Drawing Lewis Electron Dot Structure or Formula
A Lewis electron dot structure shows how the atoms of a molecule or an ion share their outermost electrons or valence electrons to form covalent bonds with each other.
It indicates the number of shared and unshared valence electrons.
This sharing of electrons results in their outermost shells being filled up thus attaining electronic configuration similar to those of the noble gases'.
Observe the electronic configurations of the noble gases below.
| He | 1s2 |
| Ne | 1s22s22p6 |
| Ar | 1s22s22p63s23p6 |
| Kr | 1s22s22p63s23p64s23d104p6 |
| Xe | 1s22s22p63s23p64s23d104p65s24d105p6 |
| Rn | 1s22s22p63s23p64s23d104p65s24d105p66s24f145d106p6 |
You can see that their outermost shells are filled up which account for their being highly stable.
Now, compare them to the electronic configurations of several atoms that commonly form covalent bonds with other atoms.
| H | 1s1 |
| C | 1s22s22p2 |
| O | 1s22s22p4 |
| N | 1s22s22p3 |
| S | 1s22s22p63s23p4 |
| B | 1s22s22p1 |
| P | 1s22s22p63s23p3 |
| Si | 1s22s22p63s23p2 |
| F | 1s22s22p5 |
With the exception of hydrogen, they share their valence electrons with other atoms to fill up their outermost shells which contain a maximum of eight electrons. Hence, a Lewis electron dot formula must follow the octet (eight) rule.
Take a look at the following examples.
For our purposes, we're going to use color codes in our illustrations to distinguish the individual atoms with their respective valence electrons.
Lewis Electron Dot Structure of Ethane, C2H6
The electronic configuration of carbon atom shows that it needs four more electrons to fill up its outer shell; the hydrogen atom, on the other hand, needs just one more electron for its lone electron shell.
Each carbon atom satisfies the octet rule by sharing 3 valence electrons with 3 hydrogen atoms and one with another carbon atom. The hydrogen atom fills up its shell by sharing one electron with carbon atom.
Take note that each pair of shared electrons constitue a single bond.
Lewis Electron Dot Structure of Ethylene (Ethene), C2H4
This time, the carbon atoms share 2 electrons with each other and 2 electrons with 2 hydrogen atoms.
Sharing of two pairs of electrons constitute double bond.
Lewis Electron Dot Structure of Acetylene (Ethyne), C2H2
Here the carbon atoms share 3 valence electrons with each other and 1 electron with 1 hydrogen atom.
Sharing of three pairs of electrons constitute triple bond.
So far, we've been looking at the Lewis electron dot structures of molecules. We're going to look at the electron dot structures of some ions.
Lewis Electron Dot Structure of Hydroxide Ion
From the electronic configurations given above, we can see that oxygen atom has 6 valence electrons; it needs 2 more electrons to fill up its outer electron shell.
In order to follow the octet rule, oxygen forms the hydroxide ion with the hydrogen atom by sharing one electron with one hydrogen atom and acquiring one more electron .
The gain of one electron results to the ion having a negative net charge of one.
The structure has one electron more than the total valence electrons of the ion (6 from O + 1 from H = 7 valence electrons).
Lewis Electron Dot Structure of Ammonium Ion
The nitrogen atom has 5 valence electrons as indicated by its electronic configuration; it needs three more electrons to fill up its last shell.
But the nitrogen atom in the ammonium ion satisfies the octet rule by sharing three of its valence electrons with each of the three hydrogen atoms.
The structure has one electron less than the total valence electrons of the ion (5 from N + 4 from 4 H = 9 valence electrons).
Determining The Net Charge of an Electron Dot Structure
From these examples, we can say that the:
- positive net charge of an ion is equal to the number of electrons deficient in its total valence electrons;
- negave net charge of an ion is equal to the number of electrons in excess of its total valence electrons.
The net charge of an ion is also determined by the algebraic sum of the formal charge of all the atoms of the ion.
The formal charge of an atom is the encircled + or - sign near the atom. It is calculated as follows:
formal charge of an atom = valence electrons of the atom - (number of unshared electrons of the atom + number of pairs of shared electrons by the atom)
Formal charge of hydroxide ion's atoms:
- O = 6 - (6 + 1) = -1
- H = 1 - (0 + 1) = 0
- net charge of hydroxide ion = -1 + 0 = -1
Formal charge of ammonium ion's atoms:
- N = 5 - (0 + 4) = +1
- H = 1 - (0 + 1) = 0
- H = 1 - (0 + 1) = 0
- H = 1 - (0 + 1) = 0
- H = 1 - (0 + 1) = 0
- net charge of ammonium ion = +1 + (4 x 0) = +1
Steps To Be Followed In Drawing A Lewis Electron Dot Structure
To summarize, the following steps may be followed to draw the Lewis electron dot structure of a chemical species.
