Chapter 4: The Periodic Table

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Chapter 4 The Periodic Table

• Section 3 Trends in the Periodic Table

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Key Terms

• Ionization energy removes an electron from an
  atom or ion
• Electron shielding when inner electrons cancel
  some of the charge of the nucleus lessen its
  attraction of outermost electrons
• Trend predictable change in a particular
  direction

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Key Terms 2

• Atomic radius depends on volume occupied by
  atoms electron cloud
• Bond radius half the distance between nuclei of
  atoms that are bonded together
• Electronegativity measure of how much an atom
  in a chemical com-pound can attract electrons

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Key Terms 3

• Electron affinity change in energy when a
  neutral atom gains an electron
• There are no existing bonds as with
  electronegativity.

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Things To Know/Answer

• What are the periodic trends in ionization
  energy, how are they affected by atomic
  structures?
• Answer these questions for atomic radius,
  electronegativity, ionic size, electron
  affinity, melting points (mp) boiling points
  (bp).

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Periodic Trends

• All in a group can be explained by electron
  configurations.
• Reactivity in alkali metals rise from top to
  bottom in group 1.

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Periodic Trends 2

• When one adds enough energy to overcome the
  attraction between protons electrons in an
  atom, it becomes charged (ion) and an electron
  escapes. See page 133 Figure 16.
• A ionization energy ? A e-

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Ionization Energy

• Ionization energy decreases down-ward in a group
  because the number of energy levels increases
  downward.
• Increasing energy levels have increasing distance
  from the nucleus where the positive charge that
  attracts e- in the atom is.

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Ionization Energy 2

• The farther an e- is from the nu-cleus, the less
  attraction protons there have on the e-.
• Also, the higher an e-s energy level is the more
  full levels of e-s there are between it the
  nucleus.

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Ionization Energy 3

• These e-s in the middle reduce the positive
  attraction that extends from the nucleus through
  the atom (electron shielding).
• Outermost electrons are less tightly held for
  this reason.

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Ionization Energy 4

• Ionization energy increases as you move across a
  period. See Figure 17 on page 134.
• This is because protons increase 1 at a time
  rightward on a row, but the energy level stays
  the same.

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Ionization Energy 5 / Atomic Radius

• Increasing e-s get crowded repel each other
  this counteracts the positive attrac-tion of the
  nucleus.
• Here, electrons can only get so close before
  increased positive attraction can-not overcome e-
  to e- repulsion.
• For this reason, atomic radius stops decreasing
  rightward in a row.

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Atomic Radius 2

• Increasing atomic across a row pro-duces a much
  bigger rise in positive attraction than rise in
  distance of e-s from the nucleus.
• This pattern also causes decreased atomic radius
  left to right on a row.

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Atomic Radius 3

• Rising attraction also pulls electrons closer
  to the nucleus.
• Added inner energy levels are present downward in
  groups add distance from the nucleus.

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Atomic Radius 4

• Electron shielding also blocks rises in
  attraction yields similar attraction down the
  group.
• These effects cause increasing atomic radius
  downward in a group.

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Electronegativity

• Electronegativity is relative attraction of
  electrons by nuclei in bonded atoms where they
  play tug of war with their shared electrons.
• Linus Pauling, one of Americas most famous
  chemists, made a scale of electronegativity
  values.

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Electronegativity 2

• In the scale, he assigned F 4.0 since it attracts
  electrons in bonds most then Pauling calculated
  values for other elements relative to this one.
• Electronegativity decreases down a group mostly
  because higher energy levels are farther from the
  nucleus.

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Electronegativity 3

• Nuclei cannot attract valence electrons in these
  distant energy levels well.
• For this reason, an element like Cs has a nucleus
  w/ more protons but weak attraction of a valence
  electron on its 6th energy level.

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Electronegativity 4

• However, an element like Li attracts a valence
  electron on its 3rd energy level more strongly.
• This makes Li more electronegative than Cs.
• Electronegativity increases sharply rightward
  across a period.

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Electronegativity 5

• This trend arises because no change in electron
  shielding happens across a row since no electrons
  get added to inner energy levels.
• As atomic rises quickly, nuclear charge does
  also and can attract bond electrons much more
  strongly.

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Electronegativity 6

• Adding inner electrons downward in a group
  increases electron shielding.
• This keeps effective nuclear charge mostly the
  same.

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Electronegativity 7

• Slight drops in electronegativity result.
• This is b/c distance from the nucleus is the key
  factor not nuclear charge.
• Slight rises in distance from the nucleus
  downward in a group have much less effect than
  boosts in nu-clear charge rightward across a row.

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Other Periodic Trends

• Effective nuclear charge electron shielding
  explain most periodic trends including ionic size
  electron affinity.
• Ionic size follows trends of atomic radii for the
  same reasons.

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Other Periodic Trends 2

• Metals tend to lose one or more electrons
  become cations (ions) whereas, nonmetals tend
  to gain e-s form anions (-ions).
• Electron affinity follows electronega-tivity
  trends (decrease down a group but increase right
  across a series) for the same reasons.

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Other Periodic Trends 3

• Mp bp do not generally rise or fall but reach 2
  different peaks as d p orbitals fill.
• For example, Cs has low mp bp b/c it only has 1
  valence electron for bonding it is far left in
  6th period.

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Other Periodic Trends 4

• As the electron increases across a row, more
  bonds can form require more energy to break.
• This effect peaks near the middle of d-block
  elements at W and Re b/c the d orbitals are half
  filled.

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Other Periodic Trends 5

• Further e-s pair in d orbitals beyond W
  decrease the of unpaired e-s that help
  strengthen bonds between atoms by forming
  multiple bonds.
• More rightward, Hg Rn have much lower mp bp
  b/c d orbitals are full.

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Other Periodic Trends 6

• Past Hg, mp bp rise again as elec-trons start
  filling p orbitals until these are half filled.
• Beyond half filled status, mp bp drop again b/c
  the p orbitals get full and unable to help
  strengthen bonds.
• By Rn, p orbitals are full also so mp bp are
  unusually low.