Lewis Structures, VSEPR Theory, and Molecular Polarity

Chemistry often feels abstract because you are trying to visualize 3D objects using 2D drawings. You draw a structure on a flat piece of paper, but in reality, that molecule has depth, angles, and a specific shape that dictates how it reacts with the world. To master chemical bonding, you need a systematic workflow to move from a flat formula to a three-dimensional prediction.

This guide bridges that gap. We will walk through the essential pipeline used by chemists: Lewis Structures $\rightarrow$ VSEPR Theory $\rightarrow$ Molecular Geometry $\rightarrow$ Polarity.

Drawing Accurate Lewis Dot Structures (Step-by-Step)

A Lewis Structure is a 2D map of the valence electrons in a molecule. It tells you how atoms are connected and where the lone pair electrons reside. Without a correct Lewis structure, all subsequent predictions about shape and polarity will be wrong.

The 5-Step Method

1. Count Total Valence Electrons: Sum the group numbers of all atoms. Add electrons for anions (-) and subtract for cations (+).
2. Determine the Central Atom: Usually the least electronegative element (never Hydrogen).
3. Form Single Bonds: Connect the central atom to surrounding atoms (subtract 2 electrons from your total for each bond).
4. Complete Octets for Surrounding Atoms: Add lone pairs to the outer atoms first.
5. Finish the Central Atom: Place remaining electrons on the central atom. If the central atom lacks an octet, convert lone pairs from outer atoms into double or triple bonds.

Illustrative Example: Carbon Tetrachloride ($CCl_4$)

Let’s apply the method to $CCl_4$ to see why it forms a stable octet.

• Step 1 (Count): Carbon (Group 14) has 4 valence $e^-$. Chlorine (Group 17) has 7 valence $e^-$. $$ \text{Total} = 4 + (4 \times 7) = 32 \text{ electrons} $$
• Step 2 (Skeleton): Carbon is less electronegative than Chlorine, so C goes in the center.
• Step 3 (Bond): Draw 4 single bonds connecting C to each Cl. This uses 8 electrons ($4 \times 2$). We have 24 left.
• Step 4 (Distribute): Place the remaining 24 electrons around the 4 Cl atoms. Each Cl gets 6 electrons (3 pairs). $$ 24 - (4 \times 6) = 0 $$
• Result: Every Chlorine has 8 electrons. The central Carbon shares 4 bonds (8 electrons). The octet rule is satisfied for all.

Advanced Case: Charged Ions ($PF_6^-$)

Don't panic when you see a charge. For the hexafluorophosphate anion ($PF_6^-$):

• Count: Phosphorus (5) + 6 Fluorines ($6 \times 7 = 42$) + 1 extra electron (due to -1 charge) = 48 valence electrons.
• Structure: P is bonded to 6 F atoms. This uses 12 electrons. The remaining 36 electrons fill the octets of the 6 Fluorine atoms.
• Exception: Phosphorus here has 12 valence electrons. This is an expanded octet, which is permissible for elements in Period 3 and below (like P, S, Se).

VSEPR Theory: Predicting Electron and Molecular Geometry

Once you have the Lewis structure, you must apply VSEPR Theory (Valence Shell Electron Pair Repulsion). The logic is simple: electron pairs (whether in bonds or lone pairs) are negatively charged. Like charges repel. Therefore, electron groups arrange themselves as far apart as possible around the central atom.

Electron Geometry vs. Molecular Geometry

This is the most common place where students lose marks. You must distinguish between the arrangement of all electron domains and the arrangement of atoms.

• Electron Geometry: Looks at all electron domains (bonding pairs + lone pairs).
• Molecular Geometry: Looks only at the atoms (ignoring lone pairs for the shape name, though the lone pairs dictate the angles).

Case Study 1: The "Bent" Molecule (Using $XCl_2$ and $CH_3OH$)

Consider a hypothetical molecule $XCl_2$ where the central atom $X$ has two bonding pairs and one lone pair.

• Electron Domains: 3 total (2 bonds + 1 lone pair).
• Electron Geometry: 3 domains arrange in a Trigonal Planar shape (120° apart).
• Molecular Geometry: Since one vertex of the triangle is invisible (the lone pair), the visible shape formed by Cl-X-Cl is Bent (or angular).
• Bond Angle: The lone pair repels more strongly than bonding pairs, compressing the angle to slightly less than 120°.

Similarly, in methanol ($CH_3OH$), looking at the oxygen atom:

• Oxygen has 2 bonds (one to C, one to H) and 2 lone pairs.
• Total domains: 4. Electron Geometry: Tetrahedral.
• Molecular Geometry: Bent (similar to water).

Case Study 2: The "T-Shaped" Molecule ($BrF_3$)

Let's look at Bromine Trifluoride ($BrF_3$).

1. Lewis Count: Br (7) + 3F (21) = 28 electrons.
2. Skeleton: Br bonded to 3 Fs (6 $e^-$ used). Remaining 22 $e^-$.
3. Lone Pairs: Fill F octets ($18 e^-$). 4 $e^-$ remain. Place them on Br as 2 lone pairs.
4. VSEPR Analysis:
   • Bonding Pairs: 3
   • Lone Pairs: 2
   • Total Domains: 5
5. Result:
   • Electron Geometry: Trigonal Bipyramidal.
   • Placement: Lone pairs always take the equatorial positions (more space).
   • Molecular Geometry: The atoms form a T-Shaped structure.

Molecular Polarity: Bond Dipoles and Overall Molecular Moment

A molecule is polar if it has a net dipole moment (an uneven distribution of electron density). This depends on two factors:

1. Bond Polarity: Are the atoms bonded together different in electronegativity?
2. Symmetry (Geometry): Do the bond dipoles cancel each other out?

Note: You cannot determine polarity from a Lewis structure alone because it is flat. You must apply VSEPR first to see the 3D shape.

Symmetry and Cancellation

If a molecule is perfectly symmetric, the pull of the electrons cancels out, making the molecule nonpolar, even if the individual bonds are polar.

• $CCl_4$ (Tetrahedral): C-Cl bonds are polar. However, the four bonds point to the corners of a tetrahedron. Vectors sum to zero. Nonpolar.
• $CS_2$ (Linear): The structure is $S=C=S$. The pull to the left cancels the pull to the right. Nonpolar.

Why Ozone ($O_3$) is Polar

Ozone consists of three identical oxygen atoms, so you might assume the bonds are nonpolar. However, $O_3$ has a dipole moment of 0.53 D. Why?

• The Geometry: The central Oxygen has a lone pair. This creates a Bent molecular geometry.
• Electron Density: The lone pair occupies a large volume of space, creating an area of high electron density on the central atom relative to the terminal atoms. Because the molecule is bent, this asymmetry does not cancel out. The result is a polar molecule.

Hybridization and Bond Types (Sigma and Pi Bonds)

To explain how these geometries form, chemists use Hybridization Theory. Atoms mix their standard atomic orbitals ($s, p, d$) to create new hybrid orbitals that match the VSEPR geometries.

Quick Guide to Hybridization

Count the electron domains (Sigma bonds + Lone pairs) around the central atom:

• 2 Domains $\rightarrow$ sp Hybridization: Linear geometry (e.g., $CO_2$).
• 3 Domains $\rightarrow$ sp² Hybridization: Trigonal planar geometry (e.g., $BF_3$, $NO_3^-$).
• 4 Domains $\rightarrow$ sp³ Hybridization: Tetrahedral geometry (e.g., $CH_4$, $H_2O$).
• 5 Domains $\rightarrow$ sp³d Hybridization: Trigonal bipyramidal (e.g., $PCl_5$).
• 6 Domains $\rightarrow$ sp³d² Hybridization: Octahedral (e.g., $SF_6$).

Sigma ($\sigma$) vs. Pi ($\pi$) Bonds

When drawing your structure, remember that multiple bonds are not just "more" of the same connection:

• Single Bond: Contains 1 $\sigma$ bond. (Head-on overlap).
• Double Bond: Contains 1 $\sigma$ bond + 1 $\pi$ bond. (Side-by-side overlap).
• Triple Bond: Contains 1 $\sigma$ bond + 2 $\pi$ bonds.

Key Takeaways: Chemical Bonding Summary

• Valence Electrons are Currency: Always start by calculating the total valence electrons. If this number is wrong, the structure is wrong.
• Lewis $\neq$ Shape: Lewis structures show connectivity. They are flat. You must count electron domains to find the 3D VSEPR shape.
• Lone Pairs Matter: Lone pairs repel more than bonding pairs. They compress bond angles and often create molecular polarity by breaking symmetry.
• Polarity Requires 3D Thinking: A molecule with polar bonds can be nonpolar overall if the geometry allows the dipoles to cancel (like in $CO_2$ or $CCl_4$).