Shapes of molecules & IMF

• Octet rule — > The tendency of atoms to prefer to have 8 electrons in the outer shell
• Dative bonding —> covalent bonds in which the shared pair of electrons has been supplied by one of the bonding atoms only.
• Lone pair —> a pair of electrons in the outer shell that are not part of a covalent bond.
• A dative (coordinate) bond occurs when an atom shares a lone pair of electrons to form a bond —> rather than both elements providing one each
• Electron pair repulsion theory —> electron pairs repel each other to be as far apart as possible. Used to explain and predict the shapes of molecules & polyatomic ions
• The amount of repulsion from greatest to least:

  • 2 lone pairs
  • lone pair and bonded pair
  • 2 bonded pairs
• The number of bonded and lone pairs will determine a molecules shape
• • The electron pairs surrounding a central atom determine the shape of the molecule or ion
  • The electron pairs repel one another so that they are arranged as far apart as possible
  • The arrangement of electron pairs minimise repulsion and holds the bonded atoms in a definite shape
  • Different number of electron pairs result in different shapes
• Wedges:

  • Chemists use wedges to help visualise structures in 3D
  • a solid line = a bond in the plane of the paper
  • a solid wedge = comes out of the plane of the paper
  • a dotted line = goes into the plane of the paper to the back
  • a ballon with 2 dots = lone pair
• 2 electron pairs = 2 bonded pairs, 0 lone pairs, linear shape, angle = 180
• 3 electron pairs = 3 bonded pairs, 0 lone pairs, trigonal planar shape, angle = 120
• 4 electron pairs = 4 bonded pairs, 0 lone pairs, tetrahedral, angle = 109.5
• 5 electron pairs = 5 bonded pairs, 0 lone pairs, trigonal bipyramidal, angle = 120 + 90
• 6 electron pairs = 6 bonded pairs, 0 lone pairs, octahedral shape, angle =90
• 4 electron pairs = 3 bonded pairs, 1 lone pair, trigonal pyramidal shape, angle = 107
• 4 electron pairs = 2 bonded pairs, 2 lone pairs, angular (bent) shape, angle = 104.5
• Molecular shapes from 4 electron pairs:

  • 4 electron pairs around the central atom repel one another as far apart as possible into a tetrahedral arrangement
  • Lone pairs repel bonded pairs slightly closer together reducing the bond angle - the angle between the bonded pair of electrons
  • The bond angle is reduced by 2.5 degrees for each lone pair
• Molecular shapes from multiple bonds:

  • 4 bonded pairs around the central carbon atom in CO2 are arranged as 2 double bonds, which count as 2 bonded regions
  • 2 bonded regions repel one another as far apart as possible therefore in a linear shape
• Shapes of ions:

  • An ammonium ion has the same number of bonded pairs of electrons around the central atom as a methane molecule
  • Ammonium has the same tetrahedral shape and bond angles (109.5) as a methane molecule
  • Carbonate and ammonium ions have 3 regions of electron density surrounding the centre atom
• Electronegativity —> the attractive pull of an element on a pair of electrons in a covalent bond. The more electronegativity = greater the pull
• • Electronegativity increases across a period as the increase of protons in the nucleus increases the nuclear charge
  • Electronegativity increases up a group as there is less electron shielding of the nucleus. Less electrons around the nucleus to repel the outer shell electrons
  • Common elements with increases electronegativity = nitrogen, oxygen, chlorine and flourine
• Bond polarity —> atoms with the same or very similar electronegativity, the electron pair will be shared equally
• • Covalent bonds with a large difference in electronegativity between atoms causes the electron pair to shift towards the more electronegative element, crating a dipole
  • A small negative charge (delta -) on the more electronegative element
  • A small positive charge (delta +) on the less electronegative element
• Polar molecules:

  • Symmetrical molecules will not be polar as the dipoles cancel each other out
  • Polar molecules occur if the bond dipoles can’t cancel out, therefore giving the overall molecule a dipole
• Permanent dipole - dipole:

  • Polar molecules contain permanent dipoles due to the electronegativity of elements within them and their shape
  • Atoms/ groups that have a delta + charge will be attracted to those in another molecule with a delta - charge
• Induced dipole - dipole:

  • All molecules, polar & non-polar, will experience instantaneous dipoles caused by the movement of electrons . Electrons move closer to one side of the molecules creating a delta - charge
  • The instantaneous dipole that forms can the induce a dipole in another nearby molecule. An instantaneously induced dipole
  • This is the weakest type of London force
• Effect of properties:

  • The type of intermolecular force present will affect the properties of a chemical

  1. Melting/boiling point —> the stronger or more numerous the forces present, the more energy is required to break them, raising the melting/ boiling point
  2. Solubility —> “Like dissolves like” ; polar substances can easily dissolve into polar solvents as they can easily produce permanent dipole-dipole interactions
• Simple molecular lattices:

  • Simple molecules can form lattices with molecules arranged in a repeating pattern
  • The lattice is only held together by the weak induced dipole-dipole forces present
• Sublimation:

  • Iodine and carbon dioxide for simple molecular lattices held together with induced dipole-dipole forces
  • When heated the weak intermolecular forces are easily broken allowing molecules to escape the lattice
• Hydrogen bonding:

  • Strongest of intermolecular forces
  • For hydrogen bonding to occur between molecules they need a lone pair of electrons and a hydrogen covalently bonded to either a fluorine, oxygen or nitrogen atom.
  • Hydrogen covalently bonded t elements with high electronegativity
• When drawing hydrogen bonding you must:

  • Show dipoles in the atoms involved
  • Draw dashed lines to show the bond
  • the bond must go from the delta + hydrogen to he lone pair on the delta - atom
• • Water is less dense as a solid
  • Liquid water - hydrogen bonds constantly broke and reformed
  • Solid water - molecules arranged in hexagonal rings to maximise the number of stable hydrogen bonds
  • The more open interlocking ring structure creates a larger volume, lowering the density