- 1. Draw all possible structures by using single, double or triple pairs of electrons for atom-to-atom bonds.
- 2. Make sure that the structures contain the total valence electrons of all the atoms of the species.
- 3. If the species has a net charge, subtract electrons from the structure as indicated by the + charge or add electrons to the structure as indicated by the - charge.
- 4. Compute the formal charge of the atoms.
- 5. The structure that satisfies the octet rule for all atoms must be the correct Lewis structure.
Following the above steps, we're going to draw the electron dot structures of oxygen molecule (O2) and cyanide ion (CN-)
Drawing the Lewis Structure of a Molecule
Here are the 3 possible Lewis structures of N2 molecule.
Structures I and II are incorrect because they violate the octet rule:
- structure I has nitrogen atoms each containing only 6 electrons;
- structure II has nitrogen atoms each containing only 5 electrons.
Structure III is the correct Lewis electron dot structure of N2 molecule.
Drawing the Lewis Structure of an Ion
For the structures of cyanide ion as illustrated below, we added one electron to the total valence electrons of 9 (4 from C + 5 from N = 9) since it has a net charge of -1.
We rule out structures I and II because their atoms are electron deficient.
Structure III is the correct Lewis electron dot structure of cyanide ion.
The Resonance Theory
Sometimes, the step by step procedure we used above does not suffice because there are chemical species that have more than one possible Lewis electron dot structures.
One such species is the benzene molecule.
The Lewis structures shown below are equivalent but neither of them represents the actual structure of benzene molecule.
(We used single and double bonds here for easier viewing of the illustrations.)
Instead, benzene is represented by this structure:
This representation of a chemical species by a structure which is intermediate of 2 or more equivalent Lewis structures is called resonance.
The above structure is called resonance hybrid and structures I and II above are called resonance structures.
It should be noted here that resonance structures retain the same atomic arrangement; the only difference among these resonance structures is the arrangement of their electrons.
The Resonance Rules
There are several resonance rules that are used to determine the stability of resonance structures.
The most stable structures are the most important ones.
According to the resonance rules, the most stable structure is the one which has:
- 1. the greatest number of covalent bonds;
- 2. the least number of formal charges;
- 3. the - sign on the more electronegative atom and the + sign on the more electropositive atom, if formal charges are present.
We will use these rules to help us determine the most important Lewis structures of compounds having several resonance structures.
Determining The Most Stable Resonance Structure
Given below are the possible electron dot structures of CO2.
We strike off structures I, II and III because of their electron deficient atoms.
Our choices then boil down to structures IV - VI.
We can see that structures V and VI are equivalent; we can use any of the two to compare with structure IV.
Being equivalent in the number of covalent bonds, we compare their number of formal charge.
Structure IV is more stable than structure V or VI because it has no formal charge. Hence, structure IV is the Lewis electron dot structure of CO2.
Let's have another example: an ion this time.
The possible structures of thiocyanate ion (SCN-) are given below.
How did we determine that C atom is the central atom?
We take the atom with the lowest valence electrons (except H) as the central atom because it has the maximum number of unpaired electrons that can be used to form covalent bonds with other atoms.
We exclude structures I - III from our choices because they don't follow the octet rule.
Since the above 3 structures have all equal number of covalent bonds, we will evaluate their stability based on the number of formal charge.
Structures IV and V are equivalent and are more stable than structure VI because they have less number of formal charge.
We choose structure IV as the Lewis structure of thiocyanate ion because the - sign is on the more electronegative atom.
Conclusion
By using the procedure and rules given above, you are now ready to draw the Lewis electron dot formula of most compounds having covalent bonds. Try the problems below.
Problems
The best way to improve your problem solving skill in chemistry is to practice solving as many chemistry problems as possible. Here are some common ions and molecules. Draw their Lewis electron dot structures.
For more problems on other chemistry topics, go to www.tutorpartner.blogspot.com.
| 1. | C2H5NH3+ |
| 2. | CH3COO- |
| 3. | H2O2 |
| 4. | HSO3- |
| 5. | HCO3- |
| 6. | CH3CN |
| 7. | C2O4-2 |
| 8. | HNO2 |
| 9. | HCOCl |
| 10. | N2 |
| 11. | CH3CONH2 |
| 12. | SO3-2 |
| 13. | C4H6 |
| 14. | NO3- |
| 15. | NO2- |
