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

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)

1. B (boron) - only element in the group that is not a metal; has semimetal and nonmetal characteristics.
2. 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
3. Other metals - Ga (gallium), In (indium), and Tl (thallium) - very scarce active metals

Group 14 (IVA)

1. 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
2. S (silicon) and Ge (germanium) - semimetals
   • Silicon is the second most abundant element in the earth's crust
3. 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)

1. 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.
2. As (arsenic) and Sb (antimony) - semimetals
3. Bi (bismuth) - metal

Group 16 (VIA)

1. 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.
2. 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!
• 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"