Lewis Structures and Molecular ShapesBackground
Non-metal atoms bond covalently, resulting in the formation of either neutral molecules or polyatomic ions. A
covalent bond is formed when non-metal atoms share their valence electrons, which they do in order to achieve
filled valence orbitals like their nearest noble gas neighbor. This means that most bonded non-metal atoms will
acquire a total of eight valence electrons via the sharing process – often referred to as the octet rule. A notable
exception is hydrogen, which only needs to acquire two electrons to be like its nearest noble gas neighbor, helium.
A Lewis Structure is a representation of covalent molecules (or polyatomic ions) where all the valence electrons
are shown distributed about the bonded atoms as either shared electron pairs (bond pairs) or unshared electron
pairs (lone pairs). A shared pair of electrons is represented as a short line (a single bond). Sometimes atoms can
share two pairs of electrons, represented by two short lines (a double bond). Atoms can even share three pairs of
electrons, represented by three short lines (a triple bond). Pairs of dots are used to represent lone pair electrons.
The rules for drawing Lewis structures can be found in the Procedure Section of this handout.
The shape of a molecule depends on the distribution of atoms in space about the central atom, and their bond
angles. Bond pair electrons and lone pair electrons repel one another, thus they will be arranged around a central
atom as far apart as possible in order to minimize repulsions. This is known as Valence Shell
Electron Pair Repulsion theory, or VSEPR theory.
The following VSEPR table shows the most common molecular shapes that you will encounter. More molecular
geometries do exist; however, they are beyond the scope of this course.
Only two outer atoms surround the central atom. There are
no lone pairs on the central atom. Outer atoms are arranged
opposite to each other. The bond angles are exactly 180.
Three outer atoms surround the central atom. There are no
lone pairs on the central atom. The central and outer atoms
all lie in the same plane (molecule is flat). Bond angles are
Two outer atoms and one lone pair surround the central
atom. Bond angles are slightly less than 120.
Four outer atoms surround the central atom. There are no
lone pairs on the central atom. The four outer atoms are
evenly arranged in 3D around the central atom as if at the
corners of a regular tetrahedron. The bond angles are exactly
Three outer atoms and one lone pair surround the central
atom. Here the central atom is located slightly above the
three outer atoms, like a tripod. The bond angles are slightly
less than 109.5.
Two outer atoms and two lone pairs surround the central
atom. Bond angles are slightly less than 109.5.
Drawing Lewis Structures
1. Draw Lewis structures for CH4, NH3, H2O, CO2, H2CCH2, . Clearly show all bond pair electrons as lines and
lone pair electrons as pairs of dots.
Rules for Drawing Lewis Structures
1. Total the number of valence electrons that each atom contributes to the molecule/polyatomic ion.
The quickest way is to find the group number for each atom.
For polyatomic anions, you must add electrons (equal to the negative charge) to the total number of
valence electrons. For polyatomic cations, you must subtract electrons (equal to the positive charge)
from the total number of valence electrons.
2. Draw a stick structure for the molecule.
Most molecules/polyatomic ions consist of one central atom bonded to 2, 3 or 4 other atoms.
The least electronegative atom is the central atom. Hydrogen is the only exception to this, as it
forms only 1 bond. The central atom will usually need to form multiple bonds.
The other atoms are arranged around the central atom, and are attached to it via single bonds.
3. The octet rule must be obeyed for all elements except hydrogen (follows a “duet” rule).
Each bond in the stick structure contains two electrons, which must be subtracted from the
total number of valence electrons.
Starting with the outside atoms, place the remaining electrons around each atom until it has a total of
8 electrons (except hydrogen – it only requires 2 electrons).
If there are not enough electrons available to obey the octet rule using single bonds, this indicates
that double or triple bonds between two atoms are required in your structure. If short by two
electrons, try a double bond, and if short by four electrons, try a triple bond or two double bonds.
Constructing Models, Determining Molecular Shapes
1. Use your molecular model kit to construct a three-dimensional model of each of these molecules and
polyatomic ions. Sketch a reasonably detailed picture of this model.
Rules for Constructing Molecules with the Model Kit
1. Each colored ball in your kit corresponds to different atom or different group of atoms. Refer to the
inside of the container lid for the correct color key.
2. Use the short sticks for single bonds.
3. Use two long flexible sticks for a double bond and three long flexible sticks for a triple bond.
2. Compare each model constructed to the molecular shapes described in the Theory section. Then identify and
record its correct shape name and its bond angles. If the molecule has no definite central atom, then only
record the bond angles.
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