Module 2, Chapter 6

6.1 Shapes of Molecules and Ions

Electron-pair Repulsion Theory

An electron has a negative charge, so electron pairs repel one another. The electron-pair repulsion theory predicts and explains the shapes of molecules.

  • The electron pairs surrounding a central atom determine the shape of the molecule/ion.
  • The electron pairs repel one another so they are arranged as far apart as possible.
  • The arrangement of electron pairs minimises repulsion so holds the bonded atoms in a definite shape. So different numbers of electron pairs result in different shapes.

Bonding-pair and Lone-pair Repulsions

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.

Relative repulsions between lone pairs and bonding pairs: BP/BP < BP/LP < LP/LP

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6.1 Shapes of Molecules and Ions

Shapes of Molecules

  • Linear:                         2BP          180 degrees                          CO2       
  • Trigonal Planar:           3BP          120 degrees                          BF3       
  • Trigonal bipyramidal:  5BP          90 & 120 degrees                   PF5       
  • Octahedral:                 6BP           90 degrees                            SF6       
  • Tetrahedral:                 4BP          109.5 degrees                       CH4       
  • Pyramidal:                   3BP 1LP   107 degrees                          NH3       
  • Non-linear:                  2BP  2LP  104.5 degrees                       H2O      
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6.1 Shapes of Molecules and Ions

Ammonium ion

NH4+ has the same number of bonded pairs of electrons around the central atom as a methane molecule. 

It has 4 bonding pairs, therefore, is tetrahedral shaped with bond angles of 109.5 degrees.

Carbonate, nitrate, and sulphate ions

Carbonate ions, CO3- , and nitrate ions, NO3- , have 3 regions of electron density surrounding the centre atom. They have the same shape as a BFmolecule; they are trigonal planar with bond angles of 120 degrees.

Sulphate ions, SO42-, have four areas of electron density around the central sulphur atom the same as a methane molecule. They are tetrahedral shaped with bond angles of 109.5 degrees.

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6.2 Electronegativity and Polarity

Electronegativity

In molecules of elements, the atoms are the same element and the bonded electron pair is shared evenly. When the bonded atoms are different elements:

  • the nuclear charges are different
  • the atoms may be different sizes
  • the shared pair of electrons may be closer to one nucleus than the other.

The shared pair of electrons in the covalent bond may now experience more attraction from one bonded atoms than the other.

Pauling Scale

The scale used to compare the electronegativity of atoms of different elements. Electronegativity increases as the nuclear charge increases and the atomic radius decreases. Therefore fluorine is the most electronegative element with a Pauling Value of 4.0.

The difference in electronegativity affects whether the bond is covalent (0), polar covalent (0 to 1.8), or ionic (>1.8) 

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6.2 Electronegativity and Polarity

Non-polar Bonds

  • The bonded pair of electrons is shared equally between bonded atoms.
  • Occur when bonded atoms are the same, or have the same/very similar electronegativity

Polar Bonds

  • The bonded pair of electrons is shared unequally between the bonded atoms.
  • Occurs when the bonded atoms are different and have different electronegativity values.
  • When the molecule becomes polarised, the opposite charges separate and form a permanent dipole. 

Polar Molecules

  • In a molecule with more than 2 atoms, there may be two or more polar bonds.
  • These dipoles may reinforce each other or they may cancel each other out, depending on the shape of the molecule.
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6.3 Intermolecular Forces

Induced dipole-dipole Interactions (London Forces)

London forces are weak intermolecular forces which exist in all molecules.

  • An instantaneous dipole is formed which exists with a constantly shifting position.
  • This induces a dipole on a neighbouring molecule, which induces dipoles on its neighbouring molecules which attract one another.

The more electrons in each molecule:

  • The larger the instantaneous and induced dipole.
  • The greater the induced dipole-dipole interactions.
  • The stronger the attractive forces between molecules.
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6.3 Intermolecular Forces

Permanent dipole-dipole Interactions

  • Act between the permanent dipoles in different polar molecules.
  • Can only occur in polar molecules.
  • Permanent dipole-dipole interactions are stronger than london forces. 
  • Molecules with permanent dipole-dipole interactions have london forces too so they have higher melting and boiling points than molecules with just london forces as extra energy is required to break the additional permanent dipole-dipole interactions.

Simple Molecular Substances

  • Made up of small molecules, which in the solid state, form a simple molecular lattice and have weak intermolecular forces and strong forces between atoms within each molecule.
  • Low melting and boiling points; weak intermolecular forces broken whilst covalent bonds are strong and don't break.
  • Non-polar substances will dissolve in non-polar and polar solvents
  • Polar substances will not dissolve in non-polar solvents, as there's not enough attraction to break the 'lattice'.
  • Simple molecular substances don't conduct electricity due to no mobile charged particles.
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6.4 Hydrogen Bonding

Hydrogen Bonding

  • Special type of permanent dipole-dipole interaction between molecules containing:
  • an electronegative atom with a lone pair of electrons (N, O, F)
  • a hydrogen atom attached to an electronegative atom. (H-F, H-O, H-F)
  • The strongest type of intermolecular forces

Water (H2O) - anomalous properties due to hydrogen bonds

Ice (solid) is less dense than water (liquid).

  • Hydrogen bonds hold water molecules apart in an open lattice; can form 4 hydrogen bonds, two lone pairs on oxygen and two hydrogen atoms.
  • The water molecules in ice are further apart than in water because the holes in the lattice structure decrease the density of the water when it freezes. 

Relatively high melting and boiling points

  • Due to extra forces from the hydrogen bonds, more energy is required to break the bonds.
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