Chapter 6: Shapes of molecules and intermolecular forces

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6.1 Shapes of molecules and ions

Electron repulsion theory:

Electron pairs around the central atom repel each so that they are as far apart as possible

The repulsion between electron pairs and the number of electron pairs determines the shape of the molecule or ion, the repulsion holds bonded atoms in a definate shape

Lone pairs repel more strongly than bonding pairs because they are slightly closer to the central atom and occupy more space

Lone pairs repel bonding pairs slightly closer together, decreasing the bond angle by about 2.5 degrees for each lone pair

  • Linear:                          2BP            180     degrees
  • Trigonal planar:          3BP            120     degrees
  • Tetrahedral:                 4BP            109.5  degrees
  • Octahedral:                  6BP            90      degrees
  • Trigonal pyramidal:    3BP 1LP     107     degrees
  • Non-linear:                  2BP 2LP     104.5  degrees
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6.1 Shapes of molecules and ions

Ions can be drawn with the charge on the atoms or the shape in brackets with the overall charge

  • CO3^2- Trigonal planar (3BP 120) 1 double bond, 2 single bonds with charge
  • NO3^-   Trigonal planar (3BP 120) 1 double bond, 1 single bond with charge, 1 dative
  • SO4^2- Tetrahedral (4BP 109.5) 2 double bonds, 2 single bonds with charge

Determining the shape of a molecule / ion:

  • Draw the dot and cross diagram
  • Electron pairs repel each other so that they are as far apart as possible
  • Lone pairs repel more than bonded pairs 
  • Count the number of bonded pairs and the number of lone pairs
  • Draw the 3D shape and quote the name and its bond angle
  • (include 4 examples; with shape, name and bond angle)
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6.2 Electronegativity and polarity (table)

Electronegativity: a measure of the attraction of a bonded atom for the pair of electrons in a covalent bond

  •                                  Across a period        Down a group
  • Electronegativity          increases                 decreases
  • Nuclear charge            increases                 increases
  • Atomic radius              decreases                increases
  • Shielding                    stays the same         increases

Pauling electronegativity values:

  • Non-metals, F, O, N, Cl, are the most electronegative elements
  • Group 1 metals, Fr, Cs, are the least electronegative elements
  • Noble gases are not given a value because they have full outer shells and tend not to form compounds
  • EN difference = 0 ---> pure covalent
  • EN difference = 0 to 1.8 ---> polar covalent
  • EN difference = >1.8 ---> ionic
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6.2 Electronegativity and polarity

  • Pure covalent bond: an equally shared pair of electrons between 2 atoms
  • Occurs when bonded atoms have the same or similar electronegativities
    • E.g. H2, O2, Cl2 or hydrocarbons, C6H14
  • Polar covalent bond: an unequally shared pair of electrons between 2 atoms
  • Occures when bonded atoms are different and have different electronegativities
    • Example: HCl
  • Cl is more electronegative than H
  • so Cl has a greater attraction for the shared pair of electrons
  • a polar covalent bond if formed between the 2 atoms
  • H-Cl bond is polarised with delta+ charge on H and delta- charge on Cl

Dipole: a separation of electrical charge so that one end of a polar molecule has a d+ charge and the other end has a d- charge

  • Atom with the larger EN value will have the delta negative charge
  • Atom with the smaller EN value will have the delta positive charge
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6.2 Electronegativity and polarity

Polar molecules:

Each bond e.g. the O-H bond in water, has a permanent dipole

The 2 dipoles don't cancel so a polar bond is formed

Non-polar molecules:

Each bond e.g. the C = O bond in CO2, has a permanent dipole

Dipoles exactly oppose one another, the dipoles cancel (because the molecule is symetrical) so the molecule is non-polar

A molecule can have polar bonds without being a polar molecule

E.g. SF4 has polar bonds (different electronegativities) but is non-polar because the molecule is symetrical and the dipoles cancel

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6.3 Intermolecular forces

Intramolecular forces: strong forces of attraction within molecules

E.g. the covalent bonds within H2 molecules (H - H)

IN forces are responsible for chemical properties

Intermolecular forces: weak forces of attraction between (dipoles of) different molecules

E.g. the london forces between H2 molecules (H-H - - - - H-H)

IM forces are responsible for the physical properties

3 types of intermolecular forces (in order of increasing strength of attraction):

  • London forces
  • Permanent dipole-dipole interactions
  • Hydrogen bonding
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6.3 Intermolecular forces

London forces:

  • Exist between all molecules
  • Only important between atoms (e.g. noble gases) or pure covalent molecules (e.g. H2)

How they arise:

Electrons are constantly moving causing a temporary dipole in one molecule which induces and dipole on a neighbouring molecule

The induced dipole induces further dipoles on neighbouring molecules which then attract one another

Strength: the more electrons in each molecule... (stronger forces)

  • the larger the temporary and induced dipoles
  • the greater the induced dipole-dipole interactions
  • the stronger the attractive forces between molecules
  • the larger the temperature required to overcome the forces
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6.3 Intermolecular forces

Permanent dipole-dipole interactions:

Occurs between permanent dipoles in different (polar) molecules

The larger the difference in EN, the larger the dipole and the stronger the attraction

Simple molecular substances:

As a solid, molecules are arranged in a simple molecular lattice

Weak intermolecular forces between molecules

Strong covalent bonds between atoms within the molecule

Melting / boiling point:

Weak IM forces, so only need a small quantity of energy to overcome the forces

Low melting and boiling points

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6.3 Intermolecular forces

Solubility of non-polar simple molecular substances

Non-polar simple molecular substances are soluble in non-polar solvents

  • Non polar substance in a non-polar solvent
  • IM forces form between substance molecules and the solvent
  • these interactions weaken the IM forces in the simple molecular lattice
  • the IM forces in the lattice break and the substance dissolves

Non-polar simple molecular substances are insoluble in polar solvents

  • Non polar substance in a polar solvent
  • there is little interaction between molecules in lattice and solvent molecules
  • the IM forces within the solvent is too strong to be broken
  • so the lattice doesn't break down or dissolve
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6.3 Intermolecular forces 6.4 Hydrogen bonding

Solubility of polar simple molecular substances:

  • Polar substances may dissolve in polar solvents
  • Molecules of each attract in the same way as in the dissolving of ionic compounds
  • Solubility depends on dipole strength so hard to predict e.g. ethanol contains both polar and non-polar parts so can dissolve in polar and non-polar solvents

Electrical conductivity: there are no mobile charged particles; SM substances can't conduct

Hydrogen bonding: 

the attraction between the positive dipole on a hydrogen atom and the lone pair of electrons on a negative atom

  • Hydrogen must be bonded to N, O or F      H d+                       H d+
  • Strongest type of IM force                        I                             I
  • Hydrogen bond                                    ** O d- -- H d+ - - - - ** O d- -- H d+
  • Covalent bond                                         **                           **
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6.4 Hydrogen bonding

Ice is less dense than water

When water freezes, hydrogen bonds hold molecules apart in an open lattice

Water molecules in ice are further apart than in water so ice is less dense / floats

The holes in the open lattice are larger and decrease the density of water on freezing

When ice melts, the lattice collapses and the molecules move closer together (denser)

There are 2 lone pairs (on O) per H2O molecule and 2 hydrogen atoms per H2O molecule so each H2O molecule can form 4 hydrogen bonds

Melting and boiling point

  • Higher than expected because water has hydrogen bonding
  • Hydrogen bonding is the strongest intermolecular force
  • Requires more energy to overcome the bonds than if the were just london forces
  • When ice melts, the lattice breaks down
  • When water boiling the hydrogen bonds break completely
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