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 |
