Building Chemical Intuition by Building Lewis Structures:

Lewis Structures are much more than Step One of VSEPR.

Dave Doherty

Click to play the animation. For optimal viewing, select full-screen mode.

Sneak Peek: Keeping track of the electrons in NH3 + H2O ⇄ NH4+ + HO-

Experience First … Memorize Later

[Innovative, simple, automatic!, efficient, intuitive] [rules, recipes, algorithms, procedures, steps] for writing correct Lewis structures are a dime a dozen. In fact, you can find at least 851 articles about them in J. Chem. Educ.

A procedure for writing Lewis structures.

Figure 1. One example of the dizzying overabundance of variations of “rules” for writing Lewis structures.

And – make no mistake – students quickly decide that the value of these rules is ephemeral: soon to disappear from their minds after that one time they need them (on the test). How do we know that? In an average month, students conduct thousands of Google searches for the Lewis structure of just one of many molecules: NH3.

This habit cheats students out of the true value of the Lewis structure and its ancillary idea, the octet rule:

The Lewis structure is so much more than a series of steps to be memorized – it is a powerful, dynamic model that should be poked and prodded to investigate and absorb the gestalt of chemical bonding.

Sure, at some point (before the test) your students will need to know some version of the rules for writing Lewis structures. But first, we suggest that they spend time building and manipulating Lewis structure models to gain a feeling for what chemical bonding really means.

As students do so, they will build chemical intuition – a deeper comprehension of how electrons arrange in atoms, molecules, and ions.

Some will even be able to use this newly acquired “intuition” to devise their own versions of the “rules” (ready to be fed into the pipeline at J. Chem. Educ. 😉).

A Dynamic Lewis Structure Model

Lewis structures are a simple, symbolic representation of chemical bonding that can be used to explain covalent and ionic bonding via the octet rule. But don’t be fooled: ”simple“ does not imply “superficial”. As we will demonstrate below, the use of Lewis structure models can promote understanding of many core chemical concepts – some of which you may not have thought of.

To illustrate, we’ll show you examples using the Atomsmith Molecule Builder, a tool that helps you and your students work smarter and not harder when learning chemical bonding:

Atomsmith Molecule Builder

Figure 2. The Molecule Builder presents a palette of atoms (represented as Lewis symbols displaying each atom’s valence electrons). The available atoms, arranged in a periodic table, can be dragged and dropped into the model.

How the Model Works

According to Gilbert Lewis 1 , there are two simple “operations” that can be applied to electrons to form full octets:

  • sharing electrons to form covalent bonds
  • transferring electrons between atoms to form ionic bonds

The Molecule Builder allows the exact same two simple operations with electrons:

Click to play the animation. For optimal viewing, select full-screen mode.

Figure 3: Creating Ions by Transferring Electrons.

To create a pair of ions, drag an electron from one atom (or molecule) and drop it onto another atom (or molecule). Once the electron is transferred, there is one blue electron on the chlorine atom and the charges on the ions are displayed.

At any time, you can select Lewis Structure in the Labels menu to display (valence, shared, and lone pair) electrons as dots – without the circles that represent atoms.

Notice that once the single valence electron is removed from the (n=2) shell of the Na atom, a thin circle remains to emphasize that the original valence shell is now empty. And the inner (n=1) shell now displays a full octet.

You can reverse the ionization by dragging a valence electron from the negatively charged atom (or molecule) back onto the positively charged atom (or molecule).

Click to play the animation. For optimal viewing, select full-screen mode.

Figure 4: Creating Covalent Bonds by Sharing Electrons.

To share a pair of electrons between a pair of atoms, drag an electron from one atom and drop it onto any valence electron on another atom. The shared pair of electrons that represent a covalent bond will appear between the two atoms.

You can create multiple (double, triple) bonds by dragging additional valence electrons from one atom onto additional valence electrons on another atom.

