Periodic Table

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Periodic Table
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Periodic Table

• The periodic table is a systematic arrangement of
  the elements by atomic number (protons)
• Elements with similar properties fall into
  vertical columns

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History of the Periodic Table

• Three men recognized patterns in the elements.
  They attempted to organize the elements according
  to these patterns..

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History of the Periodic Table

• Johann Wolfgang Döbereiner
• Noticed patterns in atomic mass recurring in sets
  of three elements
• Became known as Döbereiner's triads

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History of the Periodic Table

• John Newlands
• noticed every eighth element had similar
  properties.
• Known as 'law of octaves'

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History of the Periodic Table

• Dmitri Mendeleev
• developed the first periodic table
• Found the repeating pattern by atomic mass and
  arranged them so that groups of elements with
  similar properties fell into vertical columns in
  his table.
• Found a problem
• Some elements fell into the wrong column
• Examples Te I Co Ni

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Mendeleevs Periodic Table
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History of the Periodic Table

• Henry Moseley Fixed Mendeleevs problem by
  rearranging the modern table by atomic number
• Used X-ray spectrometer to find the atomic numbers

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Arrangement of Periodic Table

• Periodicity
• trends of properties
• as you go across the table
• or down a column

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Periods

• horizontal rows
• there are 7
• Period number tells which energy level holds the
  valence electrons

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2
3
4
5
6
7
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Groups/Families

• vertical columns
• groups 1-18
• elements in the same group share chemical
  properties
• Main group elements
• Groups 1,2 13, 14, 15, 16, 17,
  18

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Types of Elements
Noble gases
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Metals

• Found on the left side of table
• Have 1, 2 or 3 valence electrons
• Lose electrons to form positive ions (cations)
• Most are silver, shiny, malleable, ductile good
  heat/electrical conductors, solid

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Nonmetals

• Found on the right side of table
• Have 5, 6, or 7 valence electrons
• Gain electrons to form negative ions (anions)
• Brittle, dull, non-conductors, and exist in all
  three states (solids, liquids, gases)

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Metalloids

• Elements found along the stair-step between
  metals and nonmetals, NOT Al
• Properties are in between metals nonmetals
• Silicon (Si) is probably the most well-known
  metalloid.

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Noble Gases

• odorless,
• colorless, 
• monatomic gases 
• low chemical reactivity.

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Color Groups of the Periodic Table
Metalloids
Noble Gases
Alkali Metals
Halogens
Alkaline Earth Metals
Transition Metals
Also called inert gases because they do not react
Inner Transitional Metals
Lanthanide Series
Actinide Series
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Properties and Electron Configuration

• All elements in a group (column) end with the
  same electron configuration. That determines many
  of the physical properties that the group share.

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Group 1

• Based on the video Alkali Metals with Water
• What properties of Alkali metals are observed?
• What trend is observed as samples are tested with
  water?

• Why werent hydrogen
• and francium tested?

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Group 17

1. What are some of the physical properties of the
   halogens? (Watch video and list some physical
   properties of halogens.)

Halogen
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Group 18

• Note In the video Group 0 is an old name for
  Group 18.
• Why are the noble gases un-reactive?
• If all neon signs were made of pure neon gas,
  what colors would we have?
• What are uses for noble gases other than in neon
  lights?
• How can a physical property be used to tell the
  difference between noble gases?
• Radon was not tested. Predict what a balloon
  filled with Radon would do when dropped from the
  roof and why.

Noble Gases
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Summary of Groups, Props. Electrons

• NOVA Video
• What is the relationship between electron
  configuration and group number on the periodic
  table?
• Why are halogens and alkali metals highly
  reactive, but not the noble gases?

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Periodic Table Trends

• Patterns on the periodic table
• Atomic Radius
• Ionic Radius
• Electronegativity
• Ionization Energy
• Reactivity

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Periodic Trends- similarities of elements based
on where they are in the table

• Depend on two things
• Effective Nuclear Charge (Zeff) -
• The attraction the valence
• electrons have for the protons
• in the nucleus.
• Electron Shielding Effect-
• Inner shell electrons 
• blocking valence electrons from the nucleus.

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Effective Nuclear Charge
Effective Nuclear Charge is abbreviated
Zeff Smart folks have noticed that the Zeff for
each group is equal to the number of valence
electrons.
Watch this video
And this
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Atomic Radius

• Atomic radius is half the distance between the
  centers of two atoms measured in angstroms.

larger
The more energy levels, the ________ the atomic
radius. (larger/smaller) The higher the
effective nuclear charge , the ________ the
atomic radius. (larger/smaller)
smaller
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Atomic Radius Trend

• Atomic radius increases as you move down a group
• Atomic radius decreases as you move from left to
  right in a period

Across the period the number of protons increases
while the number of shielding electrons stays the
same. This make the nucleus pull in the valence
electrons. That makes a smaller atom.
Down the group the number of energy levels
increase so the number of shielding electrons
increase. The nucleus cannot pull in the valence
electrons. That makes a bigger atom.
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Ions

• Anions
• Form from nonmetals
• Gain electrons
• Nonmetals have high effective nuclear attraction
  on the valence electrons

• Cations
• Form from metals
• Lose electrons
• Metal have low effective nuclear charge holding
  on to the valence electrons.

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Electronegativity

• Electronegativity is a measure of how strongly
  atoms attract bonding electrons to themselves
• An assigned number rates the electronegativity
  (from 0.7 to 4.0)
• Low electronegativity cannot attract valence
  electrons
• High electronegativity can attract valence
  electrons

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Electronegativity Trend

• Electronegativity values increase as you move
  from left to right in any period.
• Group 18 (noble gases) have EN 0.
• Within any group, electronegativity values
  inrease as you go up.

Biggest IE Fluorine Smallest IE Francium
Electronegativity
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Ionization Energy

• Ionization Energy the energy needed to remove
  the outermost electron in an atom. How hard is
  it to steal an electron
• Increases as you go across a period
• Larger nuclear charge more protons pulling on
  the electrons
• Atom is smaller outer electrons are closer to
  the nucleus easier to pull in electrons
• Increases as you go up in a group
• less energy levels Radius is smaller outer
  electrons are closer to the nucleus and more
  strongly attracted to nucleus

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Ionization Energy Pattern
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Ionization Energy
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Metal (Re)activity Trend

• Metal Activity depends on the attraction the
  metal has for the nonmetals electrons.

• Trend
• Increases as you move down a group
• Increases as you move from right to left in a
  period

The most reactive metal is francium
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Nonmetal Activity Trend

• Non-Metal Activity refers to how easily nonmetals
  gain e- to form anions

increasing nonmetal activity

• Trend
• Increases as you move up a group
• Increases as you move from left to right in a
  period

increasing nonmetal activity
The most reactive nonmetal is fluorine