Atomic Electron Configuration and Periodic Periodicity

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Atomic Electron Configuration and Periodic
Periodicity
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Orbitals

• s
• p
• d
• f

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Quantum Numbers

• Act as an address to indicate area of probable
  location of finding an electron
• Quantum Number symbol designates
• 1st (principal) n energy
  levels (1-7)
• 2nd l
  sublevels (s, p, d, f)
• 3rd m
  orientation in space
• 4th s
  spin

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Electron Spin Magnetic Quantum Number ms

• Electrons have a spins
• If two electrons occupy the same suborbital
• they have opposite spins

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Paramagnatism

• diamagnetic substances that are slightly
  repelled by a strong magnet
• paramagnetic substances that are attracted to a
  magnetic field
• Arises from electron spin
• 1electron system 2 orientations possible
• Aligned with field or opposed
• Paramagnetism the attraction to a magnetic
  field in which the constituent ions or atoms
  contain unpaired electrons

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Ferromagnetism

• Occurs when spins of unpaired electrons in a
  cluster (domain) of atoms in a solid align
  themselves in the same direction.
• Once aligned in a magnetic field, metal is
  permanently magnetized
• Only Fe, Co, Ni and a few others exhibit this
  property

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• Principal sublevels sublevel
  total e
• energy level type
  2n2
• (n) (orbital)
• 1 s
  2
• 2 s, p
  8
• 3 s, p, d,
  18
• 4 s, p, d, f
  32

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• Orbital Number of Number of e- in
• type orbitals orbital
• s 1 2
• p 3 (along x, y, z axis) 6
• d 5 10
• f 7 14

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Quantum Rules

• Pauli Exclusion Principle
• No more than 2 es can occupy one orbital
• Aufbau Principle
• es occupy lowest level available
• Hunds Rule (school bus rule)
• each one of a given orbital is filled with one
  electron before a second enters.
• All 3p orbitals contain one electron before a
  second enters

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Electron Orbitals
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Energy Diagrams

• Orbital Filling Sequence
• 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s,
  4f, 5d, 6p, 7s, 5f, 6d, 7p

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Electron Configuration and Orbital Box Diagrams
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Condensed Electron Configuration ( Noble Gas
Notation)
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Some exceptions

• Cr Ar 4s13d5
• Cu Ar 4s13d10
• Ag Kr 5s14d10
• Why Stability of electron arrangement
• Filled s and p orbitals most stable
• Filled sublevels
• Half filled sublevels

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Electron Configuration of Ions

• Cations- formed when valance electrons removed
  (electrons from highest energy level-n)
• If there are several subshells within nth level,
  electron(s) of maximum l value removed
• Fe Ar 4s23d6 ? Fe2 Ar 3d6 2e-
• Fe2 Ar 3d6? Fe3 Ar 3d5 2e-
• Chemical and physical properties of transition
  metal cations determined by presence of electrons
  in d orbitals
• example paramagnetism
• look at example 8.4 pg 304

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Valence Electrons

• Furthest electrons from nucleus
• Involved in chemical reactions
• Electrons in the highest energy level (valence
  shell)
• Highest value of n.
• Na 1s2 2s2 2p6 3s1
• Ge 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p2
• I Kr 5s24d105p5

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Effective Nuclear Charge Zeff

• Nuclear charge experienced by a particular
  electron in a multielectron atom
• Outer electrons are attracted toward the nucleus
  by the nuclear charge but are pushed away by the
  repulsion of inner electrons.
• As a result, the nuclear charge actually
• felt by outer electrons is diminished,
• and the outer electrons are shielded
• from the full charge of the nucleus
• by the inner electrons

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Periodic Table

•  Groups/Families
• vertical columns
• Similar chemical and physical properties.
• Periods
• horizontal rows.
• Increasing atomic number
• Numbering system
• NEW - The families are labeled 1 18.
• OLD - The families are labeled 1A 8A for s and
  p blocks
• labeled 1B 8B (not logical order)

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• Metals elements to the left of the line
• Non-metals elements to the right of the line
• Semi-metals elements touching the line

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Atomic Properties and Periodic Trends

• Relate electron configuration to physical and
  chemical properties of elements
• These properties change in a predictable manner
  as go move left to right across a period or down
  a group

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Atomic Radius

• Radius the nucleus center to the outermost
  electron.
• Increases going down a group
• Outermost electrons found in orbitals with higher
  values of n
• Electrons in outer shell further from nucleus
• Decreases going across a period
• n remains the same
• Zeff increases as go across period (effect of
  additional proton more important than addition of
  electron

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Ionic Radius

• the nucleus center to the outermost electron.
• Compare ion with the atom
• Anions are bigger than the atom
• Cations are smaller than the atom

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Ionization Energy (IE)

• The energy needed to remove one electron from an
  atom in the gas phase (endothermic process)
• Mg(g) ? Mg(g) e- ?E IE
  738kJ/mol
• Atoms can have a series of ionization energies
• Provides experimental evidence of electron shell
  structure of atoms
• First ionization
• Mg(g) ? Mg(g) e- ?E IE 738kJ/mol
• Second ionization
• Mg(g) ? Mg2(g) e- ?E IE 1451kJ/mol
• Third ionization
• Mg2(g) ? Mg3(g) e- ?E IE 7733kJ/mol

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Ionization Energy

• First ionization energies increase as go across
  period
• Increase in Zeff
• First ionization energies decrease as go down a
  group
• Electron being removed is further from nucleus-
  less nucleus electron attractive force

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Electron Affinity (EA)

• Energy of process in which electron is acquired
  by an atom in the gas phase
• F(g) 1e- ? F-(g) ?E EA -328kJ/mol
• Increase in EA as go across period
• Gets easier to add electron
• IE values get more negative
• Not a smooth trend
• Decrease in EA as go down group
• Attractive force between nucleus and electron
  decreases
• Element with high ionization energy usually have
  a high electron affinity
• The addition of a second electron is an
  endothermic process

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Electronegativity

• Electronegativity
• ability to attract electrons in a chemical bond
• Increases as go across period (excluding noble
  gases)
• Decreases as go down a group
• Unitless
• Fluorine is most electronegative element
• Used to predict polarity of bond