You can also break a bond by dragging an electron from a shared pair back onto an atom. The remaining electron from the shared pair will once again become a valence electron on the other atom.

As you add atoms and move electrons, the model automatically adjusts the geometry to maxmize, in two dimensions (2D), the distances between electron domains. This is similar to how VSEPR predicts 3D molecular geometries. And as you will see below, the Molecule Builder can also convert 2D Lewis structures to 3D VSEPR molecular geometries.

How To Check Your Work

As you build a Lewis structure model in the Molecule Builder, it keeps track of the number of valence electrons around each atom in the model. At any time you can click the Check Model button to see if the model is correct – that all of the octets around each atom (or duets) have been filled.

Click to play the animation. For optimal viewing, select full-screen mode.

Figure 5: Checking the Model. When you click the Check Model button, the Molecule Builder will analyze the current Lewis structure and advise you if, where, and how your model is incorrect. Otherwise, it will advise you that the model is correct 👌!

If the model is incorrect, the Molecule Builder will inform you of any unfilled octets, and will highlight their locations by drawing pink circles where they occur. You can check the model as often as you like.

There is another way to know when you have built a correct Lewis structure model. If you select Molecule Formula/Name in the Labels menu (see Figure 5), the model will draw the formula for each “fragment” in the model. Once the Molecule Builder sees that you have built a correct model, it will label all of the molecules/ions with their IUPAC names.

[Note: The Molecule Builder allows “expanded octets” (i.e., “hypervalent structures”) for certain third row elements. It also allows for “electron-deficient” molecules (e.g., BF3). The Check Model button declares such structures as correct, and also explains that they are exceptions to the octet rule.]

Tools for Teachers

As a teacher, you may be wondering,

How could I use the Molecule Builder to demonstrate in front of the class?

If so, you may also be thinking,

I would not like to have to always take the time to build each of those Lewis structures from scratch.

That’s what the Molecule Builder’s Teachers Panel is for:

Click to play the animation. For optimal viewing, select full-screen mode.

Figure 6: Teachers Panel. The Teachers Panel (not available to students) contains a large collection of pre-built Lewis structures.

For example, here we use the Teachers Panel to load the Lewis structures (and molecular geometries) of vinegar (acetic acid) and baking soda (sodium bicarbonate).


Here’s how to use the Molecule Builder Lewis structure model to explore and explain the basis of many common chemical concepts.

Building Lewis Structures for Fun!

The purpose of the Molecule Builder is to enable students to improve their understanding of chemical bonding, and to strengthen their confidence in that understanding, by repeatedly building Lewis structures and receiving instant feedback on the structures they have built. . And we could not think of a better way to do this than to turn the Molecule Builder into a game.

So we did.

It’s called Octets!

Click to play the animation. For optimal viewing, select full-screen mode.

Figure 7: Playing the Octets! Game. Octets! is a Jeopardy®-like game board where students earn points by building molecules and ions, working their way to “Atomsmith Master” status.


Far and away, the most common use for writing Lewis Structures is as the first step in predicting molecular geometries.

As in:

  1. Draw a Lewis structure
  2. Count the electron domains
  3. Look up the geometry in a table

By now, how to perform steps 1 and 2 in the Molecule Builder should be clear.

When it comes to step 3, however, the Molecule Builder offers an alternative (Figure 8).

Click to play the animation. For optimal viewing, select full-screen mode.

Figure 8: VSEPR. Once you build a correct Lewis structure, click the Molecule Builder’s Load 3D Model button to load a 3D molecular model into Atomsmith’s 3D Model Window. Then click on connected atoms to measure lengths and angles. The H-O-H bond angle is confirmed to be 104.5°.

If you can build a correct Lewis structure model, the Molecule Builder can then build and display a 3D molecular geometry (see Figure 8). By default, the 3D model is displayed as a 3D Lewis Structure model: with bonds as cylinders and lone pairs as small spheres. (Btw, we invented this 3D Lewis structure model!)

Next you can measure bond lengths and bond angles by clicking on bonded atoms in the 3D model (also Figure 8). A table with the selected data appears.

We suggest this as a precursor to using the interactive molecular geometry table in the Atomsmith VSEPR Lab (Figure 9):

Click to play the animation. For optimal viewing, select full-screen mode.

Figure 9: VSEPR Lab. 3D structures can be loaded from the VSEPR Lab’s interactive Molecular Geometry table.

Periodic Trends

An additional way to use the Molecule Builder to reinforce periodic trends. Have your students examine the periodic table that contains the ”atoms” that get dragged into the model. Both:

  • the radii of these atoms, and
  • their numbers of valence electrons

follow the expected patterns – both across and up-and-down the table. The more time students spend using the Molecule Builder, the more time they will have to recognize and internalize these trends.

At a higher level of difficulty, you and your students can look for similarities and differences in bonding and geometries of central atoms from different Groups:

Click to play the animation. For optimal viewing, select full-screen mode.

Figure 10: Periodic Trends. Comparing the bonding and molecular geometries (both bent) of the Group 16 hydrides: H2O and H2S.

Bonus: If you wish to explain the different bond angles of these two molecules and the effects of their lone pairs, you can label the Lewis structure by their electronegativities.

Ionic Charges

Assigning the charges on a pair of monatomic ions is pretty intuitive if you understand what the columns of the periodic table mean.

But do your students struggle with where the charge on a polyatomic ion comes from?

Instead of suggesting that they just memorize it, have them build the ion’s Lewis structure first.

Then memorize it. For example:

Click to play the animation. For optimal viewing, select full-screen mode.

Figure 11: Ionic Charges. Building a (monatomic and polyatomic) pair of ions: Na+ and HO-. That blue electron on the oxygen atom is the extra electron that gives hydroxide its charge.

Ionic Formulas

Closely related to the concept of ionic charges is the concept of ionic formulas. Building Lewis structures can help with these also. For example:

Click to play the animation. For optimal viewing, select full-screen mode.

Figure 12: Ionic Formulas. First we build an Na+ ion and an S- ion. When we check the structure, the Molecule Builder tells us that the S atom does not have a full octet. So we build another sodium ion by dragging its valence electon onto the sulfur ion. And the Molecule Builder is happy. (Notice the two blue electrons on the sulfide ion.)

By building the Lewis structures, students learn that the formula unit of sodium sulfide is Na2S. Assign several of these to your students!

While you’re here, you can also use the Atomsmith Molecule Library to show your students that the Na2S formula unit is actually part of a large crystal lattice.

Click to play the animation. For optimal viewing, select full-screen mode.

Figure 13: Na2S Crystal Lattice. Search for “hydrogen sulfide” in the Molecule Library.

Chemical Reactions

Lewis structures are static, symbolic representations of matter – atoms, molecules and ions.

But a big part of chemistry concerns how matter changes.

You may recall that we mentioned above that – not only can you form bonds in the Molecule Builder – you can also break bonds.

When you dissolve ammonia in water, it forms a basic solution. Dissolved ammonia reacts with water to form ammonium and hydroxide ions.

Let’s see how to use Lewis structures to account for these chemical changes:

(Click to play the animation. For optimal viewing, select full-screen mode.)

Figure 14: A Chemical Reaction. NH3 + H2O ⇄ NH4+ + HO-

Note: This understanding can be reinforced using Atomsmith’s Reaction Lab.

Hopefully this blog post provides you with some ideas to help your students build their chemical intuition by building Lewis structures in the Atomsmith Molecule Builder. It is a powerful teaching and learning tool.

If you are not already a subscriber and you want to try out the Atomsmith Molecule Builder – as well as a multitude of other models – just request a trial license here:


  1. Lewis, G N. Valence and the Structure of Atoms and Molecules. New York: The Chemical Catalog Company, Inc., 1923.