A level Chemistry exam revision resources





 
Molecules of Methane, CH4, ammonia, NH3, water, H2O and hydrogen fluoride, HF

The directional nature of covalent bonds is shown in the diagrams of molecules above. The shape of the methane molecule is tetrahedral because the four bonding pairs of electrons repel each other equally, and the equilibrium position of all four bonding electron pairs is tetrahedral.

It is possible to work out the shape of a small molecule that has a formula XYn by applying a few simple rules. We will use ammonia as an example to illustrate the idea.

 
  • Rule 1. First find the number of bonding pairs of electrons in the molecule. The number of bonding pairs of electrons in the molecule NH3 can be seen in the formula. There must be three bonding pairs of electrons holding the three hydrogens onto the nitrogen.

  • Rule 2. Find the number of valence electrons (electrons in the outer energy level) on an atom of the central atom (The one of which there is only one.) Nitrogen is in group V, so the nitrogen has five electrons in the outer energy level.

  • Rule 3. Find the number of lone pairs on the central atom by subtracting the number of bonding pairs (3) from the valence electrons (5) to find the number of electrons (2) that will make up lone pairs of electrons. Divide this number by 2 to find the number of lone pairs, (2/2 = 1).

  • Rule 4. Distribute all the electron pairs around the central atom and learn the angles they will make from molecules with no lone pairs.

  • Rule 5. Learn that the repulsion between lone pairs of electrons is greater than the repulsion between bonding pairs, and subtract 2o from the bond angles for every lone pair.

  • Rule 6. Learn the names of the shapes. The shapes are named form the position of the atoms and not the position of the orbitals.

The shape of ammonia
The Shape of Ammonia

Table of Shapes

Table showing shapes of molecules

There is one more rule to learn, and it concerns the shape of polyatomic ions.

 
  • Rule 2a. If the molecule is an ion, e.g. ammonium (NH4+), subtract 1 from the number of valence electrons for every + charge on the ion and add 1 to the valence number for every - charge, then proceed as before.

 

Some more Examples



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