Chemical Bonding and Molecular Structure

Here are some rules to follow when drawing Lewis structures—you should follow these simple steps for every Lewis structure you draw, and soon enough you’ll find that you’ve memorized them. While you will not specifically be asked to draw Lewis structures on the test, you will be asked to predict molecular shapes, and in order to do this you need to be able to draw the Lewis structure—so memorize these rules! To predict arrangement of atoms within the molecule

1. Find the total number of valence electrons by adding up group numbers of the elements. For anions, add the appropriate number of electrons, and for cations, subtract the appropriate number of electrons. Divide by 2 to get the number of electron pairs.
2. Determine which is the central atom—in situations where the central atom has a group of other atoms bonded to it, the central atom is usually written first. For example, in CCl4, the carbon atom is the central atom. You should also note that the central atom is usually less electronegative than the ones that surround it, so you can use this fact to determine which is the central atom in cases that seem more ambiguous.
3. Place one pair of electrons between each pair of bonded atoms and subtract the number of electrons used for each bond (2) from your total.
4. Place lone pairs about each terminal atom (except H, which can only have two electrons) to satisfy the octet rule. Leftover pairs should be assigned to the central atom. If the central atom is from the third or higher period, it can accommodate more than four electron pairs since it has d orbitals in which to place them.
5. If the central atom is not yet surrounded by four electron pairs, convert one or more terminal atom lone pairs to double bonds. Remember that not all elements form double bonds: only C, N, O, P, and S!

Example

Which one of the following molecules contains a triple bond: PF₃, NF₃, C₂H₂, H₂CO, or HOF?

Explanations

The answer is C₂H₂, which is also known as ethyne. When drawing this structure, remember the rules. Find the total number of valence electrons in the molecule by adding the group numbers of its constituent atoms. So for C₂H₂, this would mean C = 4 × 2 (since there are two carbons) = 8. Add to this the group number of H, which is 1, times 2 because there are two hydrogens = a total of 10 valence electrons. Next, the carbons are clearly acting as the central atoms since hydrogen can only have two electrons and thus can’t form more than one bond. So your molecule looks like this: H—C—C—H. So far you’ve used up six electrons in three bonds. Hydrogen can’t support any more electrons, though: both H’s have their maximum number! So your first thought might be to add the remaining electrons to the central carbons—but there is no way of spreading out the remaining four electrons to satisfy the octets of both carbon atoms except to draw a triple bond between the two carbons. For practice, try drawing the structures of the other four compounds listed.

Example

How many sigma (s) bonds and how many pi (p) bonds does the molecule ethene, C₂H₄, contain?

Explanations

First draw the Lewis structure for this compound, and you’ll see that it contains one double bond (between the two carbons) and four single bonds. Each single bond is a sigma bond, and the double bond is made up of one sigma bond and one pi bond, so there are five sigma bonds and one pi bond.

Sometimes you’ll come across a structure that can’t be determined by following the Lewis dot structure rules. For example, ozone (O₃) contains two bonds of equal bond length, which seems to indicate that there are an equal number of bonding pairs on each side of the central O atom. But try drawing the Lewis structure for ozone, and this is what you get:

We have drawn the molecule with one double bond and one single bond, but since we know that the bond lengths in the molecule are equal, ozone can’t have one double and one single bond—the double bond would be much shorter than the single one. Think about it again, though—we could also draw the structure as below, with the double bond on the other side:

Together, our two drawings of ozone are resonance structures for the molecule. Resonance structures are two or more Lewis structures that describe a molecule: their composite represents a true structure for the molecule. We use the double-directional arrows to indicate resonance and also bracket the structures or simply draw a single, composite picture.

Let’s look at another example of resonance, in the carbonate ion CO₃^2-:

Notice that resonance structures differ only in electron pair positions, not atom positions!

Example

Draw the Lewis structures for the following molecules: HF, N₂, NH₃, CH₄, CF₄, and NO⁺.

Explanation