Octet Rule and Lewis Dot Structures
CH2000: Introduction to General Chemistry, Plymouth State University, Fall 2014
Introduction:
In the late 1800s, chemists were working to make sense of the large numbers of compounds and
ions which they had discovered. They compared many formulas and found that for main group
elements, such as H, C, N, O, F, Cl, P, S and Si, neutral compounds could almost always be described
by assuming that each element made a certain number of “bonds” – one bond each for H, F, Cl, Br and
I, two for O and S three for N and P, etc. They began to make line drawings to indicate bonds:
By the early 20th century, it became clear that numbers of bonds had something to do with the
numbers of electrons. Since the patterns repeated from row to row in the periodic table, not all
electrons were important in deciding what was happening. Gilbert Lewis introduced the idea of
“valence level electrons” (we normally just say “valence electrons”). He showed that if a bond contains
two electrons, the numbers of bonds and ion charges we observe for main group elements mostly
follow an “octet rule.” In other words, if you only count electrons in the current row of the periodic
table around an atom (the valence electrons), these plus any others they share in bonds add up to 8 e-.
(For H and He, the total is 2 e-.) (We now know that the number 8 is special because it is the number of
electrons needed to fill all the s and p atomic orbitals at each principal energy level--we will discuss
this later in class). Why is this? Quick test: count the valence electrons in the noble gases Ne and Ar.
You should see that there are 8. In fact, [ignoring the "metal" block in the middle of the table] you
cannot find an element with more than 8 valence electrons. Atoms, molecules and ions with a full set
of “outer shell” electrons have a special stability we call "noble gas configuration."
In order for elements to have an octet of electrons, they often need to have electrons that are not in
bonds. These are known as "non-bonding electrons" and they generally reside in pairs. Thus, the
correct Lewis structures of the molecules above, with the central atoms fulfilling the octet rule
(hydrogen has its own special version of the octet rule: only two electrons) would be:
Note that each line between atoms represents a bond, which contains 2 electrons.
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Steps to Drawing Lewis Dot Structures:
1. Count the total number of valence electrons in the molecule (make a little table to help you do
this!). If the "molecule" is an ion, add or subtract electrons as appropriate.
2. Build the "skeleton" of the molecule. Draw ONLY one bond between neighboring atoms
a) The first atom listed in a molecular formula is usually the central atom (for example, in NH3
the "N" is the central atom to which all the H's are bound).
b) The least electronegative atom is often the central atom (for example, in NO2, the N is the
central atom).
c) Hydrogen (H) will only bond to ONE other atom. Hydrogen is NEVER the central atom.
3. On the skeleton structure, add all the remaining valence electrons as non-bonding electrons to
fulfill the octet rule to the extent possible.
a) Work from the outside atoms to the inside atoms. If you have any electrons left, add them
to the central atom.
b) Do NOT place more total electrons into your structure than you counted in Step 1!! EVER!
c) Do NOT have more than one bond [TWO electrons total] to hydrogen!! EVER!
d) If you have run out of electrons, but the central atom does not have an octet, do NOT add
more electrons (see 3b above). Move on the step 4.
e) Do NOT start drawing double or triple bonds onto your skeleton structure.
4. If you have used all the valence electrons and all the atoms have octets, your skeleton structure
is the Lewis dot structure; move to step 5. If not:
a) If the central atom has more than 8 electrons, check it's location on the periodic table. If it
is in the second row (i.e. C, N, O, etc) it CANNOT have more than 8 electrons. Your
structure has a mistake. If the element is in a lower row (e.g. S, As, I), it is allowed to have
an expanded octet. You may be OK. Move to step 5.
b) If the central atom has less than an octet, but one or more of its neighbors has a lone pair of
electrons, you will need to draw new structures that convert those lone pairs into bonding
pairs. You will very often need to draw more than one structure. Leave the skeleton
structure alone to act as a guide for drawing these. Draw every possible new structure
separately showing double bonds to the atoms with lone pairs of electrons (remember to
"remove" the lone pair that is moved into a bond).
c) Some elements are OK with less than 8 electrons, including those in the first 2 columns (e.g.
H, He, Li, Be, Na, etc.) and in the boron (B) column.
5. Determine the formal charges of all the atoms in all the structures you have. The formula for
formal charge is:
Valence electrons - (total # non-bonding electrons + number of bonds) = formal charge
(note: "number of bonds" is technically "half the number of bonding electrons", but it is easier
to count bonds than to count bonding electrons ,then divide by 2; you get the same answer in
the end!)
a) All the formal charges should sum up to the total charge on the ion. Check this!
b) The "best" structure is the one with the most "zero" charges. Formal charges of 1+ and 1are OK, but higher charges (2 +/-, 3 +/-) usually are not. Avoid large charges as much as
possible.
© Copyright Plymouth State University and Jeremiah Duncan. May be distributed freely for education purposes only.
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6. If there are several reasonable structures that all have the same formal charges and only differ
by the placement of a double or triple bond, you have identified resonance (alternate structures).
All resonance structures MUST be listed as the answer.
THERE ARE MANY EXCEPTIONS TO THE OCTET RULE!! The best way to learn is to practice:
Exercises
1. Draw the Lewis dot structures of the following. Show your valence electron count and
calculate the formal charge of all the atoms in each structure. If alternate structures (resonance)
exists, show all the possible resonance structures with resonance arrows between them.
(a) PCl4+
(b) XeF62+
(c) BBr3
(d) AlCl3
(e) CN-
(f) COF2
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2. (a) Fill in the Lewis dot symbols for H, C, N, O and Cl. Based on this information predict how
many bonds and how many lone pairs each of these atoms will possess.
element
Lewis symbol
bonds
lone pairs
H
C
N
O
Cl
(b) Based on your answers in the table, complete the structure of Tylenol (below). Show all
multiple bonds and lone pairs.
3. Draw three structures for NO (two with double bonds and one with a triple bond). Determine
the formal charge for the atoms in each structure. Determine which structure contributes most
to the true bonding picture of NO.
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4. Draw all possible resonance structures for each of the following. Determine the formal charge
for each atom in the molecules, and use this information to determine whether the resonance
structure contributes significantly to the structure of the molecule.
(a) N2O
(b) NO3-
(c) HNO3
(d) CO32-
(e) HCO3-
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(f) SO3
(g) SO42-
(h) HSO4-
5. Draw the structures for O2NNCO and ONNCO2. Determine whether one is more likely or both
are equally likely.
NOTE: Next week's work depends on answers in this worksheet. Complete this and bring it to class.
© Copyright Plymouth State University and Jeremiah Duncan. May be distributed freely for education purposes only.
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