Structural formula: A chemical formula which shows the groupings of atoms in a compound.
Lewis structures,chemical formulae (diagrams) which show the arrangement of the atoms and valence electrons of discrete molecules and ions, are an essential part of the chemical language. They are used to show details of structure and of chemical reactions. The ability to write Lewis structures for basic molecules and ions is an essential skill for chemists. The following rules require knowledge only of the number of valence electrons of an atom of each element, and the arrangement of the nuclei, i.e. which atoms are joined or bonded together.
Rules for drawing Lewis structures
- Determine the total number of valence electrons in the molecule or ion by adding together the numbers of valence electrons of each atom, and if an anion, by adding the overall charge of the ion, and if a cation, by subtracting the overall charge of the ion.
- Place the atoms in their relative positions.
- Draw a line representing a single bond containing two electrons between joined atoms.
- Distribute the remaining electrons in pairs on the outer atoms so that these have up to eight electrons (except for hydrogen which has two.) Any still not used after this should be placed on the central atom.
- If the central atom is now surrounded by fewer than eight electrons, move sufficient non-bonding pairs from outer atoms other than halogens to between joined atoms, thus making them bonding, to bring the number on the central atom up to a maximum of eight.
- Count the number of electrons "owned" by each atom assuming bonding electrons are shared evenly. To evaluate the formal charge at that atom, compare the result with the number of valence electrons of the neutral atom. Show only non-zero charges.
- For central atoms from the third or later rows of the periodic table, move further non-bonding pairs to bonding positions to lower the positive formal charge on the central atom to one or zero.
Formal Charge: The electric charge of an atom in a molecule or ion assuming perfect covalent bonding.
Examples of these rules
Example: in NH4+, N "owns" four electrons and has the formal charge of +1 as an N atom has five valence electrons.
Of course, many molecules or ions do not have a "central atom".
Organic molecules with more than one C are obvious examples. Provded the structure is known, (i.e. which atoms are bonded to which), extending the rules with common sense will work.
Example: C2H2, H-C-C-H,
|put pairs of electrons on C atoms evenly,
|move both pairs to between the C atoms to form a triple bond.
Condensed structures. The information provided by a Lewis structure can frequently also be provided by a modified or condensed form that is easily understood once some basic ideas are mastered. Often the non-bonding electrons pairs are not shown. In many organic compounds (carbon compounds), chains of joined carbon atoms are present.
Thus butan-1-ol (1-butanol) can be represented as CH3CH2CH2CH2OH,
where atoms joined to a carbon are
shown to its right before the adjoining atom is given.
Butan-2-ol is CH3CH2CH(OH)CH3.
Ethanoic acid is written as CH3COOH or CH3CO2H.
A degree of condensation appropriate
for the required information can be chosen.
The next chapter has further information of this topic.
The structure of sulfuric acid may be written
as (HO)2SO2, which is much more meaningful than
the usual representation (H2SO4) in that it shows that the
Hs are joined to oxygen.
Familiarity with these types of formulae comes with practice and experience based on the electron structure and valency of the atoms of the elements.
Molecular shape. The shape of a simple molecule or ion describes the relative positions of the nuclei of the atoms in the molecule.
H2O is described as bent because the H-O-H atoms do not lie in a straight line.
NH3 is described as trigonal pyramid because the N atom is above the plane of the three H's which are at the corners of an equilateral triangle.
CH4 is described as tetrahedral because the C atom is in the middle and surrounded tetrahedrally by the four H atoms.
BF3 is trigonal planar, because the boron atom is in the centre of an equilateral triangle with fluorines at each corner.
CO2 and C2H2 are linear the O-C-O atoms and H-C-C-H atoms lying in straight lines.
It is difficult to give a simple description of shape of more complex molecules such as butanol, ethanoic acid and sulfuric acid.
However we can describe the arrangement of the atoms around a single atom.
Thus in butanol above each C atom is surrounded tetrahedrally by four other atoms.
In ethanoic acid the carbon atom bonded to two O atoms and a C is in the centre in an equilateral triangle, the other carbon atom being tetrahedrally surrounded by the three H's and one C.
Valence-shell electron pair repulsion (VSEPR) theory: A theory which gives a simple rationale of shape (reason for the particular arrangement or shape). The "groups" of electrons (i.e. a non-bonding electron pair, a bonding pair, the four electrons of a double bond, the six electrons of a triple bond) around the central atom (as shown in the Lewis structure) keep as far away from each other as possible.
Examples: NH3 is a trigonal pyramid - there are four pairs of electrons tetrahedrally around the N,
but the shape description only involves the atoms;
H2O is bent - there are four pairs of electrons tetrahedrally around the O;
the carbonate ion, CO32¯, is a planar triangle;
sulfur dioxide, SO2 is bent.
In sulfuric acid, the S is tetrahedrally surrounded by four O atoms, and the H-O-S atoms do not lie on a straight line.
Draw Lewis structures for the following species.
Unless otherwise stated, the first element in the formula is central
and individually bonded to the other atoms.
Remember to show non-zero formal charge.
1. tetrafluoromethane, CF4
2. ammonia, NH3
3. sulfur trioxide, SO3
4. sulfite anion, SO32¯
5. nitrate anion, NO3¯
6. nitrite anion, NO2¯
7. ammonium cation, NH4+
8. nitrous oxide, N2O (one N central, linear)
Answers 1 to 8
Draw the full Lewis structures for the following compounds
Answers 9 to 18