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Learning Objectives
- Draw the Lewis dot structure of a given molecule or ion.
- Draw resonance structures of some molecules.
- Assign formal charge to an atom in a dot structure.
- Assess the stability of a structure by considering formal charges of atoms.
- Give examples for molecules and ions that do not follow the octet rule.
Lewis Dot Structures
| Lewis symbols of the main group elements | |||||||
| \(\ce{H\cdot}\) | \(\textrm{He:}\) | ||||||
|---|---|---|---|---|---|---|---|
\(\underset{\:}{\ce{Li\cdot}}\) | \(\underset{\:}{\ce{\cdot Be \cdot}}\) | \(\ce{ \cdot \underset{\:}{\overset{\Large{\cdot}}{B}} \cdot}\) | \(\ce{ \cdot \underset{\Large{\cdot}}{\overset{\Large{\cdot}}{C}} \cdot}\) | \( \underset{\Large{\cdot\,}} {\overset{\Large{\cdot}} {\textrm{:N}\cdot}} \) | \( \underset{\Large{\cdot\cdot\,}} {\overset{\Large{\cdot}} {\textrm{:O}\cdot}} \) | \( \underset{\Large{\cdot\cdot}} {\overset{\Large{\cdot\cdot}} {\textrm{:F}\cdot}} \) | \( \underset{\Large{\cdot\cdot}} {\overset{\Large{\cdot\cdot}}{\textrm{:Ne:}}} \) |
| \(\ce{Na}\) \(\ce{K}\) \(\ce{Rb}\) \(\ce{Cs}\) | \(\ce{Mg}\) \(\ce{Ca}\) \(\ce{Sr}\) \(\ce{Ba}\) | \(\ce{Al}\) \(\ce{Ga}\) \(\ce{In}\) \(\ce{Tl}\) | \(\ce{Si}\) \(\ce{Ge}\) \(\ce{Sn}\) \(\ce{Pb}\) | \(\ce{P}\) \(\ce{As}\) \(\ce{Sb}\) \(\ce{Bi}\) | \(\ce{S}\) \(\ce{Se}\) \(\ce{Te}\) \(\ce{Po}\) | \(\ce{Cl}\) \(\ce{Br}\) \(\ce{I}\) \(\ce{At}\) | \(\ce{Ar}\) \(\ce{Kr}\) \(\ce{At}\) \(\ce{Rn}\) |
G.N. Lewis used dots to represent the valence electrons in his teaching of chemical bonding. He eventually published his theory of chemical bonding in 1916. He put dots around the symbols so that we see the valence electrons for the main group elements. Formation of chemical bonds to complete the requirement of eight electrons for the atom becomes a natural tendency. Lewis dot symbols of the first two periods are given here to illustrate this point. In fact, the entire group (column) of elements have the same Lewis dot symbols, because they have the same number of valence electrons.
| \(\ce{CF4}\), \(\ce{H2O}\) and \(\ce{CO2}\) dot structures | ||
|---|---|---|
| .. : F : . . . . . . : F : C : F : . . . . . . : F : ' ' | .. .. O / \ H H | .. .. O::C::O .. .. or . . . . |
Lewis dot structures are useful in explaining the chemical bonding in molecules or ions. When several dot structures are reasonable for a molecule or ion, they all contribute to the molecular or ionic structure making it more stable.
The representation of a molecular or ionic structure by several structures is called resonance. The more stable the dot structure is, the more it contributes to the electronic structure of the molecule or ion.
You need to know what dot structures represent, how to draw them, and what the formal charges for the atoms in the structure are. When several dot structures are possible, consider the resonance structures to interpret the real structure. Apply some simple rules to explain which of the resonance structures are major contributors to the electronic structure.
Drawing Lewis Dot Structures and Resonance Structures
Follow these simple steps to draw Lewis dot structures:
- Draw the atoms on paper and put dots around them to represent valence electrons of the atom. Be sure to have the correct number of electrons.
- If the species is an ion, add or subtract electrons corresponding to the charge of the ion. Add an electron for every negative (-) charge, and subtract an electrons for every positive (+) charge.
- Consider bonding between atoms by sharing electrons; some may come from one atom.
- If possible, apply the octet rule to your structure. Some structures don't obey the octet rule, but explain why.
- Assign formal charges to atoms in the structure.
Exercise
Draw Lewis dot structures for \(\ce{CH4}\), \(\ce{NH3}\), \(\ce{HF}\), \(\ce{OF2}\), \(\ce{F2}\), \(\ce{O2}\), \(\ce{N2}\), \(\ce{Cl-}\) and some compounds you know.
Formal Charge
The formal charge on any atom in a Lewis structure is a number assigned to it according to the number of valence electrons of the atom and the number of electrons around it. The formal charge of an atom is equal to the number of valence electrons, Nv.e. minus the number of unshared electrons, Nus.e. and half of the bonding electrons, ½ Nb.e..
\(Formal\: charge = N_{\large{v.e.}} - N_{\large{us.e.}} - \dfrac{1}{2} N_{\large{b.e.}}\)
Some practice of assigning formal charge is necessary before you master this technique. Some examples of drawing Lewis structure and assigning formal charge are given below.
The formal charge is a hypothetical charge from the dot structure. The formal charges in a structure tell us the quality of the dot structure.
Formal charge rules
Often, many Lewis dot structures are possible. These are possible resonance structures, but often we should write a reasonable one, which is stable. The formal charge guides us about the stability of the dot structure. The guidance is called formal charge rules:
- Formulas with the lowest magnitude of formal charges are more stable.
- More electonegative atoms should have negative formal charges.
- Adjacent atoms should have opposite formal charges.
Example 1
Draw Lewis dot structure for \(\ce{SO2}\).
Solution
Put down number of valence electrons:
\(\mathrm{
:\overset{\Large{..}}O :
:\overset{\Large{..}}S :
:\overset{\Large{..}}O :}\)
Put all atoms together to make a molecule and check to see if it satisfies the octet rule.
\(\begin{alignat}{1}
:&\overset{\Large{..}}{\ce O} :
:&&\overset{\Large{..}}{\ce S} :
:&&\overset{\Large{..}}{\ce O} : &&\textrm{ <= octet rule not satisfied}\\
&\,0 &&\,0 &&\,0 &&\textrm{ formal charge}
\end{alignat}\)
Adjust bonding electrons so that octet rules apply to all the atoms.
\(\begin{alignat}{1}
&:\underset{\Large{..}}{\overset{\Large{..}}{\ce O}}
&&:\overset{\Large{..}}{\ce S} :
:&&\overset{\Large{..}}{\ce O} : &&\textrm{ <- octet rule satisfied}\\
&\,{-1} &&\,{+1} &&\,0 &&\textrm{ formal charge}
\end{alignat}\)
Since the left \(\ce{O}\) has 6 unshared plus 2 shared electrons, it effectively has 7 electrons for a 6-valence-electron \(\ce{O}\), and thus its formal charge is -1.
Formal charge for \(\ce{O}\) = 6 - 6 - (2/2) = -1.
Formal charge for \(\ce{S}\) = 6 - 2 - (6/2) = +1.
There is yet another structure that does not satisfy the octet rule, but it's a reasonable structure:
\(\begin{alignat}{1}
&:\underset{\Large{..}}{\overset{\Large{..}}{\ce O}}
&&:\overset{\Large{..}}{\ce S} :
&&\underset{\Large{..}}{\overset{\Large{..}}{\ce O}} : &&\textrm{ <- octet rule not satisfied}\\
&\,{-1} &&\,{+2} &&{-1} &&\textrm{ formal charge}
\end{alignat}\)
Resonance Structures
When several structures with different electron distributions among the bonds are possible, all structures contribute to the electronic structure of the molecule. These structures are called resonance structures. A combination of all these resonance structures represents the real or observed structure. The Lewis structures of some molecules do not agree with the observed structures. For such a molecule, several dot structures may be drawn. All the dot structures contribute to the real structure. The more stable structures contribute more than less stable ones.
For resonance structures, the skeleton of the molecule (or ion) stays in the same relative position, and only distributions of electrons in the resonance structures are different.
