Bonding

• going along period -> melting point increases

  • more delocalised electrons & greater nuclear charge

  going down period -> melting point decreases

  • same no. of delocalised electrons & more shielding -> less attraction to other nuclei
• metallic bonding = electrostatic attraction between positive metal ions and negative delocalised electrons

  • can conduct electricity
  • malleable -> layers can slide over each other
  • high melting point -> strong metallic bonds
• ionic bonding = electrostatic attraction between positive & negative ions

  • high melting points -> strong ionic bonds
  • can't conduct electricity when solid but can when molten/in aqueous solution
  • ions are free to move
  • tend to dissolve in water
  • polar water molecules pull ions away from lattice
• covalent bonding = when non-metals share a pair of bonds from their outer shells

  • giant covalent structures (e.g. graphite, diamond, silicon dioxide)
  • graphite can conduct electricity
  • simple molecules
  • less energy needed to break these bonds -> weak intermolecular forces
  • can't conduct electricity
• coordinate bonding = a covalent bond in which both electrons in the shared pair are provided by one atom

  • the lone pair of electrons from 1 atom are donated to an electron deficient atom
  • lone pair of electrons = pair of electrons not involved in bonding
  • gives an overall charge of +1 to compound
• electronegativity = power of an atom to attract a pair of electrons in a covalent bond

  • going along the period -> electronegativity increases
  • going down group -> electronegativity decreases
• factors affecting electronegativity

  1. atomic radius

  • smaller radius -> outer shell electrons closer to nucleus -> greater electron withdrawing (pull) effect -> greater electronegativity

  2. nuclear charge

  • higher nuclear charge -> greater tendency to withdraw (pull) electrons -> greater electronegativity

  3. electron shielding

  • makes electrons easier to remove & lower electronegativity
• Polar bond

  1. More electronegative element -> electrons closer to atom
  2. H has slight positive charge & F has slight negative charge

  • HF has a dipole across the bond (HF molecule is polar)
  • dipole = difference in charge
  • molecule must be asymmetrical
  • otherwise dipoles cancel out

  
• intermolecular forces = attractive forces between covalent molecules


  1. dipole-dipole forces

  • molecules with permanent dipoles have partial charges
  • partial charges of 1 molecule may attract (or repel) the partial charges of another molecule

  2. van der Waals forces

  • weakest intermolecular forces
  • electron movement in a molecule forms an instantaneous dipole
  • this dipole induces a dipole in a nearby atom/molecule
  • all chemicals have van der Waals forces
  • increase van der Waals strength by increasing no. of electrons
• intermolecular forces: hydrogen bonding

  • only 3 cases -> with 3 most electronegative elements (O/N/F)
  • special kind of dipole-dipole interaction & has 1/10 of strength of covalent bond as they are the most polar
  • need very electronegative atom with lone pair covalently bonded to a hydrogen atom
  • H is very electron deficient as O is very electronegative
  • protons have very strong electric field due to small size
  • lone pair from 1 oxygen atom strongly attracted to H in another water molecule
• Hydrogen bonding in water

  • water in liquid state -> H bonds break & reform easily
  • when water freezes -> H bonds hold water molecules in fixed (crystalline) positions
  • to fit this structure -> molecules slightly less closely packed together in liquid water
• VESPR - Valence Electron Shell Repulsion Theory

  • (lone) electron pairs exists in charge clouds thus high chance of finding other electrons
  • electrons repel until they're as far apart as possible to minimise repulsion between them
  • lone pair - lone pair repulsion > lone pair - bond pair repulsion > bond pair - bond pair repulsion
• Shape of molecules pt 1

  2 bond pairs:

  • name: linear bond angle: 180°

  3 bond pairs:

  • name: trigonal planar bond angle: 120°

  4 bond pairs:

  • name: tetrahedral bond angle: 109.5°

  5 bond pairs:

  • name: trigonal bipyramidal bond angle: 120° & 180°

  6 bond pairs:

  • name: octohedral bond angle: 90°
• Shapes of molecules pt 2

  3 bond pairs & 1 lone pair:

  • name: pyramidal bond angle: 107°

  2 bond pairs & 2 lone pairs:

  • name: non-linear/bent bond angle: 104.5°

  4 bond pairs & 1 lone pair:

  • name: seesaw shape bond angle: >120° & >90°/86.5° & 102°

  3 bond pairs & 2 lone pairs:

  • name: T-shape bond angle: 86°

  5 bond pairs & 1 lone pair:

  • name: square pyramidal bond angle: 85° & 90°

  4 bond pairs & 2 lone pairs:

  • name: square planar bond angle: 90°
• Shape of molecules diagram

  A) linear

  B) trigonal planar

  C) non-linear

  D) pyramidal

  E) tetrahedral

  F) trigonal bipyramidal

  G) octahedral

  
• Shape of molecules diagrams pt 2

  A) seesaw

  B) t-shape

  C) square pyramidal

  D) square planar

  
• Finding shape of unknown molecules

  1. Decide on the central atom (usually only 1 of them)
  2. How many electrons does the central atom have in it's outermost shell?
  3. how many atoms is it bonded to?
  4. does it have a negative (add) or positive (subtract) charge?
  5. add them all up & divide by 2
  6. how many bond pairs are there?

  A) O

  B) 6

  C) 3

  D) positive

  E) subtract