Periodic Table

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Periodic Table

• Periodic table is arranged according to
  increasing atomic number by Henry Moseley.
• Originally arranged by order of increasing atomic
  mass by Demetri Mendelev.
• The word periodic means pattern. Chemical and
  physical of atoms properties are a function of
  atomic number The Periodic Law.
• Groups
• Periods

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Metals

• Electron donors (in ionic bonds)
• Will only be found in ionic bonds not covalent
• Conduct heat and electricity
• Metals are found to the far left of the periodic
  table(before staircase) the further left the
  more metallic in nature.

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Non Metals

• Electron acceptors (negative charge in ionic
  bonds)
• Insulators
• Non metals are found to the far right of the
  periodic table.
• Non metal to non-metal bonding is covalent.

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Names of Groups

• Group 1- Alkali Earth Metals
• Group2- Alkaline Earth Metals
• Group3-12 Transition Metals
• Post Transition Metals
• Group 13- Boron Group
• Group 14- Carbon Group
• Group15-Nitrogen Group
• Group 16- Oxygen group
• Group 17- Halogen group
• Group 18- Noble gases

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Groups contd

• Metalloid- elements with both metallic and
  nonmetallic properties
• Inner transition metals- are located in the f
  subshell. Lanthanide series(4f) and actinide
  series( 5f)
• Hydrogen - contains the element hydrogen and has
  a 1 and -1 charge
• Noble gases are inactive gases
• Halogen group- is the most active non metal group
• Alkali Metals- are the most active metal group.

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Octet Rule

• Also known as the rule of 8
• Atoms want the outer most shell to be completely
  filled.
• For the s and p blocks this means a valence of 8
  (or 2) electrons.
• They will react in a way either ionically or
  covalently to acquire that valence of 8.
• In nature because of the octet rule you never
  find elements by themselves only in compounds
  except for noble gases. These are called diatomic
  elements H2, N2 , O2, F2, Cl2, Br2, I2

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Oxidation StatesThe charge and atom will have in
order to fulfill the octet rule in an ionic bond.

• Group1 1
• Group2 2
• Group13 3
• Group 14 4 or 4
• Group 15 -3
• Group 16 -2
• Group 17 -1
• Group 18 0
• Transition (and Post-Transition) Metals will
  have variable oxidation states. Not as
  predictable as the main group elements (know the
  common ones Fe, Cu, Cr, Pb, Sn).

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Key Concepts

• Valence electrons- total electrons in the outer
  most shell
• Core electrons- everything else except valence
  electrons.
• Groups- goes from top to bottom on the periodic
  table.
• Periods- goes from left to right on the periodic
  table.

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Effective Nuclear Charge

• As you add more protons the pull of the nucleus
  gets stronger as felt by the valence electrons.
• Effective nuclear charge increases as you go
  across the periodic table from left to right.
• It decreases as you go down the periodic table
  from top to bottom because of the increase in
  valence electron distance and electron shielding.
• Example Lets compare Li and F

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Electron Shielding

• The more core electrons there are the less of a
  pull the valence electrons will feel from the
  nucleus.
• Shielding stays constant as you go across the
  periodic table (left to right) and increases as
  you down the periodic table.
• Example Lets compare Li and Cs

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Atomic Radius

• Size of the atom.
• As you go from left to right the atomic radius
  decreases because of the stronger effective
  nuclear charge.
• As you go from top to bottom the atomic radius
  increases because of the shielding valence
  distance.
• As you go down the group you are adding more
  energy levels and this increases the size of the
  atom.

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Ions
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Ionization Energy

• Energy needed to remove the outermost electron.
• As you go from left to right on the periodic
  table the ionization energy increases because of
  stronger effective nuclear charge.
• As you go from top to bottom the ionization
  energy decreases because of more electron
  shielding an increase in the distance of the
  valence eletctrons.
• Example Put in order of increasing ionization
  energy, Be, Li, F

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Electronegativity

• Measurement of the ability of an atom to attarct
  an electron.
• As you go from top to bottom the
  electronegativity decreases because of shielding
  and distance of valence.
• As you go from left to right the
  electronegativity increases because of effective
  nuclear charge.
• Example