7.3 Lewis Symbols and Structures

Lewis Symbols and Structures

Lewis symbols and structures give you a way to visualize valence electrons and how atoms bond together in molecules. Understanding how to draw and interpret these structures is foundational for predicting molecular shapes, reactivity, and properties throughout chemistry.

Lewis Symbols and Drawing Lewis Structures

A Lewis symbol represents an individual atom's valence electrons as dots placed around its chemical symbol, one dot per valence electron. Valence electrons are the electrons in the outermost shell of an atom, and they're the ones that participate in chemical bonding.

A Lewis structure goes further: it shows how atoms in a molecule or ion are connected and where all the valence electrons end up. Bonds between atoms are drawn as lines:

• Single bond: one shared pair of electrons (one line)
• Double bond: two shared pairs (two lines)
• Triple bond: three shared pairs (three lines)

Any valence electrons not involved in bonding are shown as lone pairs (dots) on individual atoms.

How to draw a Lewis structure:

1. Count total valence electrons. Add up the valence electrons from every atom in the molecule. For polyatomic ions, add one electron for each negative charge or subtract one for each positive charge.
2. Identify the central atom. The least electronegative atom goes in the center (hydrogen is never the central atom since it can only form one bond).
3. Connect atoms with single bonds. Draw a single bond from the central atom to each surrounding atom. Each bond uses 2 electrons.
4. Distribute remaining electrons as lone pairs. Place them on the outer atoms first, giving each atom (except hydrogen) an octet.
5. Check the central atom's octet. If the central atom doesn't have 8 electrons, convert lone pairs from a neighboring atom into double or triple bonds until it does.

Formal charge helps you pick the best Lewis structure when more than one arrangement is possible. It's calculated as:

$\text{Formal Charge} = \text{Valence electrons} - \text{Lone pair electrons} - \frac{1}{2}(\text{Bonding electrons})$

The most stable structure is the one where formal charges are closest to zero and any negative formal charges sit on the more electronegative atoms.

Octet Rule and Exceptions

The octet rule states that atoms tend to gain, lose, or share electrons until they have 8 electrons in their valence shell, mimicking the electron configuration of the nearest noble gas. Hydrogen is the obvious exception here: it only needs 2 electrons (a "duet") to match helium.

Ions follow this same principle. Cations lose electrons and anions gain electrons to reach a stable, noble-gas-like configuration.

However, several important exceptions exist:

• Odd-electron molecules: Molecules like $\text{NO}$ and $\text{NO}_2$ have an odd total number of valence electrons, so at least one atom can't have a complete octet. These species are called free radicals and tend to be highly reactive.
• Incomplete octets: Some atoms, especially beryllium and boron, are stable with fewer than 8 electrons. For example, $\text{BF}_3$ has only 6 electrons around boron, and $\text{BeH}_2$ has only 4 around beryllium. These electron-deficient molecules are often very reactive.
• Expanded octets: Elements in period 3 and beyond can accommodate more than 8 electrons because they have accessible d-orbitals. $\text{SF}_6$ has 12 electrons around sulfur, and $\text{PCl}_5$ has 10 around phosphorus. Only elements in period 3 or higher can do this, so you'll never see carbon or nitrogen with an expanded octet.

Resonance structures arise when more than one valid Lewis structure can be drawn for the same molecule. The actual molecule doesn't flip between these structures; instead, the real electron distribution is an average (or "hybrid") of all the resonance forms. A classic example is ozone ($\text{O}_3$), where the two oxygen-oxygen bonds are equivalent in reality, even though each individual Lewis structure shows one single bond and one double bond.

Molecular Shapes from Electron Domains

VSEPR theory (Valence Shell Electron Pair Repulsion) predicts molecular shapes based on one key idea: electron domains around a central atom repel each other and spread out as far apart as possible.

An electron domain is any region of electron density around the central atom. That includes single bonds, double bonds, triple bonds (each counts as one domain), and lone pairs.

The number of electron domains determines the electron-pair geometry:

The molecular geometry describes only the arrangement of atoms, ignoring lone pairs. Since lone pairs take up space but are "invisible" in the molecular shape, they change the geometry you'd predict. Lone pairs also compress bond angles slightly because they repel more strongly than bonding pairs.

Here are key examples to know:

• $\text{BeF}_2$: 2 bonding domains, 0 lone pairs → linear (180°)
• $\text{BF}_3$: 3 bonding domains, 0 lone pairs → trigonal planar (120°)
• $\text{CH}_4$: 4 bonding domains, 0 lone pairs → tetrahedral (109.5°)
• $\text{NH}_3$: 3 bonding domains, 1 lone pair → trigonal pyramidal (~107°)
• $\text{H}_2\text{O}$: 2 bonding domains, 2 lone pairs → bent (~104.5°)

Notice that $\text{CH}_4$, $\text{NH}_3$, and $\text{H}_2\text{O}$ all have 4 electron domains (tetrahedral electron-pair geometry), but their molecular geometries differ because lone pairs replace bonding pairs.

Molecular Structure and Bonding

To summarize the distinction: electron-pair geometry accounts for all electron domains (bonding and lone pairs), while molecular geometry describes only where the atoms are. You need the electron-pair geometry first to figure out the molecular geometry.

Hybridization explains why atoms form bonds that are equivalent in length and strength, even though their atomic orbitals (s, p, d) have different shapes. Atomic orbitals mix to form new hybrid orbitals that are all identical in energy and shape. The type of hybridization matches the electron-pair geometry:

• 2 electron domains → $sp$ hybridization (linear)
• 3 electron domains → $sp^2$ hybridization (trigonal planar)
• 4 electron domains → $sp^3$ hybridization (tetrahedral)

For example, carbon in $\text{CH}_4$ is $sp^3$ hybridized: one s orbital and three p orbitals combine to form four equivalent $sp^3$ orbitals arranged in a tetrahedron. This is why all four C-H bonds in methane are identical, even though the original s and p orbitals had different shapes and energies.