Chapter 4: The Periodic Table

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Chapter 4 The Periodic Table
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History

• Late 1820s- Johann Dobereiner noted a
  relationship between the properties of certain
  elements and their atomic masses
• Dobereiners triads Cl, I, Br
• S, Se, Te
• Ca, Sr, Ba
• Li, Na, K

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History

• 1864- John Newlands- every eighth element has
  similar properties
• (inert gases were not yet discovered)

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History

• 1869- Dmitri Mendeleev- developed a table that
  arranged elements in order of increasing atomic
  mass
• He left blank spaces for elements not yet
  discovered

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History

• 1914- Henry Mosley- arranged the table of
  elements into order of increasing atomic number
• Elements with similar properties were placed in
  the same group

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Periodic Law

• The chemical and physical properties of the
  elements are periodic functions of their atomic
  number

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Metals

• Left side of table
• Form positive ions because they lose electrons to
  gain stable electron configurations
• End up with a smaller radius when they become
  ions

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Group 1

• Alkali metals
• Form ions with a 1 charge
• Very active most active metals
• Not found free in nature

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Group 2

• Alkaline earth metals
• Form ions with a 2 charge
• Very active, but not as active as group 1

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Groups 3-12

• Transition elements
• Produce variable charges when ions are formed
• Colored solutions can be formed when transition
  elements are present

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Rare earth elements

• 2 rows at the bottom of the table
• Most are unstable and are associated with
  radioactivity

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Nonmetals

• Right side of the periodic table
• Most are gases
• Gain electrons and form an ion with a negative
  charge

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Group 17

• Halogens
• Form a 1 ion
• Only group that has a solid, liquid, and a gas
• Most active nonmetals
• Fluorine is the most active nonmetal

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Group 18

• Inert gases or Noble gases
• Have completed outer shells
• Stable elements
• Unreactive
• Last elements discovered

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Metalloids

• Found along the step line or zigzag line on the
  periodic table
• Have properties that resemble both metals and
  nonmetals
• Boron, silicon, germanium, arsenic, antimony,
  tellurium, polonium, and astatine
• Exception aluminum is a metal, not a
  metalloid!!!

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Ionization Energy

• Energy required to remove an electron
• The closer an electron is to the nucleus the
  higher the ionization energy
• Table S lists ionization energies of elements

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Electronegativity

• Attraction an atom has for an electron
• Nonmetals have higher electronegativities because
  they want to attract electrons to complete their
  shell
• The farther the electron from the nucleus, the
  weaker the attractive force
• Fluorine has the highest electronegativity (4)
• Table S lists electronegativities

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Trends from top to bottom down a group

• Metallic properties increase
• Radius increases
• Ionization energy decreases
• Electronegativity decreases
• Oxidation state remains the same
• Similar properties

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Trends from left to right across a period

• Metallic properties decrease
• Radius decreases
• Ionization energy increases
• Electronegativity increases
• Oxidation states change
• Properties are different

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