Chemical Bonding and Molecular Geometry

Overview

Chemical bonding is one of the most heavily tested topics on the MCAT's Chemical and Physical Foundations section. Understanding why and how atoms bond determines your ability to predict molecular shapes, polarities, intermolecular forces, physical properties, and reactivity. The MCAT expects you to integrate Lewis structures, VSEPR theory, orbital hybridization, and molecular orbital theory into a coherent framework for analyzing molecules.

This guide covers ionic bonding, covalent bonding, metallic bonding, Lewis structures, formal charges, resonance, VSEPR geometry, hybridization, sigma and pi bonds, molecular orbital theory, and intermolecular forces. Every organic chemistry and biochemistry topic you encounter downstream depends on these fundamentals.

Key Concepts

1. Types of Chemical Bonds

Ionic bonds form when electrons transfer from a metal to a nonmetal, creating oppositely charged ions held together by electrostatic attraction. Ionic compounds form crystal lattices, have high melting points, conduct electricity when dissolved or molten, and are generally soluble in polar solvents. Lattice energy (energy released when gaseous ions form a solid) increases with higher charge and smaller ionic radius.

Covalent bonds form when atoms share electrons, typically between nonmetals. Single bonds share one pair of electrons, double bonds share two pairs, and triple bonds share three pairs. Bond strength increases and bond length decreases with bond order: triple > double > single in strength. Polar covalent bonds occur when atoms with different electronegativities share electrons unequally, creating a dipole moment.

Metallic bonds involve a sea of delocalized electrons shared among metal cations. This model explains electrical conductivity, malleability, ductility, and metallic luster. Metallic bonding strength correlates with the number of valence electrons and the charge density of the metal cation.

2. Lewis Structures and Formal Charge

Lewis structures show valence electrons as dots or lines (bonds). To draw a Lewis structure: (1) count total valence electrons, (2) connect atoms with single bonds, (3) distribute remaining electrons to satisfy octets (starting with outer atoms), and (4) form multiple bonds if needed to complete octets. Formal charge = valence electrons - lone pair electrons - (1/2) bonding electrons. The best Lewis structure minimizes formal charges and places any negative formal charge on the most electronegative atom.

Some elements are exceptions to the octet rule. Hydrogen can only accommodate 2 electrons. Boron is stable with 6 electrons (electron-deficient). Elements in Period 3 and below can expand their octet using d orbitals (e.g., PCl5, SF6). Free radicals have unpaired electrons (e.g., NO, NO2).

3. Resonance

Resonance structures are multiple valid Lewis structures for a molecule where electrons are delocalized. The actual molecule is a hybrid of all resonance structures. Resonance stabilizes molecules by spreading charge over multiple atoms. Key examples include benzene (6 equivalent C-C bonds), the carbonate ion (3 equivalent C-O bonds), and the carboxylate group in amino acids. Resonance does not involve atoms moving — only electrons are redistributed. Greater resonance stabilization means lower energy and greater stability.

4. VSEPR Theory and Molecular Geometry

Valence Shell Electron Pair Repulsion (VSEPR) theory predicts molecular geometry by minimizing electron pair repulsion around the central atom. Count the steric number (bonding pairs + lone pairs) to determine the electron geometry, then identify the molecular geometry based on the positions of atoms only.

• Steric number 2: linear electron geometry (e.g., CO2, BeCl2), 180 degree bond angles

• Steric number 3: trigonal planar electron geometry (e.g., BF3), 120 degree bond angles; with 1 lone pair: bent (~117 degrees)

• Steric number 4: tetrahedral electron geometry (e.g., CH4), 109.5 degree bond angles; with 1 lone pair: trigonal pyramidal (~107 degrees); with 2 lone pairs: bent (~104.5 degrees)

• Steric number 5: trigonal bipyramidal (e.g., PCl5), 90/120 degree bond angles; see-saw, T-shaped, and linear geometries with lone pairs

• Steric number 6: octahedral (e.g., SF6), 90 degree bond angles; square pyramidal and square planar with lone pairs

Lone pairs occupy more space than bonding pairs, compressing bond angles below ideal values. This is why water (2 lone pairs on O) has a bond angle of 104.5 degrees rather than the tetrahedral 109.5 degrees.

5. Hybridization

Hybridization describes the mixing of atomic orbitals to form new hybrid orbitals for bonding. The hybridization matches the steric number: 2 = sp (linear), 3 = sp2 (trigonal planar), 4 = sp3 (tetrahedral). Each sigma bond or lone pair requires one hybrid orbital. Unhybridized p orbitals form pi bonds: sp2 carbons have one pi bond, sp carbons have two pi bonds.

