Periodicity

Dalton's Billiard Ball Model

1803

Composed of extremely small particles called atoms

Atoms from same element are identical, but different from other element. Compounds form by combining atoms.

Thomson's Plum Pudding Model

1904

Atom is made up of negative electrons that float in a sphere of positive charge like plums in a pudding

1897

JJ Thomson discovered electron (cathode ray experiment).

1913

JJ Thomson discovered isotopes.

Rutherford's Nuclear Model

1911

Discovered the nucleus of a gold atom with his "gold foil" experiment.

Small, dense center with a positive charge, electrons in a fixed orbit

Bohr's Planetary Model

1913

Nucleus surrounded by orbiting electrons at different energy levels, electrons have definite orbits

Lothar Meyer

Germany, 1869

Similar chemical and physical properties recur PERIODICALLY when the elements arranged in order of INCREASING ATOMIC WEIGHT

Dmitri Mendeleev

Russia, elements with similar characteristics be listed in the same COLUMN forced him to leave blank spaces in his table

Periodic Law

The properties of the elements are periodic functions of their atomic masses.

Henry Moseley

1913

atomic number as the number of protons in the nucleus of the atom

frequency increase as atomic mass increased

By bombarding different elements with high energy electrons, each element produced X-rays of unique frequency and that the frequency _________.

Atomic number

number of protons in nucleus which determines the identity of the element

Mass Number (Atomic Mass)

number of protons plus neutrons, unit is g/mol

Neutrons

Mass Number - Atomic Number

Isotopes

atoms of the same element with varying number of neutrons

Ion Charge

Protons - Electrons

Electrons

atomic number - charge

Johann Wolfgang Dobereiner

1829

Law of Triad, where elements are arranged in groups of 3's which are alike in many properties

John Newlands

1864

Law of Octaves, arranged the elements in groups of 8's and have similar properties and are seven elements apart

Periods

seven HORIZONTAL rows in the periodic table

2 electrons in s sublevel

Period 1 has 2 elements corresponding to...

8 electrons in s and p sublevels

Period 2 and 3 has 8 elements corresponding to...

18 electrons in s, p, and d sublevels

Period 4 and 5 has 18 elements corresponding to...

32 electrons in s, p, d, f sublevels

Period 6 has 32 elements corresponding to...

Period 7

still incomplete but elements fill up s, p, d, and f sublevels

Groups or Families

vertical columns in the periodic table, which are divided into A and B subgroups

Alkali Metals

Group 1A

Alkaline Earth Metals

Group 2A

Halogens

Group 7A

Noble Gases

Group 8A

Boron Group

Group 3A

Carbon Family

Group 4A

Nitrogen Family

Group 5A

Oxygen Family

Group 6A

Helium

most ideal gas element

Coulomb's Law

attractive force between an electron and the nucleus depends on the magnitude of the nuclear charge and on the average distance between the nucleus and the electron.

magnitude of nuclear charge (Z)

more protons = stronger attraction

weaker attraction

Farther electrons feel...

nuclear charge, distance

Force increases with higher ______, and decreases with _______.

effective nuclear charge (Zeff)

The net positive charge experienced by an electron in a multi-electron atom.

outer electrons

Since inner/core electrons shield outer electrons from the nucleus, the ________ don't feel the full nuclear charge.

Z

atomic number (total nuclear charge)

S

screening constant (approx. number of core electrons)

less than

Zeff is always _____ Z because shielding reduces the effective pull of the nucleus.

van der Waals radius (nonbonding atomic radius)

the radius of an atom when it is not bonded to another atom

Covalent Bonding Radius

half of the nucleus-to-nucleus distance, d

Atomic Size (Atomic Radii)

the average distance between the nucleus and valence electron

ionic radius

Distance from the center of an ion's nucleus to its outermost electron

Metals

Lose e, which means more p than e (more attraction)

Nonmetals

gain e, more e than p (not much attraction)

Cation radius

less than neutral atomic radius

Anion radius

greater than neutral atomic radius

Isoelectronic series

group of ions all containing the same number of electrons

Metallic property

ability of the atom to donate electrons

reactivity

tendency of an atom to react

Ionization energy

amount required to remove an electron from an atom or ion

electron affinity

energy change when an electron is accepted by gaseous atom to form anion

Electronegativity

defined as the relative ability of an atom of an element to attract or gain electrons

increasing (top to bottom, right to left)

atomic size, metallic property, reactivity

decreasing (top to bottom, right to left)

ionization energy, electron affinity, electronegativity