Lewis Dot Diagrams

Try the activity with 30 practice molecules, which can be found here.

The elecrtrons in atoms can be split into two main categories. First, the core electrons which do not participate in forming bonds to other atoms. Second, the valence electrons, which are involved in bonding. The number of valence electrons, which equals the number in the outermost shell of an atom, determines how an atom bonds to other atoms.

lewis.dots.PTOE (19K)
I borrowed this image from HyperPhysics.

The number of valence electrons an atom has is determined by the arrangement of the electrons shells and the total number of electrons found in a neutral atom. The complete description of how this is determined is unecessary when the aim is to be able to draw molecular structures. A simple way to count valence electrons is to look at the group number. In the figure at right the old group numbers are given as roman numerals. The number of dots arranged around each atomic symbol gives the number of valence electrons. Note that the value of the group number and the number of valence electrons is equal. Modern periodic tables have changed group number labels so that groups III – VIII (zero in this figure) are now labeled as groups 13 – 18. Still, the pattern— once seen— is easy to remember.

Atoms form bonds with one another based on the number of valence electrons they have. The general rule is known as the ‘octet rule’. This rule governs the number of electrons an atom needs to borrow from other atoms in order to have a full complement of eight electrons in its outermost shell. The reason for this rule is that eight electrons is a sort of magic number. When an atom has eight valence electrons it is in a lower energy state than when it does not. An atom in a Lewis dot structure has an octet when the sum of shared and un-shared electrons equals eight.

Conventionally, electrons are considered in pairs (an exception is a molecular radical with an odd number of electrons). Electron pairs in a bond are represented as a line between the atomic symbols: H—O—H. Non-bonding pairs (also called lone pairs) are represented by a pair of dots near the atomic symbol they belong to (for example, :N—F3). Lone pair (also called non-bonding pair) electrons are valence shell electrons that do not participate in bonding.

Given a molecular formula, here are the steps to follow in drawing a Lewis Dot Diagram for the molecule. Read and re-read them so that you learn them in all their details.

  1. Count the number of valence electrons for each of the atoms in the molecule and calculate the total. The number of valence electrons an atom has depends on the group number. Group 1 elements (H, Li, Na, etc.) have 1 electron. See the image above for other groups.
  2. If the molecule has a negative charge add enough electrons to account for it. If the molecule has a positive charge, subtract the number of electrons needed to create that charge.
  3. Organize the atoms around a central atom. Hydrogen is never the central atom. Atoms other than the central atom are referred to as ligands.
  4. Draw a single line between each ligand atom and the central atom to represent bonds between those atoms. Each bond requires one electron pair.
  5. Add lone pairs of electrons to each ligand atom as required to make an octet, giving each atom a noble gas electron configuration. Add remaining electrons as lone pairs on the central atom. Remember not to exceed the number of electrons you calculated earlier!
  6. Redistribute lone pairs as double- or triple-bonds if the central atom does not have an octet. Do not move lone pair electrons from ligand halogens.
  7. Double-check: do all atoms have an octet? If not, rearrange your electrons to make it so. If necessary, re-draw the skeleton of your molecule to make it possible to give every atom an octet.
  8. Formal charge will determine if your electron distribution is correct. Calculate it by comparing the number of electrons that belong to an atom to the number of valence electrons for that atom. Do this as follows:
    1. Count one half of the total number of bonding electrons attached to the atom.
    2. Count the total number of lone pair electrons.
    3. Add up the previous two quantities and subtract from the number of valence electrons. This is formal charge.

An easy way to find formal charge is to draw a circle around an atom in a Lewis Structure and count the electrons inside the circle (counting only half of each bond as being inside the circle). Subtract the result from the number of valence electrons to find formal charge.

Note formal charges next to atoms by using a small + or – number in a circle.

Total formal charge for a molecule must equal the overall charge of the molecule.

Minimize formal charges so that they are all as small as possible and so that as few atoms as possible have formal charges. For central atoms from period 3 or higher it is permitted to allow further bonds that exceed an octet as long as the formal charge is zero or at most +1.

Formal charge is simply a way to decide whether electrons are equally or unequally shared in a molecule. When formal charge is zero it means that the number of electrons shared or owned by an atom equals the number of valence electrons. When an atom has a formal charge that is negative it means that the atom has taken electrons away from its neighbors and is sharing unfairly by taking more electrons. Usually atoms with high electronegativity will do this. In other words, non-metals and especially non-metals at the furthest right-hand extreme of the periodic table. When an atom has a formal charge that is positive it means that the atom has given away electrons to its neighbors and is sharing more generously. Usually atoms with low electronegativity will do this. In other words, metals may be found to occur in some molecules with a positive formal charge.

Not all molecular model drawings (Lewis diagrams) require the calculation of formal charge. You only need to calculate formal charge under any one of the following conditions:

  1. Calculate formal charge on all atoms in a molecule that is a polyatomic ion. Remember, formal charges must all add up to the overall ionic charge.
  2. Calculate formal charge on any atom in a molecule that violates the octet rule. Some molecules cannot be diagrammed without giving an atom more or fewer than 8 electrons. When this is the case the diagram is most likely to be correct when the formal charge is zero. Sometimes, the formal charge may be allowed to be +1 but only rarely is it higher. When the central atom which violates the octet rule has a positive formal charge other atoms in the molecule will likely have a compensating negative formal charge.
  3. Calculate formal charge, finally, when at atom does not follow its usual bonding pattern. In the table below you will find the symbols for atoms from Period 2 along with the number of bonds they are usually found to make and the number of non-bonding pairs (lone pairs) of electrons they usually have. Use this as a guide when drawing molecules and when a molecule seems to require a deviation from the usual bonding pattern you should calculate formal charge.
Usual Bonding Patterns
Group Number Number of
Valence Electrons
Example Atoms Number of
Bonds
Number of
Lone Pairs
Examples
1 1 H Li Na K 1 0 Usual.Bonding.Li (1K)
2 2 Be Mg Ca 2 0 Usual.Bonding.Mg (1K)
13 3 B Al 3 0 Usual.Bonding.Al (1K)
14 4 C Si 4 0 Usual.Bonding.C (1K)
15 5 N P As 3 1 Usual.Bonding.N (1K)
16 6 O S Se 2 2 Usual.Bonding.O (1K)
17 7 F Cl Br I 1 3 Usual.Bonding.F (1K)
Three molecules with a total of 8 valence electrons and one with 16.
Lewis.Dot.Example.Molecules (6K)
I based my original set of rules for drawing structures on the rules I learned as a student at the University of Southern Maine. I have modified them (May 1, 2012) according to the following JChemEd article:
Lewis structures, formal charge, and oxidation numbers: A more user-friendly approach
John E. Packer and Sheila D. Woodgate
J. Chem. Educ., 1991, 68 (6), p 456
A useful supplement to this page (an alternate way to draw structures) can be found here
Lewis Diagrams Exercises
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