D7.1 Hydrocarbons: Why Use Lewis Structures?

Let’s start with what molecules actually look like—well, our best guess. Shown below is space-filling models of simple molecules made only of carbon and hydrogen: methane (CH₄), ethane (C₂H₆), and propane (C₃H₈). These are all examples of hydrocarbons—molecules composed entirely of carbon and hydrogen atoms.

As you examine their structures, some patterns should stand out:

• Carbon (dark atoms) tends to form four connections.
• Hydrogen (light atoms) always forms one connection.
• The shapes are not flat—they’re three-dimensional, with atoms arranged to minimize repulsion between electron regions.

These 3D space-filling representations give you a tangible sense of how atoms pack together. They show approximate relative sizes, angles, and how tightly atoms are bonded. But they’re also hard to sketch and annotate, especially as molecules get larger or more complex.

To make structure easier to communicate, chemists use a 2D system called Lewis structures—a symbolic shorthand for representing molecules in two dimensions. Lewis structures represent bonding and valence electrons using simple conventions:

• Bonds are shown as lines.
• Lone electron pairs (if any) are shown as dots.
• Atom connectivity is drawn in flat space.

Instead of memorizing these conventions right away, we will unpack them from the 3D models. You can see, for instance, that in every structure above, carbon appears bonded to four other atoms. Lewis structures encode that information with four lines extending from each carbon.

3D space-filling models are great for visualizing how atoms occupy space and how they’re arranged. But they’re not so great for drawing on a whiteboard, annotating with electron counts, or quickly comparing two different molecules. That’s why chemists developed Lewis structures. Let’s go back to our example molecules: methane, ethane, and propane. If we translate their 3D structures into 2D Lewis structures, we notice some clear conventions:

• Each carbon makes four lines (bonds).
• Each hydrogen connects with one line (bond).
• We show electrons only when they’re not in bonds.

These are not arbitrary rules—they’re patterns that emerge from the molecular models themselves. Lewis structures don’t create the structure; they reflect it.