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Introductory Chemistry
Drawing Lewis
Structures
Demonstration of process: Lets draw the Lewis structure for the carbonate ion, CO32- |
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Step 2: Calculate the
total number of valence electrons contributed by
all the atoms.
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Step 3: Check to see if
there is a charge on the formula unit. Adjust the
total in Step 2 by the number of electrons added or
removed because of any electrical charge. This
gives you the final total of valence electrons that
are available to hold the particle
together. total of 4 (carbon)+(3x6)(oxygen)+2(negative charge)=24 valence electrons. |
Step 4: Identify central
atom/atoms in the formula unit. |
Step 5: Think of your
structure as a target. Write down the symbol for
the central atom in the "bullseye". Then arrange
the symbols for the oxygen atoms evenly in the next
ring around the center atom. |
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Step 6: Divide the total
number of valence electrons for the formula unit
(Step 3) by 2 so you know the number of "electron
pairs" for the particle. For example, in carbonate ion, Steps 1, 2, and 3 would tell you that the molecule contains 24 valence electrons. Therefore its bond, octets, and duets are done with 12 pairs of electrons. |
Step 7: Connect the center
atom with a single bond (1 pair of electrons) to
each neighboring atom in the second ring. Then,
using a single bond, connect each atom in the
second ring to its neighbor in the third ring.
Count the number of electron pairs used.
For example, in carbonate ion you would connect the C to
each of the O's. This would use a total of 3 electron
pairs.
Step 8: Subtract the
number of electron pairs used in Step 7 from the
total number of electron pairs for the particle
(Step 6). Use these remaining electron pairs to
fill out the octets for atoms in the second
ring. Step 9: The almost last
step is verification that the atoms each have an
octet or duet. If they do you are finished.
If the central atom lacks
an octet when only single bonds are used you need
to move a lone electron pair to make one or more
multiple bonds. This may mean double or triple
bonds as appropriate. NOTE: Any lone pair can be
used here. The negative two charge is spread
over the entire ion.
For example, in carbonate ion you used 3 pairs of
electrons in Step 7, leaving