1.2 Lewis structures – the next step
You will recall Lewis structures, developed by American chemist G.N. Lewis in 1916, are essentially drawings of where the atoms comprising a molecule are, and where their valence electrons are located. And, later in this course you will find they are used to predict the shape of molecules and infer a lot about the chemical and physical properties of compounds.
Essentially, when constructing a Lewis structure, the shared electron-pair bond is used to re-express structural formulas in an electronic form in which each atom has the shell structure of a noble gas.
Consider for example ammonia (NH3). From the structural formula you know there is a central nitrogen atom to which is attached three hydrogen atoms.
How many electrons are there in the valence shell of nitrogen?
Five (2s2 2p3)
So in order to achieve the shell structure of the nearest noble gas – neon, a nitrogen atom must gain three electrons. Likewise hydrogen attains the shell structure of helium, with 2 outer electrons.
Therefore, if nitrogen shared three of its valence electrons with hydrogen atoms, it can indeed acquire an octet in its valence shell. This is shown in structure 4.14.
Note that the electron pairs in these Lewis structures are of two types.
The pairs shared between atoms represent chemical bonds and are called bond pairs. But there are also pairs that remain on just one atom and are unshared. These are called non-bonded electron pairs or lone pairs.
How many bond pairs and non-bonded electron pairs are there in the ammonia molecule?
NH3 contains three bond pairs and one non-bonded pair.
Now let’s turn to another simple molecule, water (H2O).
Draw the Lewis structure for the water molecule (H2O). How many bond pairs and non-bonded pairs are there?
Oxygen has six electrons in its valence shell, and so it can acquire an octet by sharing electrons with two hydrogen atoms, which in turn fills its valence shell, see structure 4.15 resulting in two bond pairs and two non-bonded pairs.