chem module 2

• linear

  • 180 degrees
  • 2 bond pairs
  • no lone pairs
• trigonal planar

  • 120 degrees
  • 3 bond pairs
  • no lone pairs
• tetrahedral

  • 109.5 degrees
  • 4 bond pairs
  • no lone pairs
• octahedral

  • 90 degrees
  • 6 bond pairs
  • no lone pairs
• pyramidal

  • 107 degrees
  • 3 bond pairs
  • 1 lone pairs
• angular

  • 2 bond pairs
  • 104.5 degrees
  • 2 lone pairs
• • electron pairs repel eachother
  • lone pairs repel more than bond pairs
• properties of giant metallic lattices and structure

  • structure: delocalised electrons throughout and can move, charges must balance
  • high mp + bp: attraction between + ions and - ions (delocalised electrons) is very strong -> high temp needed to overcome metallic bonds
  • good electrical conductivity: delocalised electrons move freely
  • ductile: can be drawn out and stretched
  • malleable: can be hammered into shapes -> layers slide past eachother
• metallic bonding

  • electrostatic attraction between positive metal ion (cation) and delocalised electrons
  • strength of bond and mp depends on number of delocalised electrons and size of ion
• smaller charge -> more delocalised electrons -> greater attraction
• intermolecular forces (london forces)

  • caused by constant random movement of electrons in shells -> unbalances distribution of charge
  • non polar
  • instantaneous dipole -> induces dipole in neighbouring molecules
  • small induced dipoles attract one another -> causes weak intermolecular forces (london forces)
  • more electrons -> bigger attraction/stronger forces -> larger dipoles (more electrons)
• effect of london forces on boiling points

  • bonding within molecules stay the same (when heated)
  • molecules become separated -> no longer attracted to eachother
  • as number of electrons increases -> so does strength of forces
  • no dipoles present in any molecules before they interact (london forces)
• permanent dipole-induced dipole

  • polar bonds present
  • slightly negative - and + end
  • when near non polar molecules -> able to cause electrons in shells in nearby molecules to shift slightly (by being repelled by - side of attracted to + side) -> causes non-polar molecule to become slightly polar -> attraction occurs
  • permanent dipole has induced dipole in another molecule
• permanent dipole- permanent dipole

  • molecules with permanent dipoles will also be attracted to other molecules with permanent dipoles
  • opposite ends attract to one another
• hydrogen bonding

  • molecules containing: N-H, F-H, O-H bonds
  • these are polar with permanent dipoles and dipoles are strong
  • attraction between lone pair of electrons on a highly electronegative atom (O, N or F)
• permanent dipole

  • difference in electronegativity -> ability of an atom to attract a bonding pair of electrons
  • bigger difference in electronegativity -> the bigger and stronger the forces between two atoms ->
  • electrons will be pulled closer to each atom
  • atom gains small negative charge
  • other atom gains small positive charge
  • this produces a polar covalent bond
• london forces (induced dipole-dipole interactions)

  • electrons move around
  • at any instant, its possible more electrons will lie to one side of atom/ molecule than the other
  • instantaneous dipole produced
  • induces a dipole in nearby atoms/ molecules -> more electrons, stronger forces
• hydrogen bonding

  • H bonded to O, F of N
  • O, F or N have to have a lone pair
  • extension of dipole- dipole interaction given even higher boiling points
  • bonds between H and 3 most electronegative elements F, N or O are very polar
  • small sizes of f,O,N,H concentrates are partial charges in smaller volume leading to high charge density
  • intermolecular attractions are greater, leading to higher boiling points
• • solids = moles=mass/Mr
  • liquids = moles=concentration x volume
  • gas = moles=volume/24
• molar mass


  mass per mole of a substance
• • acid: donate protons (hydrogen ions H+)
  • base: accepts protons (via lone pair)
  • strong: one mole of HCL they would all 'split' to form one mole of H+ ions and one mole of Cl- ions
  • weak: 1% split
• • dilute acid: acid molecules mixed with large amount of water so there's only a low concentration of H+ ions
  • concentrated acid: acids have little to no water molecules mixed with acid molecules -> concentration of H= ions is high
• basic oxidation rules

  • oxidation state of atoms in elements is 0
  • oxidation state of ions is their charge
  • complex ions the sum of all oxidation states is equal to overall charge on ion
  • in compounds, sum of all oxidation states is 0. balance out. one part positive and one negative
• basic oxidation rules 2

  • group 1 elements = +1
  • group 2 elements = +2
  • group 3 elements = +3
  • fluorine = always -1
  • hydrogen = +1
  • oxygen = -2
  • chlorine = -1
• molality


  ratio of moles of solute to the volume of solvent in kilograms
• molarity


  ratio of moles of solute to the volume of solution in litres
• molality equation


  molality = molarity x molar mass
• relative isotopic mass 


  mass of an atom of an isotope compared with 1/12 of mass of an atom of carbon -12
• electronegativity


  measure of the attraction of an electron in a covalent bond
• percentage yield


  (actual amount/theoretical amount) x 100
• atom economy


  (mass of desired product/mass of all products) x 100
• permanent dipole


  small charge differences across a bond that results from a difference in electronegativities of bonded atoms
• hydrogen bond


  strong dipole-dipole attraction between, an electron deficient hydrogen atom on one molecule and a lone pair of electrons on a highly electronegative atom on a different molecule
• relative atomic mass


  weighted mean mass of an atom of an element compared with 1/12 of the mass of an atom of carbon -12
• strong base


  completely dissociate into ions in aqueous solutions resulting in loads of OH- ions
• strong acid


  completely dissociate into ions in aqueous solutions resulting in many H+ ions
• weak acid


  partially dissociates into ions and are excellent at accepting H+ ions
• weak base


  partially dissociate into ions and reacts with aqueous solution to form OH- ions
• water of crystallisation


  water molecules that form an essential part of the crystalline structure of a compound
• number of particles = moles x avagrados constant

  avagrados constant = 6.02 x 10^23