Periodic Table

Chapter 5, Section 1 - Organizing the Elements                                                                                                                 

1869 - Dmitri Mendeleev published the first periodic table.
  • He put elements in order by their atomic masses, and noticed that certain properties repeated themselves.
  • He put the elements in rows, and started a new row when the properties repeated.
  • If an element didn't follow the pattern, he left a space and moved it to a column with similar properties.
    • He predicted that new elements would be discovered to fill these gaps, and even predicted their properties.
      • Germanium, Gallium and Scandium were all predicted by Mendeleev.
The modern periodic table organizes elements by atomic number, not atomic mass.  (This fixed the few problems with Mendeleev's table.)

Periodic Law - When elements are arranged in order of their atomic numbers, elements with similar properties appear at regular 
intervals.  (On the Periodic Table, elements with similar properties are listed in the same column.)

Periods - The horizontal rows on the Periodic Table.
  • Metals are on the left, nonmetals are on the right.
  • Elements get less metallic as your move across the table (left to right).
Groups - The vertical columns on the Periodic Table.
  • Elements in the same group have similar properties.

    Pg. 150 #2-6.

Chapter 5, Section 2 - Exploring the Periodic Table                                                                                                           

The patterns of properties in the periodic table are due to the number and arrangement of valence electrons.
Valence electrons are the electrons in the atom's outermost shell.  They determine the atom's chemical properties.
  • Atoms in the same group have the same number of valence electrons, so they have similar properties.
    • They have different numbers of protons and inner-shell electrons, so their properties are not exactly the same.
An element's location in the periodic table is related to the arrangement of its electrons...
  • Each period represents an energy level.
  • Hydrogen has one e- in its 1s orbital.
  • Helium has two e- in its 1s orbital, filling this energy level.
  • Carbon has 2e- in its 1s orbital, 2e- in its 2s orbital, and 2e- in its 2p orbitals.
Atoms that do not have a full outer orbital can gain or lose electrons so that their outermost orbital is full.  This is called ionization.

Ion - An atom or molecule that has gained or lost electrons and, therefore, has a net positive or negative charge.

Cation - A positive ion.
  • Group 1 elements (lithium, sodium, etc) have one valence electron, which they lose easily to become 1+ ions.
Anion - A negative ion.
  • Group 17 elements (fluorine, chlorine, etc) have seven valence electrons, so they gain one electron to become 1- ions.

Elements can be grouped into three main categories...
  • Metals
    • Good conductors of heat and electricity.
    • Ductile - Can be drawn into a wire.
    • Malleable - Easily shaped and formed.
    • Most are shiny.
    • Solid (except mercury).
  • Nonmetals
    • Insulators - Poor conductors of heat and electricity.
    • Not ductile or malleable - solid nonmetals are usually brittle.
    • Most are not shiny.
    • Can be solid, liquid or gas.
  • Semiconductors
    • Share some properties with metals, some with nonmetals.
    • Conduct electricity under certain conditions.  (Better conductor than nonmetal, not as good as metal.)
    • Also called metalloids.
      • Often used in computer chips and electronic devices.

    Pg. 155 #1, 3-7.

Chapter 5, Section 3 - Families of Elements                                                                                                                       

Groups of elements can be classified into families based on their properties and their valence electrons.

Alkali Metals
  • Group 1 elements (except hydrogen).
  • Soft, shiny, react violently with water.
  • One valence electron - make 1+ ions.
  • Not found as pure elements in nature.
Alkaline-Earth Metals
  • Group 2 elements.
  • Harder, denser and stronger than alkali metals.
    • Higher melting points.
  • Two valence electrons - make 2+ ions.
  • Less reactive than alkali metals, but still make compounds easily.
Transition Metals
  • Groups 3-12 elements.
  • Harder, denser, stronger, and have higher melting points than alkaline-earth metals.
    • Except for mercury!
  • Can use electrons from inner shells to make bonds,...
    • this lets them lose different numbers of electrons and form different kinds of atoms.
    • examples:  Fe+2 / Fe+3,  Au+ / Au+3
Noble Gases
  • Group 18 elements.
  • Their orbitals are full, so they do not gain or lose electrons.
    • Inert - They do not react with other elements (under normal conditions).
  • Exist as single atoms - not molecules.
  • Group 17 elements.
  • Most reactive nonmetals.
  • Have seven valence electrons - only need one to be full!
  • Combine easily with metals to form salts.
Other Nonmetals
  • Carbon, nitrogen, oxygen, phosphorus, sulfur, selenium.
  • Form compounds and/or negative ions.
  • Carbon can form millions of different carbon compounds, because it can bond with up to four other atoms at a time.
  • Boron, silicon, germanium, arsenic, antimony, tellurium.
  • Can conduct heat and electricity under certain conditions.
  • Unique because it has one electron and once empty space in its orbital.
  • About 3/4 of all atoms are hydrogen.

    Pg. 164 #1-5.

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