     Periodic table is arranged according to increasing atomic number by Henry Moseley.

Originally arranged by order of increasing atomic mass by Demetri Mendelev.

The word “periodic” means pattern. Chemical and physical of atoms properties are a function of atomic number – The Periodic Law.

Groups: Periods:

    Electron donors (in ionic bonds) Will only be found in ionic bonds not covalent Conduct heat and electricity Metals are found to the far left of the periodic table(before staircase) – the further left the more metallic in nature.

    Electron acceptors (negative charge in ionic bonds) Insulators Non metals are found to the far right of the periodic table.

Non –metal to non-metal bonding is covalent.

         Group 1- Alkali Earth Metals Group2- Alkaline Earth Metals Group3-12 – Transition Metals  Post Transition Metals Group 13- Boron Group Group 14- Carbon Group Group15-Nitrogen Group Group 16- Oxygen group Group 17- Halogen group Group 18- Noble gases

      Metalloid- elements with both metallic and nonmetallic properties Inner transition metals- are located in the f subshell. Lanthanide series(4f) and actinide series( 5f) Hydrogen - contains the element hydrogen and has a +1 and -1 charge Noble gases are inactive gases Halogen group- is the most active non metal group Alkali Metals- are the most active metal group.

  Also known as the rule of 8:  Atoms want the outer most shell to be completely filled.

  For the s and p blocks this means a valence of 8 (or 2) electrons.

They will react in a way either ionically or covalently to acquire that valence of 8.

In nature because of the octet rule you never find elements by themselves only in compounds except for noble gases. These are called diatomic elements H 2, N 2 , O 2 , F 2, Cl 2 , Br 2 , I 2

         Group1: +1 Group2: +2 Group13: +3 Group 14: +4 or – 4 Group 15: -3 Group 16: -2 Group 17: -1 Group 18: 0 Transition (and Post-Transition) Metals – will have variable oxidation states. Not as predictable as the main group elements (know the common ones Fe, Cu, Cr, Pb, Sn).

    Valence electrons- total electrons in the outer most shell Core electrons- everything else except valence electrons.

Groups- goes from top to bottom on the periodic table.

Periods- goes from left to right on the periodic table.

    As you add more protons the pull of the nucleus gets stronger as felt by the valence electrons.

Effective nuclear charge increases as you go across the periodic table from left to right.

It decreases as you go down the periodic table from top to bottom because of the increase in valence electron distance and electron shielding.

Example: Let’s compare Li and F

   The more core electrons there are the less of a pull the valence electrons will feel from the nucleus.

Shielding stays constant as you go across the periodic table (left to right) and increases as you down the periodic table.

Example: Let’s compare Li and Cs

    Size of the atom.

As you go from left to right the atomic radius decreases because of the stronger effective nuclear charge.

As you go from top to bottom the atomic radius increases because of the shielding & valence distance.

As you go down the group you are adding more energy levels and this increases the size of the atom.

    Energy needed to remove the outermost electron.

As you go from left to right on the periodic table the ionization energy increases because of stronger effective nuclear charge.

As you go from top to bottom the ionization energy decreases because of more electron shielding & an increase in the distance of the valence eletctrons.

Example: Put in order of increasing ionization energy, Be, Li, & F

    Measurement of the ability of an atom to attarct an electron.

As you go from top to bottom the electronegativity decreases because of shielding and distance of valence.

As you go from left to right the electronegativity increases because of effective nuclear charge.

Example: