to form anions with a charge of -1.
e.g.  F^-, Cl^-, and Br^-
In compounds, their oxidation numbers range from -1 to +7.
None are found in nature in their elemental forms; instead they are found as salts of the halide ions.
Properties (at room temperature)
F₂ - highly toxic, colorless gas, most reactive element known.
So reactive that it can even form compounds with noble gases (once thought to be inert)
Used in manufacturing Teflon, (C₂F₄)_n
Used to make freons which are used in refrigerators
Cl₂ - highly toxic, pale yellow-green gas, strong oxidizing agent
Used commercially as a bleaching agent an as a disinfectant
Br₂ - reddish-orange liquid with a bad odor
Its name, bromine, is derived from the Greek bromos which means "stench"
Used in preparing fire-extinguishing agents, sedatives, insecticides, and antiknock agents for gasoline
I - deep purple solid with a metallic-like luster
Sublimes directly into a violet gas (I₂) from the solid phase when heated without passing through the liquid phase
Used as a disinfectant, catalysts, drugs, and dyes
AgI (silver iodide) is used in photography
Iodine deficiency in the human body can lead to a goiter, a swelling of the thyroid gland.
Group 18 (VIIIA) - Noble (Rare) Gases
He (helium - "sun"), Ne (neon - "new one"), Ar (argon - "lazy one"), Kr (krypton - "hidden one"), Xe (xenon - "stranger"), and Rn (radon)
Mistakenly labeled as "inert gases" until about 30 years ago because it was thought that these gases did not react with anything.
In 1962, Neil Barlett isolated the first compound containing a noble gas: [Xe⁺][PtF₆^-]
Since then, compounds containing Kr, Xe, and Rn have been isolated, but none containing He, Ne, or Ar.
The oxidation numbers of the rare gases in compounds include +1, +2, +4, +6, and +8.
Noble gases have filled valence (outermost) shells.
Check out periodic tables with the following properties.
Atomic Radii
Atomic Volumes
Atomic Weights
Boiling Points
Covalent Radii
Densities at 300 K
Electrical Conductivities
Electronegativities
First Ionization Energies
Heat of Fusion
Heat of Vaporization
Melting Points
Specific Heat Capacities
Thermal Conductivities
To learn more about individual elements, check out these cool periodic table sites.
Interactive Periodic Table of Elements
Displays information for elements, including atomic radii, electronegativities, electron configurations, densities, melting and boiling points, common oxidation numbers and even who discovered the element.
Click on an element in the table to view a picture.
WebElements Periodic Table
Includes brief descriptions and pictures of all the elements.
Displays images of electron distribution of the elements.
Even includes some movies and comics!
Dynamic Periodic Table
Select which properties you want displayed.
Click an element to view properties, picture, and isotopes (with decay processes)!
Graph trends of properties between any range of elements.
Also has links to alternate styles of periodic tables.
Next:  "Periodic Trends"