7.6 Molecular Structure and Polarity

Molecular Geometry and Polarity

VSEPR theory predicts molecular shapes by figuring out how electron pairs arrange themselves to stay as far apart as possible. Once you know the shape, you can determine whether a molecule is polar or nonpolar, which directly affects how it behaves in reactions and solutions.

Molecular Structures Using VSEPR Theory

VSEPR stands for Valence Shell Electron Pair Repulsion. The core idea: electron pairs around a central atom repel each other, so they spread out as far as they can. That arrangement determines the molecule's shape.

Electron pairs come in two types:

• Bonding pairs are shared between two atoms
• Nonbonding pairs (lone pairs) sit on a single atom and aren't shared

Electron pair geometry counts all electron pairs (bonding and lone) around the central atom:

Molecular geometry only describes where the atoms are, ignoring lone pairs. This is why molecular geometry can differ from electron pair geometry. For example, water ($H_2O$) has 4 electron pairs (tetrahedral electron pair geometry), but since 2 of those are lone pairs, the molecular geometry is bent.

Common molecular geometries that differ from their electron pair geometry:

• Bent: 2 bonding pairs + 1 or 2 lone pairs ($H_2O$, $SO_2$)
• Trigonal pyramidal: 3 bonding pairs + 1 lone pair ($NH_3$, $PH_3$)

Lone pairs take up slightly more space than bonding pairs, so they compress bond angles. In $CH_4$ the bond angle is 109.5°, but in $NH_3$ (one lone pair) it drops to about 107°, and in $H_2O$ (two lone pairs) it's about 104.5°.

Polar vs. Nonpolar Covalent Bonds

Covalent bonds form when atoms share electrons. Whether that sharing is equal depends on electronegativity, which measures how strongly an atom attracts electrons in a bond. Fluorine is the most electronegative element, followed by oxygen, nitrogen, and chlorine.

• Nonpolar covalent bonds form between atoms with equal or very similar electronegativities. The electrons are shared roughly equally. Examples: $H-H$ in $H_2$, $C-H$ in $CH_4$.
• Polar covalent bonds form between atoms with different electronegativities. Electron density shifts toward the more electronegative atom, creating a partial negative charge ($\delta^-$) on that atom and a partial positive charge ($\delta^+$) on the other. Examples: $O-H$ in $H_2O$, $N-H$ in $NH_3$.

The dipole moment ($\mu$) quantifies how polar a bond is. It's calculated as:

$\mu = Q \times r$

where $Q$ is the charge separation and $r$ is the bond length. A larger dipole moment means a more polar bond. For the hydrogen halides, polarity follows this trend: $HF > HCl > HBr > HI$, because fluorine's electronegativity creates the largest charge separation.

Molecular Polarity Analysis

A molecule's overall polarity depends on two things: the polarity of its individual bonds and its geometry. A molecule can have polar bonds and still be nonpolar overall if those bond dipoles cancel each other out due to symmetry.

Nonpolar molecules either have no polar bonds or have polar bonds arranged symmetrically so the dipoles cancel:

• $CO_2$ has two polar $C=O$ bonds, but its linear shape points them in exactly opposite directions. They cancel, so $CO_2$ is nonpolar.
• $CCl_4$ has four polar $C-Cl$ bonds, but its tetrahedral symmetry means the dipoles cancel perfectly.

Polar molecules have polar bonds arranged asymmetrically, producing a net dipole moment:

• $H_2O$ has two polar $O-H$ bonds in a bent shape. The dipoles point in roughly the same direction and don't cancel, making water polar.
• $NH_3$ has three polar $N-H$ bonds in a trigonal pyramidal shape. The lone pair on nitrogen prevents the dipoles from canceling.

To determine whether a molecule is polar, follow these steps:

1. Identify the polar bonds by comparing electronegativities of the bonded atoms.
2. Determine the molecular geometry using VSEPR theory (count bonding and lone pairs).
3. Check whether the dipole moments cancel based on the geometry.
   • Symmetric arrangement → dipoles cancel → nonpolar
   • Asymmetric arrangement → net dipole → polar

Lewis Structures and Valence Electrons

Lewis structures show how valence electrons (the outermost electrons) are arranged in a molecule. They're the starting point for predicting molecular geometry because you need them to count bonding pairs and lone pairs before applying VSEPR theory.

To draw a Lewis structure:

1. Count the total valence electrons from all atoms in the molecule. For ions, add electrons for negative charges or subtract for positive charges.
2. Place the least electronegative atom in the center (hydrogen is always on the outside).
3. Draw single bonds between the central atom and each surrounding atom. Each bond uses 2 electrons.
4. Distribute remaining electrons as lone pairs, starting with the outer atoms, to satisfy the octet rule.
5. If the central atom lacks an octet, convert lone pairs on outer atoms into double or triple bonds.

Hybridization describes how atomic orbitals mix to form new hybrid orbitals that match the observed geometry. For this course, the key connections are:

• 2 electron groups → $sp$ hybridization → linear
• 3 electron groups → $sp^2$ hybridization → trigonal planar
• 4 electron groups → $sp^3$ hybridization → tetrahedral