Atoms bond chemically to form molecules. Lewis structures are a way to represent this bonding on two dimensional paper and determine the molecular geometry of a structure.
Review of bonding
Covalent molecules share electrons while ionic compounds transfer electrons from one atom to another.
Lewis Structures of atoms
The element symbol is drawn to represent the nucleus and core electrons. The valance electrons are drawn around the symbol—one on each side before doubling up.
Exceptions to the Octet Rule
Most atoms are the most stable with 8 electrons in their valence shell, and will bond until this is reached. However, hydrogen and helium can only hold 2 electrons in their valence shell. Boron and Beryllium can be stable with only 6 valence electrons. Any element in the third row or below can hold more than 8 in the empty d subshells.
Arranging atoms in a Lewis Structure
It is often difficult to know in what order to place the atoms. There are some general rules that can be followed:
- For molecules with only 2 elements, arrange the atoms symmetrically
- “COOH” is a carboxylic acid (both O’s bond to the C and the H goes on one of the O’s)
- Hydrogen and halogens cannot go in the middle
- Write the remaining atoms in the order they appear in the formula
- Write the hydrogen and halogen atoms around the element they are written next to in the formula
Drawing Lewis Structures for covalent compounds
Once the atoms are arranged, a system can be used to complete the Lewis Structure:
- Arrange the atoms as above
- Determine the # of valence electrons for each atom
- Draw the valence electrons—do not double up where a bond is going to form between two atoms
- Count to see if all atoms have full valences
- If two atoms adjacent to each other do not have full valences, move in an electron from each to form a double bond. Repeat for triple bond if necessary.
- If two atoms that are not adjacent to each other need to double bond, try moving a hydrogen to one of them to cause two atoms adjacent to each other to need the double bond.
Another approach to drawing Lewis Structures
There is a second method that is also commonly used to arrive at the same structure:
- Arrange the atoms as above.
- Determine the total # of valence electrons for the whole molecule
- Put one bonding pair between each set of atoms to be bonded.
- Place remaining electrons in lone pairs, starting with the most electronegative element.
- If atoms do not have full valence shells, move a lone pair from an adjacent atom in to double, or triple, bond.
Ionic Structures
Ionic bonds are formed from the transfer of electrons from the metal atom to a non-metal atom or polyatomic ion. When drawing ionic structures, do not draw the atoms as sharing the electrons. Rather, remove the electrons from the
Valence Shell Electron Pair Repulsion Theory
Bonds are made of electrons and electrons are negative and therefore repel each other. Bonds and lone pairs form as far apart from each other as possible. This theory can be used to determine the electron structure (the 3D shape based upon electron regions—bonding regions and lone pair regions—of the central atom) or molecule structure (the 3D shape based on the electron regions, but named after the bonded atoms only).
A = central atom; X = ligands; E = lone pairs
Electron regions |
Molecular Formula |
Name |
2 |
AX2 |
Linear |
3 |
AX3 |
Trigonal Planar |
3 |
AX2E |
Bent |
4 |
AX4 |
Tetrahedron |
4 |
AX3E |
Trigonal pyramidal |
4 |
AX2E2 |
Bent |
5 |
AX5 |
Trigonal bipyramidal |
5 |
AX4E |
See-saw |
5 |
AX3 E2 |
T-shaped |
5 |
AX2E3 |
Linear |
6 |
AX6 |
Octahedron |
6 |
AX5E |
Square pyramidal |
6 |
AX4 E2 |
Square planar |
