 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: