Lewis Structures

Now that we know a bit about how electrons are used in bonds, we will begin to draw up lewis structures, which are a way of illustrating the bonds between atoms in a molecule by drawing every valence electron.

Here is an example of a lewis structure of water:

H:O:H

Here, oxygen has six valence electrons. In water, two participate in two O–H bonding pairs, each with an electron from a hydrogen atom. The other four oxygen valence electrons remain as two lone pairs. Count all eight valence electrons in the H₂O molecule: six from O and one from each H. A line-bond shorthand may omit dots, but it does not make the lone pairs disappear.

If we look at the periodic table, we can know how many valence electrons very easily for non-metals, alkali metals, and alkali earth metals by looking at the column it is in. The first column (the alkali metals) each have one valence electron. The second (alkali earth metals) have two. From there we skip the central section of the periodic table and jump to column 13, which starts with Boron. That column has 3 valence electrons, and then from there on, the number of valence electrons simply increases by one. Finally, the noble gases have 8 valence electrons.

Even though noble gases have 8 electrons, they typically don’t form compounds very easily. This is due to something called the octet rule, which says that atoms tend to form bonds in order to attain 8 valence electrons. This can be seen on the oxygen molecule in the H2O example. The oxygen molecule has 8 electrons in its perimeter.

Well, once you have the valence electron number, you begin to add dots around the atom’s label, on each side of the atomic label (i.e. top, bottom, left, right). You only put one on each side until you run out of electrons, and if you have more, then you can begin pairing them. We will go ahead now and show you a gallery of all the lewis structures for the elements Li through Neon, to show you what we mean:

Once you have the structures drawn, it is time to begin bonding atoms. Note that you sometimes have a side with only one electron and other times you have sides with two electrons. These represent electron pairs. For example, Oxygen has two electron pairs and two unpaired electrons. Typically, the only bond-forming electrons are the unpaired electrons. Therefore, Oxygen can form 2 bonds.

Unpaired electrons can bond with unpaired electrons from other atoms. The following gallery will guide you through drawing the lewis structure of water (H2O).

One technique for drawing Lewis Structures that is very useful for the upcoming special circumstances is outlined in the following example problem. It entails finding the total number of valence electrons, then making the initial sets of bonds, and finally adding any leftover electrons.

The other major aspect of putting together Lewis Structures is the octet rule. This applies to most of the atoms in the first couple rows of the periodic table, with some exceptions (H, S, P for example). The rule states that an atom will bond such that it forms a structure that provides it with 8 electrons. These include both electrons involved in a bond, at least in the case of covalent bonds (which the bulk of this section deals with).

Formal charge is bookkeeping, not a direct measurement of an atom’s physical charge. Assign each atom all of its lone-pair electrons and half the electrons in each bond; compare that assigned count with the neutral atom’s valence-electron count. Formal charge helps us compare possible Lewis structures, while bond polarity concerns unequal electron density and oxidation state uses a different assignment convention.

Example Problem 4.2.1

Draw the Lewis Structure for Carbon Tetrachloride (CCl4)

You should now be a whiz at putting together basic Lewis Structures.

A few more rules now, and then you’ll be ready to draw any compound’s structure.

If the initial single-bond skeleton leaves a central atom short of an octet, consider converting a neighboring lone pair into an additional shared electron pair. That creates a double or triple bond without adding atoms or electrons. Recount electrons and formal charges after the change, and use the specified charge and known octet exceptions to assess the structure.

Example Problem 4.2.2

Draw the Lewis Structure for the Carbon Monoxide (CO) molecule.

An ion will be similar to any other compound, except that it will have extra electrons (if it is negative) or less electrons (if it is positive).

Note that a positively charge ion is called a cation and a negatively charged ion is an anion.

Example Problem 4.2.3

Draw the Lewis Structure for the Ammonium (NH4+) ion.

We are nearly done talking about the Lewis Structures of covalent bonded molecules. One last topic we have to cover is resonance.

Sometimes, we have multiple options that satisfy all the rules we have previously described. In that case, each possible structure is correct. This is called resonance. The compound will have a few different representations, and you have to illustrate them all if you want to fully represent the structure of the molecule.

Resonance contributors are alternative Lewis drawings for the same arrangement of nuclei. The actual species is a delocalized hybrid, not a molecule switching rapidly between the drawings. Equivalent contributors can contribute equally; non-equivalent contributors do not necessarily have equal weights. Use measured bond lengths and a consistent electron count to explain why a single localized drawing is incomplete.

An example of resonance is the SO2 molecule, whose Lewis Structure(s) looks like this:

<Image 1>

Notice that the structures are each bracketed and linked by the double-headed arrow. This is what resonance entails. The arrow represents that the true form exists somewhere in between those given forms.

On the next page, we will go through an example of how to find the resonance structure of a complex ion: the acetate ion (CH3COO-).

Example Problem 4.2.4

Draw the Lewis Structure for the Acetate (CH3COO-) ion.

That just about does it for Lewis Structures, but you might also need to draw Lewis Structures for ionic bonds. We don’t need an example problem for this, so instead you’ll just get the quick version. As there is no sharing of electrons in an ionic bond, the Lewis Structure of ions involved in an ionic bond must show the distinction.

Also, when drawing the Lewis Structure, remember to include the formal charges of the involved ions.

A couple example structures are:

MgO: <Image 1>

CaCl2: <Image 2>

BaBr2: <Image 3>

Al2O3: <Image 4>