1. Formation of the Periodic Table Mendeleev: arranged the periodic table in order of increasing atomic MASS (didn’t know about protons) –Started new rows when properties were similar –Very similar to today’s, but was turned 90 ° Moseley: rearranged it in order of increasing atomic NUMBER Periodic law: physical and chemical properties tend to repeat themselves in a systematic manner based on increasing atomic number

2. Organization: Periods: –ROWS on the periodic table –Valence electrons all in same energy level Groups: –COLUMNS on the periodic table –All have similar electron configurations (and number of valence electrons) –Similar properties!

3. Metals/nonmetals/metalloids Metals: –Left side and below stairsteps –React with nonmetals to form salts –Malleable, ductile, conduct electricity, have luster –Solid at room temperature (except Hg) Metalloids –Along stairstep line: B, Si, Ge, As, Sb, Te, Po, At –Have some properties of metals, some of nonmetals Ex: silicon: shiny (has luster) but brittle Nonmetals: –Right side and above stairsteps –Varying states of matter (gases in red, liquids in blue, solids in black) –Varying properties, DON’T conduct electricity, solids NOT shiny, solids brittle

4. Group names/properties Alkali metals: group 1/IA –Highly reactive, form basic solutions when reacted with water Alkaline earth metals: group 2/IIA –Still reactive, but not as much as Grp IA Halogens: group 17/VIIA –React with metals to form salts, most reactive of nonmetals Noble gases: group 18/VIIIA –Nonreactive, have full valence shells

5. Large regions Main block (or representative) elements: –The “s” and “p” blocks Transition elements: –The “d” block Inner transition elements: –The “f” block –Also called the lanthanides (lanthanoids) and the actinides (actinoids

6. Reasons for trends Distance of valence electrons from nucleus: based on which energy level the electrons are in (higher level = farther) Coulomb’s Law: the forces between positive and negative charges, which DECREASE as the distance between them INCREASES Shielding effect: Electrons in inner energy levels “shield” the electrons in the outer energy levels from the pull of the nucleus Nuclear charge: the pull of the nucleus on the electrons (increases as number of protons increases)

7. Patterns (trends) on the Periodic Table Atomic radius: –Half the distance between the nuclei of a diatomic molecule (because we can’t measure to the edge of the electron cloud accurately) –Tells the size of the atom As atomic number increases in a group (top to bottom) INCREASES As atomic number increases in a period (left to right) DECREASES

9. Electron Affinity, Ionization energy and Electronegativity Electron affinity: the ability of an atom to “hold on to” its electrons Ionization energy: The amount of energy required to remove an electron from an atom Electronegativity: The tendency to attract electrons during bonding (ability to “steal” or “keep” the electrons)

10. Electron Affinity, Ionization energy and Electronegativity The GENERAL trend: As atomic number increases in a group (top to bottom) DECREASE As atomic number increases in a period (left to right) INCREASE

13. 1 st ionization energy, 2 nd ionization energy, etc. 1 st : the amount of energy to remove one electron 2 nd : the amount to remove another (a second) electron 3 rd : the amount to remove a third electron Determined by which electron will be removed! Draw orbital diagrams to decide. Remember that full sublevels (or completely ½ full sublevels) are more stable than partially filled Ex: Na –1 st would remove the 3s 1 electron – relatively easy –2 nd would remove the last electron from 2p 6 – much more difficult!

14. Ionic radius periodicTbl2 Ionic radius of a POSITIVE ion (cation) is SMALLER than its parent atom because there are MORE protons than electrons (because it LOST electrons), pulling the electrons in tighter Ionic radius of a NEGATIVE ion (anion) is LARGER than its parent atom because there are more electrons than protons (because it GAINED electrons) and they pull away more

15. Paramagnetic/Diamagnetic atoms Paramagnetic: the atom has at least one unpaired electron Diamagnetic: all electrons are paired within the atom Which groups are the atoms paramagnetic?

16. Lewis Dot diagrams: valence electrons! Examples: