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Lewis Dot Structures

What are lewis structures?

• Lewis structures are used to describe and visualize molecules.

• Lewis structures are used to show the bonds between atoms as well as the electrons surrounding certain atoms.

• Lewis structures can help determine the formal charge of a molecule.

Steps to drawing lewis dot structures:

1. Determine which atoms are connected to each other.

2. Determine the number of valence electrons in the molecule.

3. Place two electrons between each atom in place of the bonds.

4. Add the rest of the available valence electrons to complete the octet of the surrounding atoms.

5. Determine if the central atom has a complete octet, if it does not, add double or triple bonds to complete its octet.

Walkthrough of drawing lewis dot structures

Determine the lewis dot structure for $SO_{3}$$^2$$^-$ (meaning there are two extra valence electrons in this molecule).

1) Determine how the atoms connect to each other.

• Because there's one sulfur and three oxygen atoms, it is safe to assume that the sulfur is the central atom, surrounded by the three oxygens.

2) Determine the number of valence electrons in the molecule.

Sulfur: 1 x 6 = 6 valence electrons.

Oxygen: 3 x 6 = 18 valence electrons.

Total valence electrons: 6 + 18 + 2 extra = 26 valence electrons.

$\color{green}{Note:}$ The periodic table shows you how many valence electron each element has.

3) Place two electrons between each atom in place of bonds.

4) Add the rest of the available valence electrons to complete the octet of the surrounding atoms.

26 - 6 = 20 valence electrons left over.

• These leftover electrons get placed on the oxygen atoms.

5) Determine if the central atom has a complete octet, if it does not, add double or triple bonds to complete its octet.

• In this case, Sulfur has only 6 electrons, so it needs 2 more electrons to complete its octet, thus we will form a double bond between one of the oxygens to complete its octet! This is done by shifting one of the electron pairs from the oxygen atom to form another bond with sulfur.

Now we can draw the actual bonds instead of the dots and we are finished!

How to determine the formal charge using the lewis dot diagram

• To determine the formal charge, simply just use the following equation to calculate the formal charge!

Let's find the formal charge for the molecule from the question above ($SO_{3}$)

Formal charge of Sulfur = $6-2-\frac{8}{2}$ = 0

Formal charge of oxygen = $6-6-\frac{2}{2}$ = -1

Formal charge of oxygen = $6-6-\frac{2}{2}$ = -1

Formal charge of double bond oxygen = $6-4-\frac{4}{2}$ = 0

Thus the total formal charge is: -1 + -1 = -2

(Video) Lewis Diagrams Made Easy: How to Draw Lewis Dot Structures

by ketzbook

This video gives a great explanation on how to draw Lewis Dot Diagrams. You'll also learn five simple steps to keep in mind while drawing your own diagrams.

Sometimes, we need to visualize a molecule. To help our visualization, we draw a Lewis dot structure. When we draw lewis dot structures, we are looking at covalently bonded molecules. The atoms are sharing pairs of electrons to fill their octets. Moving forward, the Lewis structure will help us determine the VSEPR geometry, ideal bond angles and angle deviations.

Steps for Drawing a Lewis Structure

1. Determine which atom is the central atom. Use the molecular formula to determine the central atom. It is always the first atom listed. Often there will only be one atom.
2. Determine the total number of valence electrons in the molecule. You can find the number of valence electrons by looking at each atom's atomic number on the periodic table. Add each atom's valence electrons to find the molecule's valence electrons.
3. Determine the number of lone pairs on the central atom. We use the formula $LP$=$\frac{VE-8n}{2}$. LP is the number of lone pairs. VE is the total valence electrons from step two. And n is the total atoms bonded to the central atom. The exception to the formula is hydrogen. Instead of 8n, the term is 2n. It is 2n because hydrogen has a full octet with two electrons.
4. We can now begin to draw our lewis structure. Write the atomic symbol for the central atom.
5. Write the atomic symbols for the other elements on each side of the central atom. Leave a half-inch gap between the central atom and the surrounding elements. If you have three surrounding elements, you leave one side blank. If you have two elements, draw them to the right and left of the central atom.
6. Place any lone pairs on the blank sides of the central atom. To draw a lone pair, write two dots directly next to each other. You do not need to leave a gap between the central atoms and the lone pairs.
7. Draw a line between the central atom and each surrounding atom. The line represents a single covalent bond.
8. Determine how many valence electrons you have remaining. Subtract two electrons for each lone pair and each single bond.
9. Place the remaining electrons around the non-central atoms. Draw lone pair dots on each non-bonded side of the atom.
10. Make sure that every atom has a complete octet. When an atom does not have an octet, convert a lone pair into another bond.
11. Minimize the total number of formal charges. You can turn lone pairs into another bond to reduce the number of formal charges. Remember that the total formal charge is the charge of the molecule.

