Unit 3 – The Periodic Table, Nomenclature and Bonding
Classification of Matter
History of the Periodic Table
Dobereiner’s Triads: (1817) Groups of three elements with similar properties
Cannizzaro’s method for determining atomic mass: (1860) Was adopted as the standard for determining atomic mass at the First
International Congress of Chemists in Karlsruhe, Germany
Newland’s Law of Octaves: (1863) When arranged by atomic mass, every 8
element had similar properties
Mendeleev’s Periodic Table: (1869) Arranged elements in a table by atomic mass.
Elements with similar properties were
grouped in columns.
Mendeleev left empty spaces for elements that were “undiscovered”.
He used his table to predict
the properties of these elements (Sc, Ga, Ge) right!
Why did some elements (Ar & K, Co & Ni, I & Te) not follow the
atomic mass pattern?
Why were the properties periodic?
The Modern Periodic Table
Moseley’s Atomic Number: (1911) led to the arrangement of the periodic table by increasing atomic number; retained columns
with similar properties
Modern Periodic Law: The properties of the elements are periodic functions of their atomic numbers
Arrangement of the Modern Periodic Table
Physical Property: a characteristic of a substance that can be observed or measured without changing the identity of the
Chemical Property: a description of a substance’s ability to react (or not react).
Changes into a new substance as a result of the
observation or measurement.
General Properties of the Elements
Metals: left of the staircase and Al
Good conductors of heat and electricity, malleable, ductile, shiny, high melting points, mostly solids (Hg is the
exception), few valence e-
Nonmetals: right of the staircase and H
Poor conductors of heat and electricity, are gases or dull, brittle solids, low melting points, 5-7 valence e-
Metalloids (semimetals): touching the staircase except Al
Solids, have some properties of metals and nonmetals, semiconductors
Families or Groups (columns)
Alkali Metals – Group I
most reactive metals, not found free in nature, silvery, soft, react vigorously with H
O, lose e- in reactions,
general valence structure ns
Alkaline Earth Metals – Group 2
reactive, not found free in nature, harder, denser, stronger than Group I, lose e- in reactions, general valence
Halogens – Group 17
most reactive nonmetals, “salt formers”, gases (F, Cl), liquid (Br), solids (I, At), gain e- in reactions, general
valence structure ns
Noble Gases – Group 18
stable, unreactive elements (although some can form compounds), discovered between 1894-1900, general
valence structure ns
, He is the exception
Transition elements (d block) – metallic properties, less reactive, harder & stronger than s-block, sum of the outer “s” and “d”
electrons gives the group number, variable oxidation numbers (charges), often form colored ions, deviation in e-