Trends

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Trends the Periodic Table
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Trends

• more than 20 properties change in predictable way
  based location of elements on PT
• some properties - anyone know where we can find
  these numbers?!
• Density
• melting point/boiling point
• atomic radius
• ionization energy
• electronegativity

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When youre done it will look like this so leave
room for writing!
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Going down column 1
increasing energy levels as go down
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Increasing number of energy levels
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Atomic Radius

• Atomic radius defined as ½ distance between
  neighboring nuclei in molecule or crystal
• Affected by
• 1. of energy
• levels
• 2. Proton Pulling
• Power

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Increasing number of energy levels
Increasing Atomic Radius
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Cs has more energy levels, so its bigger

Li Group 1 Period 2 Cs Group 1
Period 6
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As we go across, elements gain electrons, but
they are getting smaller!
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Decreasing Atomic Radius
Increasing number of energy levels
Increasing Atomic Radius
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Why does this happen..

• As you go from left to right, you again more
  protons (the atomic number increases)
• You have greater proton pulling power
• Remember the nucleus is and the electrons are -
  so they get pulled towards the nucleus
• The more protons your have, the more Proton
  Pulling Power

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as go across row size tends to decrease a bit
because of greater PPP proton pulling power
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We can measure the Proton Pulling Power by
determining the Effective nuclear charge

• It is the charge actually felt by valence
  electrons
• The equation
• Nuclear charge - inner shell electrons
• (doesnt include valance e-)

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1

• Calculate effective nuclear charge
• protons minus inner electrons

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What the inner electrons do. They Shield the
charge felt by the valance electrons.
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H and He only elements whose valence electrons
feel full nuclear charge (pull) NOTHING TO
SHIELD THEM
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Decreasing Atomic Radius
Increasing number of energy levels
Increasing Atomic Radius
Increased Electron Shielding
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Look at all the shielding Francium's one valance
electron has. It barely feels the proton pull
from the nucleus. No wonder it will lose its one
electron the easiest. No wonder its the most
reactive metal
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Ionization Energy

• amount energy required to remove a valence
  electron from an atom in gas phase
• 1st ionization energy energy required to remove
  the most loosely held valence electron (e-
  farthest from nucleus)

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• Cs valence electron lot farther away from nucleus
  than Li
• electrostatic attraction much weaker so easier to
  steal electron away from Cs
• THEREFORE, Li has a higher Ionization energy then
  Cs

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Increased Ionization Energy (harder to remove an
electron)
Decreasing Atomic Radius
Increased Electron Shielding
Increasing number of energy levels
Increasing Atomic Radius
Decreased Ionization Energy (easier to remove an
electron)
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Electronegativity

• ability of atom to attract electrons in bond
• noble gases tend not to form bonds, so dont have
  electronegativity values
• Unit Pauling
• Fluorine most electronegative element
• 4.0 Paulings

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Increased Electronegativity
Increased Ionization Energy (harder to remove an
electron)
Decreasing Atomic Radius
Increased Electron Shielding
Increasing number of energy levels
Decreased Ionization Energy (easier to remove an
electron)
Increasing Atomic Radius
Decreased Electronegativity
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Reactivity of Metals

• judge reactivity of metals by how easily give up
  electrons (theyre losers)

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Increased Electronegativity
Increased Ionization Energy (harder to remove an
electron)
Decreasing Atomic Radius
More metallic
Increased Electron Shielding
Increasing number of energy levels
Decreased Ionization Energy (easier to remove an
electron)
Decreased Electronegativity
Increasing Atomic Radius
Most reactive metal Fr (the most metallic)
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Reactivity of Non-metals

• judge reactivity of non-metals by how easily gain
  electrons (they are winners)

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Increased Electronegativity
Increased Ionization Energy (harder to remove an
electron)
Decreasing Atomic Radius
Most Reactive Nonmetal F
More metallic
Increased Electron Shielding
Increasing number of energy levels
Decreased Ionization Energy (easier to remove an
electron)
Decreased Electronegativity
Increasing Atomic Radius
Most reactive metal Fr (the most metallic)
Nonreactive
BACK
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How do you know if an atom gains or loses
electrons?

• Think back to the Lewis structures of ions
• Atoms form ions to get a valence of 8
• (or 2 for H)
• Metals tend to have 1, 2, or 3 valence electrons
• Its easier to lose them
• Nonmetals tend to have 5, 6, or 7 valence
  electrons
• Its easier to add some
• Noble gases already have 8 so they dont form
  ions very easily

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Positive ions (cations)

• Formed by loss of electrons
• Cations always smaller than parent atom

2e
8e
8e
8e
8e
2e
2e
Ca
Ca
Ca2
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Negative ions or (anions)

• Formed by gain of electrons
• Anions always larger than parent atom

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Allotropes

• Different forms of element in same phase
• different structures and properties
• O2 and O3 - both gas phase
• O2 (oxygen) - necessary for life
• O3 (ozone) - toxic to life
• Graphite, diamond
• both carbon in solid form