PERIODIC TRENDS

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PERIODIC TRENDS

• OR
• Why we call it the PERIODIC table

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PERIODIC

• What does Periodic mean?
• What are examples of things that are periodic?

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Periodic Trends

• Across the periodic table nucleus becomes more
  positive, more attraction occurs, elements become
  more non-metallic
• Down a group the elements become more positive,
  but there is a larger nucleus due to the
  increased number of protons and neutrons.
  Therefore, there is less attraction between
  protons and outer electron levels and the
  elements have more metallic properties.

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Periodic Trends

• Both the position and the properties of the
  elements arise from their electron configuration
  (which comes from their atomic number).
• Same column similar outer level e-
  configuration
• Properties that are periodic
• metals, metalloids, nonmetals, boiling point,
  density, atomic radii, ionic radii, oxidation
  numbers, ionization energies, electron
  affinities, electronegativity

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Boiling Point of the Elements
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Alkali Metals Video

• Looking at a periodic trend of the alkali metal
  family.

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Alkali Metals Video

• What is the chemical equation for the reaction
  for each of these metals in water?
• 2Li H2O ?2LiOH H2
• 2Na H2O ? 2NaOH H2
• 2K H2O ? 2KOH H2

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A Trend for Groups

• SHIELDING
• The higher the principle quantum number, the
  more energy levels there are. The orbitals are
  further out. Inner electrons shield the positive
  charge of the protons and the electrons in the
  outer energy level come off more easily.

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A Trend for Periods

• PROTON PULL (nuclear charge)
• The greater the positive charge (due to
  increased atomic number), the greater the
  attraction (pull) within the energy level. More
  protons can be attracted to more electrons.

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ATOMIC RADII

• page 188
• As the energy levels increase, the principle
  quantum number increases , and so does the atomic
  radii.
• As the atomic number increases across a period,
  the positive charge also increases within the
  same energy level.
• TREND atomic size increases down a group,
  atomic size decreases across a period.
• Why within a group, inner level electrons shield
  outer level e- from the positive nucleus
  distance from nucleus increases within a period,
  size of nuclear charge increases and the
  attraction between protons and electrons pulls
  the atom in tighter.

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ATOMIC RADII
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IONIC RADII

• page 190
• When atoms form a compound, the compound is more
  stable than the uncombined atoms were. The outer
  energy levels are full (like the noble gases).
• TREND Metallic ions are smaller than the atoms
  they come from. Nonmetallic ions are larger than
  the atoms they come from.
• Why metal lose electrons and become the noble
  gas configuration on the energy level below.
  Nonmetals gain electrons and become the noble gas
  configuration at the end of their period.

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IONIC RADII
14
Sodium Chlorine Salt
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OXIDATION NUMBERS

• This is the charge number an atom would have if
  the valence electrons were completely
  transferred/lost.
• Group 1 (1A) 1 oxidation (H can have 1 or 1-)
• Group 2 (IIA) 2 oxidation
• Group 3-12 (IIIB) can have 1 to 8 (lose d
  and then s, one at a time)
• Group 13 (IIIA) 3
• Group 14 (IVA) 2 or 4
• Group 15-17 gain electrons (-3, -2, -1)
• Group 18 noble gases 0

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OXIDATION NUMBERS
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IONIZATION ENERGY

• page 191-192
• This is the energy required to remove an
  electron. The first ionization energy is the
  energy required to remove the most loosely held
  electron.
• TREND Ionization energy increases as atomic
  number increases across a period. In a group,
  there is a gradual decrease in the first
  ionization energy as the atomic number increases.
  Metals have low first ionization energies and
  nonmetals have high first ionization energies.
• Why? Across a period, the larger the nuclear
  charge, the greater the ionization energy Down a
  group, the greater shielding effect leads to less
  ionization energy With a bigger electron cloud
  (more energy levels), the ionization energy
  decreases an electron in a full or half-full
  sublevel requires additional energy to be removed.

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FIRST IONIZATION ENERGY
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IONIZATION ENERGIES

• (kilojoules per mole)
• 1st 2nd 3rd 4th 5th
  6th
• H 1312.0
• He 2372.3 5220
• Li 520.2 7300 11750
• Be 899.5 1760 14850 20900
• B 800.6 2420 3660 25020 32660
• C 1086.5 2390 4620 6220 37820 46990
• Al 577.5 1810 2750 11580 14820
  18360
• Ga 578.8 1980 2970 6170 8680
  11390

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ELECTRON AFFINITY

• The attraction of an atom for an electron the
  energy change that occurs when an atom gains an
  extra electron. The same factors that affect
  ionization energies affect electron affinities.
• TREND Nonmetals have large electron affinities
  (except for Noble gases). The more stable an atom
  is, the less the tendency to have an affinity for
  electrons. Metals have low electron affinities
  and Nonmetals have high electron affinities.
  Noble gases (p6) and Alkaline Earth Metals (s2)
  have negative affinities (extremely low).

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ELECTRON AFFINITY

• H He
• 72.766 (-21)
• Li Be B C N O F Ne
• 59.8 (-241) 23 122 0 141 328 (-29)
• Na Mg Al Si P S Cl Ar
• 52.9 (-230) 44 120 74 200 349 (-34)
• K Ca Ga Ge As Se Br Kr
• 46.36 (-156) 36 116 77 195 325 (-39)
• Rb Sr In Sn Sb Te I Xe
• 46.88 (-167) 34 121 101 190 295 (-40)
• Cs Ba Tl Pb Bi Po At Rn
• 45.5 (-52) 48 101 101 (170) (270) (-41)

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ELECTRONEGATIVITY

• page 194
• A tug of war between atoms for electrons in a
  chemical bond. How well the electrons tug is
  electronegativity. The better the atom tugs an
  electron from another element, the higher the
  electronegativity
• TREND increases from left to right across a
  period as the number of valence electrons
  increases and the size of the atom decreases. A
  low ionization energy and a low electron affinity
  means low electronegativity.
• Fluorine is the most electronegative element.

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ELECTRONEGATIVITY
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ELECTRONEGATIVITIES
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SUMMARY

• Noble Gases
• - valence electrons s2p6
• - ionization energy HIGH
• - Electron Affinity LOW
• - Electronegativity LOW

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SUMMARY

• Halogens
• - valence electrons s2p5
• - ionization energy HIGH
• - Electron Affinity HIGH
• - Electronegativity HIGH

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SUMMARY

• Alkali metals and Alkaline Earth metals
• - valence electrons s1 and s2
• - ionization energy LOW
• - Electron Affinity LOW
• - Electronegativity LOW

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SUMMARY - STABILITY