In biological molecules: sp3 carbon is found in saturated compounds (alkanes, amino acid alpha-carbons), sp2 carbon is found in carbonyl groups, carboxylates, and aromatic rings, and sp carbon is rare but appears in nitriles and alkynes. Nitrogen and oxygen follow the same hybridization rules — count the steric number including lone pairs.

6. Sigma and Pi Bonds

Sigma bonds form from head-on orbital overlap and allow free rotation. Pi bonds form from side-by-side overlap of unhybridized p orbitals and restrict rotation. A single bond = 1 sigma. A double bond = 1 sigma + 1 pi. A triple bond = 1 sigma + 2 pi. Pi bonds are weaker than sigma bonds individually but contribute to overall bond strength. Restricted rotation around double bonds creates cis/trans isomerism, which is critical in biological molecules (e.g., cis vs. trans fatty acids, retinal isomerization in vision).

7. Molecular Orbital Theory

Molecular orbital (MO) theory describes bonding as the combination of atomic orbitals into molecular orbitals that span the entire molecule. When two atomic orbitals combine, they form one bonding MO (lower energy) and one antibonding MO (higher energy, denoted with *). Bond order = (bonding electrons - antibonding electrons) / 2. A bond order of 0 means the molecule does not exist.

For homonuclear diatomics: O2 has a bond order of 2 and two unpaired electrons in antibonding pi orbitals (explaining its paramagnetism). N2 has a bond order of 3 and is diamagnetic. The MCAT primarily tests MO theory for diatomic molecules — know the filling order and be able to calculate bond order and predict magnetic properties.

8. Intermolecular Forces

Intermolecular forces (IMFs) determine physical properties like boiling point, melting point, viscosity, and solubility. In order of increasing strength: London dispersion forces (present in all molecules, increase with molecular weight and surface area), dipole-dipole interactions (between polar molecules), hydrogen bonding (between H bonded to F, O, or N and a lone pair on F, O, or N), and ion-dipole forces (between ions and polar molecules, critical for dissolving salts in water).

Hydrogen bonding is essential for biological systems: it stabilizes DNA base pairing, protein secondary structure (alpha helices and beta sheets), and water's unique properties (high specific heat, high heat of vaporization, surface tension, density anomaly of ice). The MCAT frequently asks you to predict or compare boiling points based on IMF analysis.

High-Yield Facts

• Bond order: triple > double > single in strength; reverse for length

• Formal charge = valence e- minus lone pair e- minus half of bonding e-

• Resonance stabilizes molecules; more equivalent resonance structures = greater stability

• VSEPR: steric number = bonding pairs + lone pairs; determines electron geometry

• Lone pairs compress bond angles below ideal values

• Hybridization: sp3 = 109.5 degrees, sp2 = 120 degrees, sp = 180 degrees

• Single bond = 1 sigma; double = 1 sigma + 1 pi; triple = 1 sigma + 2 pi

• Pi bonds restrict rotation, causing cis/trans isomerism

• MO bond order = (bonding - antibonding) / 2; O2 is paramagnetic

• IMF strength: ion-dipole > H-bonding > dipole-dipole > London dispersion

• H-bonding requires H-F, H-O, or H-N with a lone pair acceptor

• Higher molecular weight = stronger London dispersion forces

• Lattice energy increases with charge and decreases with ionic radius

• Polar molecules have a net dipole moment; symmetric molecules may cancel dipoles

Common Mistakes

• Confusing electron geometry with molecular geometry — VSEPR gives electron geometry; molecular geometry considers only atom positions

• Forgetting to include lone pairs when determining hybridization and molecular geometry

• Assuming all molecules with polar bonds are polar — symmetric arrangements (like CO2) cancel dipole moments

• Mixing up formal charge and oxidation state — they use different counting rules

• Thinking resonance structures are different molecules — they represent delocalized electrons in a single molecule

• Forgetting that bond strength and bond length are inversely related

• Not accounting for expanded octets in Period 3+ elements (S, P, Xe, etc.)

• Confusing intermolecular forces with intramolecular bonds — IMFs are much weaker

Practice Strategy

For bonding questions, develop a systematic approach: (1) draw the Lewis structure, (2) determine formal charges, (3) identify the electron and molecular geometry using VSEPR, (4) assign hybridization, and (5) analyze polarity. This five-step process answers the vast majority of bonding questions.

Practice predicting boiling points for sets of molecules by analyzing their intermolecular forces. The MCAT loves ranking problems. Build a decision tree: first check for ionic character, then hydrogen bonding capability, then polarity, and finally molecular weight for London dispersion forces.

For molecular orbital questions, memorize the MO filling order for homonuclear diatomics up to Ne2. Practice calculating bond order and predicting paramagnetism. Remember that MO theory explains phenomena Lewis structures cannot, such as the paramagnetism of O2.