periodicity

• Periodicity


  The study of the repeating patterns or trends that occur in physical or chemical properties as we move around in the periodic table
• Regions of the periodic table

  • Groups (up and down columns)
  • Periods (rows)
  • Blocks (s, p, d)
• Blocks


  • The block an element is in corresponds to the subshell that you find the outer electrons in
• Blocks


  • s block (group 1 and 2)
  • p block (final 6 groups)
  • d block (transition elements)
• Atomic radius


  The distance between the center of the atom (nucleus) and the electrons in the outermost energy level
• Atomic radius


  Larger if the atom has more occupied shells of electrons, smaller if fewer occupied electron shells
• Across period 3


  Atomic radius decreases
• Reason for decreasing atomic radius across period 3
• Electronegativity


  The ability of an atom to attract electron density or electrons in a covalent bond
• Across period 3


  Electronegativity increases
• Reason for increasing electronegativity across period 3
• Within a group


  Electronegativity increases going up the group
• Most electronegative elements are in the top right of the periodic table, least electronegative in the bottom left
• Ionization energy


  The enthalpy or energy required to remove a mole of electrons from a mole of gaseous atoms
• Across period 3


  Ionization energy increases
• Reason for increasing ionization energy across period 3
• Pattern of ionization energy across a period

  General increase with two dips, one between s block and p block, one between p block elements
• Reason for dip between s block and p block
• Reason for dip between p block elements
• p sub shell


  • Three orbitals
  • Electrons occupy the orbitals singly
• Phosphorus


  • Electrons in p sub shell occupy the orbitals singly
• Sulfur


  • One of the orbitals has two electrons
  • The two electrons repel each other
  • One electron has slightly more energy than expected
• Electron repulsion in Sulfur

  Makes one electron slightly easier to remove
• You might be asked to explain the shape of the graph, the dips, or finish the graph
• The pattern of increase, dip, increase, increase, dip, increase, increase continues across a period
• Going down a group

  • Electrons are further from the nucleus
  • Attraction between nucleus and outer electrons is smaller
  • Outer electrons are easier to lose
• Solid


  Atoms in fixed positions, vibrating about fixed points, low energy
• Liquid


  Atoms in a more random arrangement, moving around more, medium energy
• Gas


  Atoms spread out, moving rapidly, high energy, low forces between particles
• Melting point


  Temperature at which a solid turns into a liquid
• Boiling point


  Temperature at which a liquid turns into a gas
• Stronger forces between particles

  • More energy required to separate them
  • Higher melting and boiling points
• Melting point pattern across period 3

  • Three distinct regions: Metals, Giant covalent, Non-metals
• Metallic bonding


  • Outer electrons delocalized and free to move
  • Atoms are positively charged ions
  • Held together by electrostatic attraction
• Across period 3 metals

  • Melting point increases
  • Nucleus gets more positive
  • Ions get smaller
  • Charge density increases
• Silicon


  • Giant covalent structure
  • Tetrahedral arrangement
  • Requires a lot of energy to break covalent bonds
• Non-metals


  • Simple molecular or simple covalent
  • Experience van der Waals forces between molecules
• Larger molecules/more electrons


  • Stronger van der Waals forces
  • Higher melting point
• Argon


  • Monoatomic noble gas
  • Only van der Waals forces between atoms
  • Very weak forces, lowest melting point
• You could be asked to explain the shape of the melting point graph, fill in missing points, or explain/predict melting/boiling points of elements