Trends

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Trends the Periodic Table
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Trends

• more than 20 properties change in predictable
  ways based location of elements on PT
• some properties can be predicted
• density
• melting point/boiling point
• atomic radius
• ionization energy
• electronegativity

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Atomic Radius

• Atomic radius defined as ½ distance between
  neighboring nuclei in molecule or crystal
• size varies a bit from substance to substance

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X-ray diffraction pinpoints nuclei then
measures distance between them Cannot measure
electron cloud
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Trends Atoms get larger as go down column ?
principal energy levels Atoms get smaller as
move across series ? PPP proton pulling power
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Going down column 1
increasing energy levels as go down - makes
sense that atoms get larger in size
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Li Group 1 Period 2 Cs Group 1
Period 6
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Going across row 2
size atoms actually get a bit smaller as go
across row left to right - whats going on?
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What do you remember about charge?

• opposites attract
• like charges repel
• largest influence on atomic size in order
• principal energy levels
• proton pulling power (PPP)

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Effective nuclear charge

• Charge actually felt by valence electrons
• To calculate
• Atomic Number minus inner shell electrons
• Not same as nuclear charge or protons in
  nucleus
• Charge felt by valence electrons is attenuated
  (shielded) by inner shell electrons

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H and He only elements whose valence electrons
feel full nuclear charge (pull)
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Lis valence e- feels effective nuclear charge of
1
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Calculate effective nuclear charge
protons minus inner electrons
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as go across row size tends to decrease a bit
because of greater proton pulling power (PPP)
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• size ? as you go ? column
• size ? as you go ? row

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Ionization Energy

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

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Trends in ionization energy

• What do you think happens to the ionization
  energy as go down column of PT?
• As go across row?

decreases
increases
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• valence e- in atoms effective nuclear charge of
  1
• Cs valence e- lot farther away from nucleus than
  Li
• electrostatic attraction much weaker so easier to
  steal electron away from Cs

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• easier to steal electron from Li than from Ne
• Li smaller effective nuclear charge
• - valence electron farther away from
  nucleus
• Li has less proton pulling power than Ne

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Trends in ionization energy

• Ionization energy decreases as go down column
• easier to remove valence electron
• Ionization energy increases as go across row
• more difficult to remove valence electron

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Periodic properties Graph shows a repetitive
pattern (Note Doesnt have to be a straight
line)
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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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Trends in electronegativity

• Related to PPP
• Increases as go across row
• Decreases as go down column
• Remember F most electronegative element!

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Reactivity of Metals

• metals are losers!
• judge reactivity of metals by how easily give up
  electrons
• most active metals Fr (then Cs)
• For metals, reactivity increases as ionization
  energy goes down

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Trends for Reactivity of Metals or Metallic
Character

• Increases as go down column
• easier to lose electrons!
• Decreases as go across row
• more difficult to lose electrons!

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Reactivity of Non-metals

• Non-metals are winners!
• judge reactivity of non-metals by how easily gain
  electrons
• F most active non-metal
• For non-metals
• reactivity ? as electronegativity ?

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Trend for Reactivity of Non-metalsDepends on
PPP

• Increases as go across row
• Decreases as go down column
• (shielded by more inner-shell electrons)

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Ionic Size Relative to Parent Atom

• Depends if () ion or (-) ion
• How do you make a positive ion?
• How do you make a negative ion?

Remove electrons
Add electrons
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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
• Non-metals 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

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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
• examples C and O

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O2 and O3 - both gas phase

• O2 (oxygen) - necessary for life
• O3 (ozone) - toxic to life

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• Graphite, diamond
• both carbon in solid form

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Graphite and Diamond both carbon in solid form