Understanding Lewis Dot Structures  - Sam Hart, 1997 - hart@physics.arizona.edu
-------------------------------------------------------------------------------

Lewis Dot structures are a method by which chemists can visual the electron and molecular
structures of simple compounds. The procedure (with examples) for generating a Lewis Dot structure
is as follows. (If any of this is incorrect, please e-mail me at "hart@physics.arizona.edu".)


- Given a molecular formula, first take all of the elements included in the compound and calculate
the total number of valence electrons. Valence electrons are those that are in the outermost shell
of an atom. This number also corresponds to the group number of the element in question. For
example, Oxygen (O) is located in the 6A group. Thus, its outermost shell can be said to contain 6
electrons. This calculation must be done for each element in the compound,

	CO    - Carbon Dioxide
	  2

	1 4(C)  = 4
	2 6(O)  = 12
	------------
	          16 valence electrons

- If the molecule is an anion, then the ion charge of needs to be added to previous sum. If it is
a cation, then the ion charge needs to be subtracted from the previous sum. Recall that an electron
is negatively charged, and when more electrons are added, you get a higher negative charge:

	1(e-) + 2(e-) = 3(e-)

Conversely, when electrons are removed, than you get a lower negative charge:

	3(e-) - 1(e-) = 2(e-)

If the previous Carbon Dioxide molecule was a negatively charged ion,

	  -
	CO 
	  2

Then the additional electron (e-) needs to be added to the aforementionned valence electron sum,

	1 4(C)  = 4
	2 6(O)  = 12
	------------
	          16 e-
	         + 1 e-
	------------
	          17 e-

Alternately, if the Carbon Dioxide molecule was a positively charged ion,

	  +
	CO
	  2

Then the missing electron (e-) must be removed from the valence electron sum,

	1 4(C)  = 4
	2 6(O)  = 12
	------------
	          16 e-
	         - 1 e-
	------------
	          15 e-

If, of course, the compound is not an ion, then nothing is added or subtracted,

	------------
	          16 e-
	           0 e-
	------------
	          16 e-

- Now arbitrarily draw the molecule, connecting the elements by lines, bearing in mind the
following guidelines:

	-Hydrogen (H) is always an end or terminal atom. Because it only allows for one more
electron in its outer shell, it can only be connected to one other element (well, for these
purposes here, at least.)

		Example: H O       ->   H -- O -- H
		          2

	-The atom of lowest electron affinity in the molecule or ion is generally the central
atom. (The central atom is the atom of lowest electronegativity.)

	-Bonds such as C=C, C=N, C=O, S=O, P=O, and so on, are quite common.

	Example: CO
	           2

	O -- C -- O

-Each line on the previous molecule symbolizes a covalent bond, where the elements share electrons.
Each line contains two (2) electrons (e-).

-Count up how many electrons exist in your drawing (recalling that each bond [line] represents
two electrons (e-).) Subtract this from your previous sum of valence electrons.

	O -- C -- O      -> 2 electrons/bond X 2 bonds = 4 electrons

	-----------
	         16 e-
	        - 4 e- (accounted for in structure)
	-----------
	         12 e- (left to place)

-Take the remaining electrons and place lone pairs about each terminal atom (except H) to satisfy the octet rule:

	Octet Rule: Because we will be filling up the outer shell, this will give the element(s)
	the electron configuration of the next noble gas on the periodic table. Since all noble
	gases (except He) have eight valence electrons, there will have to be eight electrons
	around each element to give it a similar outer shell configuration.


	..             ..
	:O  --  C  --  O:
	''             ''

	(Note: If the central atom is from the third or higher period, it can accomodate more
	than four electron pairs.)

-If the central atom is not yet surrounded by four electron pairs, convert one or more terminal
atom pairs to another bond. Keep altering the original structure until all elements have four
electron pairs. (There are exceptions to this rule, but it is a good guideline.)

	..             ..
	 O ==== C ==== O
        ''             ''

                                _______________________________
                                -EXCEPTIONS TO THE OCTET RULE:-
                                -------------------------------

-There are compounds that have fewer than four pairs of electrons around the central atom.
Consider BeH  which only can be configured in the following manner,
            2

	H - Be - H

Also, BH  can only be configured thusly,
        3

	    H
	    |
	H - B - H

In BeH  there were only two pairs of electrons around Be, and in BH  only three are around B.
      2                                                            3

This is typical of chemistry involing boron, and makes many boron compounds highly reactive. The boron
will accomodate a fourth pair when it is accompanied by another atom. This means that compounds such as
BH  react readily with another molecule that has a non-binding electron pair.
  3
*A covalent bond in which the bonding pair originates on one of the bonded atoms (as in this case) is
called a _coordinate_covalent_bond._



-If an element of the third or higher periods is the central atom in a molecule or ion, it can be
surrounded by more than four valence electron pairs because such elements have d oribitals available.

	Example: F P
	          5

	..  .. ..
	:F :F: F:
	''\ | /''
	   \|/
	    P
	 ../ \..
	 :F   F:
	 ''   ''


See Also "Resonance Structures"