Let us return to the \(\ce{SO2}\) molecule. The molecule has a bent structure due to the lone pair of electrons on \(\ce{S}\). In the last structure that has a formal charge, there is a single \(\ce{S-O}\) bond and a double \(\ce{S=O}\) bond. These two bonds can switch over giving two resonance structures as shown below.
| 1 | 2 | 3 | 4 | |||
|---|---|---|---|---|---|---|
| .. S / \ :O: :O: ' ' ' ' | « | .. S // \ :O: :O: ' ' | « | .. S / \\ :O: :O: ' ' | « | .. S // \\ :O: :O: |
In structure 1, the formal charges are +2 for \(\ce{S}\), and -1 for both \(\ce{O}\) atoms. In structures 2 and 3, the formal charges are +1 for \(\ce{S}\), and -1 for the oxygen atom with a single bond to \(\ce{S}\). The low formal charges of \(\ce{S}\) make structures 2 and 3 more stable or more important contributors. The formal charges for all atoms are zero for structure 4, given earlier. This is also a possible resonance structure, although the octet rule is not satisfied. Combining resonance structures 2 and 3 results in the following structure:
| .. S /.' '' '.\ :O: :O: |
|---|
Exercise
Draw the Lewis dot structures and resonance structures for the following. Some hints are given.
\(\ce{CO2}\) - \(\textrm{:O::C::O:}\) (plus two more dots for each of \(\ce{O}\))
\(\ce{NO2}\) - \(\ce{.NO2}\) (bent molecule due to the odd electron)
\(\ce{NO2-}\) - \(\ce{:NO2-}\) (same number of electron as \(\ce{SO2}\))
\(\ce{HCO2-}\) - \(\ce{H-CO2}\)
\(\ce{O3}\) - (ozone, \(\ce{OO2}\); same number of electrons as \(\ce{SO2}\))
\(\ce{SO3}\) - (consider \(\ce{O-SO2}\), and the resonance structures)
\(\ce{NO3-}\) (see Example 2 below)
\(\ce{CO3^2-}\) (ditto)
Notice that some of the resonance structures may not satisfy the octet rule. The \(\ce{NO2}\) molecule has an odd number of electrons, and the octet rule cannot be satisfied for the nitrogen atom.
Example 2
Draw the resonance structures of \(\ce{NO3-}\)
Solution
| -1 :O: || N / \ :O: :O: ' ' ' ' | « | . . -1 :O: | N // \ :O: :O: ' ' | « | . . -1 :O: | N / \\ :O: :O: ' ' |
|---|
The resonance structure is shown on the right here. Note that only the locations of double and single bonds change here. What are the formal charges for the \(\ce{N}\) atoms? What are the formal charges for the oxygen atoms that are single bonded and double bonded to \(\ce{N}\) respectively? Please work these numbers out.
- Formal charges: \(\ce{N}\), +1; \(\ce{=O}\), 0; \(\ce{-O}\), -1
- The most stable structure has the least formal charge.
- In a stable structure, adjacent atoms should have formal charges of opposite signs.
The more stable the structure, the more it contributes to the resonance structure of the molecule or ion. All three structures above are the same, only the double bond rotates.
Exercise
Draw the Lewis dot structures and resonance structures for
\(\ce{HNO3}\)
\(\ce{H2SO4}\)
\(\ce{H2CO3}\)
\(\ce{HClO4}\)
\(\ce{C5H5N}\)
\(\ce{NO3-}\)
\(\ce{SO4^2-}\)
\(\ce{CO3^2-}\)
\(\ce{ClO4-}\)
\(\ce{C6H6}\) (Benzene)
\(\ce{Cl2CO}\)
You have to do these on paper, because putting dots around the symbols is very difficult using a word processor. The octet rule should be applied to \(\ce{HNO3}\), \(\ce{NO3-}\), \(\ce{H2CO3}\), \(\ce{CO3^2-}\), \(\ce{C5H5N}\), \(\ce{C6H6}\), and \(\ce{Cl2CO}\).
