THE PERIODIC TABLE

BRIEF HISTORY

• 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: Br2

• Gases: Noble gases, H2, N2, O2, F2, and Cl2

• 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 s1 configuration.

-lose 1 electron to become stable.

-have oxidation number of +1 (charge)

• GROUP 2 Alkaline Earth Metals

-have 2 valence electrons and s2 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.