Presentation on theme: "THE PERIODIC TABLE BRIEF HISTORY. Dmitri Mendeleev (1869, Russian) –Organized elements by increasing atomic mass. –Elements with similar properties were."— Presentation transcript:
Dmitri Mendeleev (1869, Russian) –Organized elements by increasing atomic mass. –Elements with similar properties were grouped together. –There were some discrepancies. –Predicted properties of undiscovered elements.
Henry Mosely (1913, British) Organized elements by increasing atomic number. Resolved discrepancies in Mendeleev’s arrangement.
I. ORGANIZATION of ELEMENTS Three major categories: -Metals -Non-metals -Metalloids
METALS Shiny Lustrous Various colors, but most are silvery Good conductors of heat & electricity Most are solids at room temperature that are: –Ductile –Malleable Mercury is a liquid at room temperature. NON-METALS Dull, not shiny. Various colors Poor conductors of heat & electricity (good insulators) Are solids, liquids & gases. Liquids: Br 2 Gases: Noble gases, H 2, N 2, O 2, F 2, and Cl 2 Remaining elements are solids.
METALLOIDS Elements that border the “staircase” on the periodic table. Have properties of both metals & non-metals. Their behavior depends on what they are bonded to chemically. –They behave like a metal when bonded to a non- metal. –They behave like a non-metal when bonded to a metal.
II. Predicting Oxidation Numbers Oxidation State: -# of electrons an atom will gain or lose to become stable. Metals always have positive (+) oxidation states. Nonmetals have negative (-) oxidation states.
s – Block Metals GROUP 1 Alkali Metals -have 1 valence electron and s 1 configuration. -lose 1 electron to become stable. -have oxidation number of +1 (charge) GROUP 2 Alkaline Earth Metals -have 2 valence electrons and s 2 configuration. -lose 2 valence electrons to become stable. -have oxidation number of +2 (charge)
p-Block Metals Must lose electrons; have 2 possibilities: 1.Remove ALL valence electrons 2.Remove only “p” valence electrons (all at once)
d – Block Metals Must lose electrons; have several possibilities: 1.FIRST, remove ALL valence electrons 2.Then, remove “d” electrons, one at a time (until stable)
Non-Metals Non-metals have only ONE choice: Will gain enough e - to make 8 valence e -.
METALLOIDS Have properties of BOTH metals & non- metals. 1.Treat like a “p” block metal (lose e - ) 2.Treat like a non-metal (gain e - )
III. Periodic Trends Anything that influences the valence electrons will affect the chemistry of the element. 1.Nuclear Charge 2.Energy Levels / # of core (inner) electrons.
1. Nuclear Charge (# of protons): A larger nuclear charge means a smaller outer level. b/c the higher positive charge of the nucleus pulls the valence e - inwards. 2. Energy Levels: Additional energy levels increase the distance between the nucleus and the valence e - Thus…the atom has more volume and a bigger radius.
3. Number of Core (inner) Electrons: More core electrons means a larger valence shell (aka atomic radius). Core e - repel the valence e - and push them farther away from the nucleus.
NUCLEAR CHARGE The # of protons in the nucleus Increases from L to R, across a period. Increases from top to bottom, down a group.
ATOMIC RADIUS Distance from the nucleus to the valence electrons. DECREASES from L to R across a period due to increasing nuclear charge. INCREASES from top to bottom down a group b/c of increased number of E levels (shielding).
Ionization Energy E required to remove an e - from an atom. INCREASES L to R across a period b/c of increasing nuclear charge. DECREASES from top to bottom down a group b/c of increased number of E levels.
Electron Affinity Ability to attract an e - (to form a anion). INCREASES L to R across a period b/c of increasing nuclear charge. DECREASES down a group due to the increase in E levels. NOTE: Noble gasses have NO electron affinity!
Ionic Radius Dist. from the nucleus to the valence e - Cations: Ionic radius is smaller than atomic radius b/c the atom has lost e - (smaller cloud) Anions: Ionic radius is larger than atomic radius b/c the atom has gained e - (larger cloud)
Electronegativity Ability to attract an e - in a chem compound INCREASES L to R across a period b/c of increasing nuclear charge. DECREASES (or stays about the same) from top to bottom down a group due to the increase in E levels.