An Easy Example: Chloroform

Draw the lewis dot structure for chloroform, $CHCl_{3}$.

1. In chloroform's molecular formula, the carbon is first, and there is only one atom of carbon. So, carbon is the central atom.
2. There are twenty-six valence electrons. Carbon has four. Hydrogen has one. Chlorine has seven. But, there are three chlorine atoms. The three chlorines contribute twenty-one electrons.
3. $LP$=$\frac{26-8(3)-2(1)}{2}$=0. The carbon will have zero lone pairs. Remember, hydrogen is the exception. For each hydrogen, we subtract two electrons.
4. Carbon is the central atom. Write a capital C.
5. In chloroform, chlorine is on three sides of the carbon. Hydrogen is on the remaining side.
6. There are zero lone pairs.
7. You should have a line between the carbon and each of the other elements.
8. We have used eight electrons. We have eighteen electrons remaining.
9. The eighteen remaining electrons will be lone pairs around the chlorine atoms.
10. Every atom has a complete octet. Carbon's four single bonds give it eight electrons. Each chlorine has eight from the one single bond and three sets of lone pairs. The hydrogen has two electrons from the single bond.
11. None of the atoms have a formal charge.

A Harder Example: Sulfur Dioxide

Try drawing the lewis structure for sulfur dioxide, $SO_{2}$. I have written the steps below to help you if you get stuck.

1. The sulfur is the central atom.
2. There are eighteen valence electrons.
3. $LP$=$\frac{18-8(2)}{2}$=1. There will be one lone pair on the sulfur.
4. Write a capital S.
5. Draw oxygen to the left and right of the sulfur.
6. Place the lone pair dots above the sulfur.
7. Draw a line between sulfur and each oxygen.
8. We have used six electrons. Two for the lone pair. Two for every single bond. We have twelve electrons remaining.
9. The twelve remaining electrons will be lone pair dots around the oxygens.
10. Each oxygen has a complete octet. The sulfur does not have an octet. Change one of oxygen's lone pairs to a bond between sulfur and oxygen. One oxygen should have a single bond to sulfur and three lone pairs. The other oxygen should have a double bond to sulfur and two lone pairs.
11. The sulfur has a positive formal charge. The single-bonded oxygen has a negative formal charge. To minimize the formal charge, we convert one of the lone pairs on the single-bonded oxygen to a double bond. Now, the formal charge on every atom is zero.

The lewis structure of sulfur dioxide has sulfur double bonded to both oxygens. The sulfur also has a lone pair of electrons.

Wait, do I obey the octet rule or minimize the formal charges?

In sulfur dioxide, sulfur is an exception to the octet rule. Some of the elements in period three of the periodic table can be exceptions to the octet rule. We know that sulfur, phosphorous, and halogens can have more than eight electrons. For the exceptions, minimizing the formal charge is more important. For the elements in period two, you must obey the octet rule.

Drawing the Exceptions to the Octet Rule

There are four exceptions to the octet rule. Here are some tips for the octet rule exceptions.

1. The transition metals and the lanthanides and actinides can use their d and f orbitals to bond. We do not draw lewis dot structures for them.
2. Molecules with an odd number of electrons have a single electron dot instead of a lone pair. A single electron can go on any of the atoms. Because a single electron can go on any atom, the molecule has resonance structures.
3. Some molecules can have more than eight valence electrons. For hypervalent molecules, we use the same steps. We will often need to minimize the formal charge for the hypervalent molecules. See the sulfur dioxide example. We should memorize which elements can have more than eight electrons.
4. For hypovalent molecules, the central atom does not have to obey the octet rule. An example is borane ($BH_{3}$). Boron only has three valence electrons. So, it only forms three bonds.

A Final Example: Lewis Structures of Ions

Try drawing the lewis structure for the oxoanion sulfite, $S$$O_{3}^{2-}$.

1. Sulfur is the central atom.
2. There are twenty-six valence electrons. Six from sulfur. Eighteen from the three oxygens. We add two more electrons because we have a negative two charge.
3. $LP$=$\frac{26-8(3)}{2}$=1
4. Write a capital S.
5. Draw oxygen to the right, left, and above the sulfur.
6. Place the lone pair below the sulfur.
7. Draw a line between sulfur and each oxygen.
8. We have used eight electrons. Six for the single bonds. Two for the lone pair. We have eighteen remaining.
9. The remaining electrons will be lone pairs on the oxygen atoms.
10. Every atom has a complete octet.
11. Each oxygen has a negative one formal charge. The sulfur has a positive one formal charge. To minimize the formal charges, turn one of oxygen's lone pairs into a double bond. Now both single-bonded oxygens have a negative formal charge. The double-bonded oxygen and sulfur do not have a formal charge.

The lewis structure of sulfite is sulfur with one double bond, two single bonds, and one lone pair.