Trends of the Periodic Table

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Trends of the Periodic Table
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Review!

• Periodic Table was first organized by
• Dmitri Mendeleev in the mid 1800s
• Mendeleev organized the elements by chemical
  reaction in rows, then by atomic mass in columns
• Henry Moseley then took Mendeleevs table, kept
  the chemical reactivities together, but placed
  them in columns instead. He also ordered the
  elements by increasing atomic number in rows.
• When Moseley did this, all the periodic trends
  just fell into place.
• Remember columns groups/families, rows
  periods

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• Periodic Trends

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Electrons

• Electrons do not freely float in space
• Orbit around nucleus Electron shells
• Each shell corresponds to an amount of energy.

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Valence Electrons

• The valence electrons are the outermost electrons
  of an atom.
• The valence electrons determine the chemical
  properties
• Number of valence electrons equals the column
  number in the A columns
• Elements with the same number of valence
  electrons are very similar chemically
• Alkali metals in Group 1A 1 valence
  electron
• Li, Na, K, Rb, Cs
• Halogens in Group 7A 7 valence electrons
• F, Cl, Br, I

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Atomic Radius

• What is Atomic Radii?
• Distance from the nucleus to the outermost level
  of e- (aka the valence shell)
• What trend do you see as you go across (left to
  right) the period?
• Atomic radius decreases
• Down the group?
• Atomic Radius increases
• WHY???

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Explaining the Trend

• As you go L to R, the atomic radius decreases
  because as you go L to R, the amount of
  attraction between p and e- increase.
• More attractions smaller atomic radius
• As you go down a column, atomic radius increases
  because the e- are farther away from the nucleus.
  There are weaker attractions.
• Weaker attractions larger atomic radius

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Electronegativity

• What is Electro-negativity?
• An atoms Luuuvvv for electrons!
• The tendency to attract another atoms electrons
• What trend do you see as you go across the
  period?
• Electronegativity increases!
• Down the group?
• Electronegativity decreases!
• WHY???

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Explaining the Trend

• As you go L to R, electronegativity increases
  because of the increase in protons. The more
  protons, the more able it will be to attract
  other atoms electrons.
• More attractions (small radius) large
  electronegativity
• As you move down a column, electronegativity
  decreases because of the increase in number
  electron an atoms already has. This means the
  atom will be less able to attract another atoms
  electrons.
• Less attractions (large radius) small
  electronegativity

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Ionization Energy

• What is Ionization Energy?
• The energy needed to remove an electron
• What trend do you see as you go across the
  period?
• Ionization E increases
• Down the Group?
• Ionization E decreases
• WHY???

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Explaining the Trend

• As you go L to R, the ionization energy increases
  because of the increase in the number of protons.
  The more protons, the more energy that is needed
  to remove an electron.
• More attractions (small radius) large
  ionization energy
• As you go down a column, the ionization energy
  decreases because of the decrease in attractions.
  
• Due to electron shielding
• More electrons, leads to outer electrons less
  tightly held.
• The less attractions, the lower the energy that
  is needed to remove an electron.
• Less attractions (large radius) small
  ionization energy

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Ionization Energy

• Amount of energy required to remove an electron
  from the ground state of a gaseous atom or ion.
• First ionization energy is that energy required
  to remove first electron.
• Second ionization energy is that energy required
  to remove second electron, etc.

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Ionization Energy

• It requires more energy to remove each successive
  electron.
• When all valence electrons have been removed, the
  ionization energy takes a quantum leap.

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Electron Affinity

• What is Electron Affinity?
• The energy needed to add an electron
• As you go across the period electron affinity
  increases .
• Electron affinity decreases down the family
• WHY???

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Explaining the trend

• As you go L to R, the electron affinity increases
  because of the increase in the number of protons.
  The more protons, the greater the attraction the
  protons have for electrons.
• More attractions (small radius) large electron
  affinity
• As you go down a family, the electron affinity
  decreases because of the decrease in attractions.
  
• Due to electron shielding
• More electrons, leads to outer electrons less
  tightly held.
• The less attractions, the lower the electron
  affinity
• Less attractions (large radius) small electron
  affinity

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Homework

• Worksheet(s)