2.2 Electrons, Bonding and Structure

• Max. number of electrons given by?


  2n^2 (where n is the shell number)
• An orbital?


  a region which can hold a maximum of two electrons with opposite spins
• Shape of s orbital?


  spherical
• Shape of p orbital?


  dumb bell
• Sub-shells orbitals

  s - 1

  p -3

  d - 5

  f - 7
• Ionic Bonding

  the strong electrostatic attraction between oppositely charged ions.
• Giant Ionic lattice

  metal and non metal bonds
• Ionic compounds have

  high melting point - ionic bonds are strong so a lot of energy is required to break them

  conduct electricity - conducts in molten or in solution because ions are free to move to carry the current, does not conduct when solid because the ions only vibrate about a fixed position in the lattice so cannot carry current
• Covalent bonding

  electrostatic attraction between a shared pair of electrons and the nuclei of the bonded atoms.
• Simple molecular lattice properties:

  weak intermolecular forces between molecules

  low melting and boiling point - weak forces so does not require much energy to break them

  doesn't conduct electricity - no free electrons or charged particles
• Giant covalent lattice (allotropes) properties:

  strong covalent bonds between molecules

  high melting point - strong bonds between molecules so more energy is required to break them

  do not conduct - except graphite
• Dative coordinate bond:

  covalent bond in which both electrons of the shared pair come from the same atom
• Metallic Bonding:

  attraction between positive metal ions and a sea of delocalised electrons

  
• Giant metallic lattice properties:

  high melting point - strong metallic bond

  conducts electricity - delocalised electrons can move through the structure so they can carry charge
• Electron pair repulsion theory:

  all electrons repel each other so that they are as far away as possible

  lone pairs repel slightly more then bonded pairs (around 2' due to each lone pair)
• Tetrahedal:

  4bp and 0lp

  109' bond angle

  
• Pyramidal:

  3bp and 1lp

  107' bond angle

  
• Non-linear:

  2bp and 2lp

  104.5' bond angle

  
• Trigonal Planar:

  3bp and 0lp

  120' bond angle

  
• Linear:

  2bp and 0lp

  180' bond angle

  
• Octahedral:

  6bp and 0lp

  90' bond angle

  
• Electronegativity?


  the ability of an atom to attract the bonding electrons in a covalent bond.
• Bond polarity:

  when one atom in a covalent bond is more electronegative than the other, the electrons are shared unequally, this is a polar bond
• Trend in electronegativity:

  increases across a period

  decreases down a group

  fluorine is the most electronegative 4.0
• London forces:

  -only occur in non-polar substances, such as noble gases

  -caused by the movement of electrons which unbalances the distribution of charge in a molecule causing a temporary dipole

  -bigger molecules have higher london forces

  -more branched molecules have weaker london forces due to more points of interaction
• Permanent Dipole-Dipole:

  -occurs in polar molecules, when the more positive end is attracted to the slightly more negative end of a neighbouring molecule

  -have stronger bonding then london forces so a higher boiling point then molecules with only london forces
• Hydrogen Bonding:

  -occurs when H bonds to either F, O or N

  -formed from a slightly more positive H atom in one molecule and the lone pair of electrons in either F, O or N in a neighbouring molecule

  -strongest type of intermolecular force
• Anomalous properties of water:

  -water has a relatively high boiling point in relation to its size due to strong hydrogen bonds

  -ice is less dense then water due to its open lattice, the atoms are pushed further apart due to hydrogen bonding, so large spaces between molecules