Exceptions to the octet rule
We can write Lewis dot structures that satisfy the octet rule for many molecules consisting of main-group elements, but the octet rule may not be satisfied for a number of compounds. For example, the dot structures for \(\ce{NO}\), \(\ce{NO2}\), \(\ce{BF3}\) (\(\ce{AlCl3}\)), and \(\ce{BeCl2}\) do not satisfy the octet rule.
| .N:::O: compare :C:::O: | . N // \ :O: :O: ' ' | .. :F: | B / \ :F: :F: ' ' ' ' | . . . . :Cl : Be : Cl: ' ' ' ' |
|---|
The above are structures for the gas molecules. The solids of \(\ce{AlCl3}\) and \(\ce{BeCl2}\) are polymeric with bridged chlorides.
| . . . . :Cl: :Cl: :Cl: \ / \ / Al Al / \ / \ :Cl: :Cl: :Cl: . . . . | Polymeric | :Cl: :Cl: :Cl: |
|---|
Alumunum chloride, \(\ce{AlCl3}\), is a white, crystalline solid, and an ionic compound. However, it has a low melting point of 465 K (192°C), and the liquid consists of dimers, \(\ce{Al2Cl6}\), whose structure is shown above. It vaporizes as dimers, but further heating gives a monomer that has the same structure as the \(\ce{BF3}\).
| Arrange dots this way | ||
|---|---|---|
| \(\begin{alignat}{5} \textrm{::Ex} &= \mathrm{\overset{\:\Large{..}}{:Ex}}\\ \textrm{:::Ex} &= \mathrm{\overset{\Large{..}}{:Ex:}} \end{alignat}\) |
In compounds \(\ce{PF5}\), \(\ce{PCl5}\), \(\textrm{:SF}_4\), \(\textrm{::ClF}_3\), \(\textrm{:::XeF}_2\) and \(\textrm{:::I}_3^-\), the center atoms have more than 10 electrons instead of 8. In compounds \(\ce{SF6}\), \(\ce{IOF5}\), \(\textrm{:IF}_5\), \(\ce{BrF5}\), \(\textrm{::XeF}_4\), \(\ce{PF6-}\) etc, the center atoms have 12 electrons.
The formulas given above follow a systematic pattern according to the positions of the elements on the periodic table. As the number of atoms bonded to it decreases, the number of unshared electrons increase.
Confidence Building Questions
- What is the total number of valence electrons in \(\ce{CO2}\)?
What about \(\ce{NO2}\)?
- What is the formal charge for \(\ce{S}\) in \(\ce{H2SO4}\) in the following structure?
O-H | (provide all dots yourself) O=S=O | O-H
The oxygen double bonded to \(\ce{S}\) has a formal charge of 0. What is the formal charge of \(\ce{C}\) in \(\ce{O-C}\textrm{:::O}\)?
- What is the formal charge for \(\ce{S}\) in \(\ce{H2SO4}\) in the following structure?
O-H | (provide all dots yourself) O-S-O | O-H
The oxygen double bonded to \(\ce{S}\) has formal charge of -1. What is the formal charge of \(\ce{C}\) in \(\ce{O-C}\textrm{:::O}\)?
- According to the formal charge rule, which structure in the two problems you have just worked on is the best for \(\ce{H2SO4}\)?
- What is the formal charge for \(\ce{N}\) in the following structure?
. N = O |:O: ..
What is the formal charge if this is the structure:.:O: | .. N = O : ..
and which one you think is the best? You can write two resonance structures for each of the two to give 4 resonance structures for \(\ce{NO2}\).
- What is the formal charge for \(\ce{B}\) in the following structure?
F \ B = F (Supply your dots for unshared electrons) /F
What about \(\ce{B(-F)3}\), all single bonds?
- What is the formal charge of \(\ce{I}\) in \(\ce{I(-Cl)3}\)? (Supply your dots)
What about \(\ce{Cl}\) in the same structure?
- Which one of the following compounds has the same number of valence electrons as \(\ce{NO2-}\):
\(\ce{CO2}\), \(\ce{NO2}\), \(\ce{O3}\), \(\ce{CO3-}\), or \(\ce{CO2-}\)? Elements of the 2nd period are: \(\ce{Li}\) \(\ce{Be}\) \(\ce{B}\) \(\ce{C}\) \(\ce{N}\) \(\ce{O}\) \(\ce{F}\) \(\ce{Ne}\).
Draw Lewis dot structure for \(\ce{O3}\) and \(\ce{NO2-}\).
Contributors and Attributions
Chung (Peter) Chieh (Professor Emeritus, Chemistry @University of Waterloo)
