Module 2: Chemical Bonding and Molecular Structure

Ionic vs. Covalent Bonding

Chemical bonding explains how atoms combine to form more complex structures — molecules, salts, and networks. The two primary types of chemical bonds tested on the MCAT are ionic and covalent. These differ in how electrons are distributed between atoms, which in turn determines the compound’s physical and chemical properties.

Ionic Bonds

Definition: Ionic bonds form when one atom donates one or more electrons to another atom, resulting in positive and negative ions that are held together by electrostatic attraction.

Typically occur between metals (cations) and nonmetals (anions).
Electron is transferred from the metal to the nonmetal.
The resulting oppositely charged ions create a strong crystalline lattice.

Example:

NaCl (table salt): Na → Na⁺ + e⁻; Cl + e⁻ → Cl⁻ → Na⁺Cl⁻

Key Characteristics of Ionic Compounds:

High melting and boiling points due to strong ionic bonds.
Brittle solids that fracture along cleavage planes.
Conduct electricity in molten or aqueous form (electrolytes).
Soluble in water, especially if the ions are small and highly charged.

MCAT Tip:

A large electronegativity difference (>1.7) between atoms generally suggests ionic bonding.

Common Ionic Compounds:

Compound	Cation	Anion	Physical State	Conductivity
NaCl	Na⁺	Cl⁻	Solid (RT)	Yes (aq/molten)
MgCl₂	Mg²⁺	Cl⁻	Solid	Yes
CaCO₃	Ca²⁺	CO₃²⁻	Solid	Low solubility

Covalent Bonds

Definition: Covalent bonds occur when two nonmetals share one or more pairs of electrons to fill their valence shells and achieve stability (octet rule).

Atoms share electrons within overlapping orbitals.
These form discrete molecules with predictable geometries.

Types of Covalent Bonds:

Nonpolar covalent: Equal electron sharing (e.g., H₂, O₂, CH₄)
Polar covalent: Unequal sharing due to electronegativity difference (e.g., HCl, H₂O)

Key Characteristics of Covalent Compounds:

Lower melting/boiling points than ionic compounds.
Poor conductors of electricity (no free ions).
Exist in gaseous, liquid, or soft solid form at room temperature.
Bond strength increases with bond order:
- Single < Double < Triple

MCAT Tip:

Electronegativity difference:
- < 0.5 → nonpolar covalent
- 0.5–1.7 → polar covalent

Covalent Bond Examples:

Molecule	Bond Type	Shape	Polarity	Notes
H₂O	Polar	Bent	Yes	Strong hydrogen bonding
CO₂	Polar bonds, nonpolar molecule	Linear	No	Dipoles cancel
CH₄	Nonpolar	Tetrahedral	No	Symmetrical sharing

Summary Table: Ionic vs Covalent Bonds

Property	Ionic Bond	Covalent Bond
Electron Movement	Transferred (metal → nonmetal)	Shared (between nonmetals)
Electronegativity Diff	> 1.7	0.5–1.7 (polar), <0.5 (nonpolar)
Physical State	Crystalline solids at RT	Gas, liquid, soft solids
Electrical Conductivity	Yes (in aqueous/molten state)	No (generally non-conductive)
Solubility in Water	Usually soluble	Varies
Examples	NaCl, MgCl₂, KBr	H₂O, NH₃, CH₄, CO₂

MCAT Strategy Tip: Master identifying bond type based on element location:

Metal + Nonmetal = Ionic
Nonmetal + Nonmetal = Covalent

Polyatomic ions often form mixed ionic/covalent compounds (e.g., NH₄Cl, NaNO₃)

Valence Electrons – Definition, Patterns, and Importance

Valence electrons are the electrons located in the outermost electron shell of an atom. These are the electrons involved in chemical bonding, determining how atoms interact, form molecules, and react chemically. On the MCAT, understanding valence electrons helps you quickly assess bonding potential, molecular shape, and periodic behavior.

Key Facts:

Valence electrons = outermost electrons in the highest principal quantum number (n).
For main group elements, the group number indicates the number of valence electrons.
For transition metals, d-electrons are sometimes included depending on chemical behavior.

Example:

Carbon (Group 14) → 4 valence electrons → forms 4 covalent bonds.
Oxygen (Group 16) → 6 valence electrons → forms 2 bonds, 2 lone pairs.

Table: Groups and Valence Electrons (Main Group Elements)

Group	Elements	Valence Electrons
1	H, Li, Na, K…	1
2	Be, Mg, Ca…	2
13	B, Al, Ga…	3
14	C, Si, Ge…	4
15	N, P, As…	5
16	O, S, Se…	6
17	F, Cl, Br…	7
18	He, Ne, Ar…	8 (He has 2)

MCAT Tips:

Master valence electron counting — it drives Lewis structures, oxidation states, and formal charges.
Remember that noble gases are inert because they already have full valence shells.
Transition metals often lose s-electrons before d-electrons — don’t be fooled by their electron configurations!

Connection to Bonding

The number of valence electrons determines how many bonds an atom will typically form.
Atoms bond to complete their octets (or duets for hydrogen), using valence electrons.

Quick Check: How many valence electrons do phosphorus (P), chlorine (Cl), and magnesium (Mg) have? Predict the number of bonds each would form.

Lewis Structures & Octet Rule

Lewis structures are two-dimensional diagrams used to represent how atoms are bonded in a molecule and where the valence electrons reside. These diagrams are essential tools for predicting molecular geometry, bond polarity, formal charges, resonance, and reactivity—all of which are fair game on the MCAT.

The Octet Rule

Most atoms tend to bond in a way that gives them eight valence electrons, achieving the stable electron configuration of noble gases. This rule drives much of covalent bonding.

Hydrogen: Exception — follows the duet rule (needs 2 electrons).
Incomplete octets: Common in boron (B) and beryllium (Be).
Expanded octets: Atoms in period 3 and beyond (e.g., phosphorus, sulfur) can accommodate more than 8 electrons due to accessible d orbitals.

Steps for Drawing Lewis Structures

Count total valence electrons:
- For polyatomic ions, add electrons for negative charges, subtract for positive ones.
Choose a central atom:
- Usually the least electronegative element (never hydrogen).
Draw single bonds (each = 2 electrons) connecting atoms to the center.
Distribute remaining electrons to outer atoms as lone pairs to fulfill octets.
Place leftover electrons on central atom.
Form double or triple bonds as needed to satisfy octets.
Check formal charges and use resonance if applicable.

Example: CO₂

Valence electrons: 4 (C) + 6×2 (O) = 16
Central atom: Carbon
Structure:

O = C = O

Each oxygen has 4 nonbonding electrons (2 lone pairs), and carbon is surrounded by 8 electrons via double bonds.

Common Lewis Structures

Molecule	Total e⁻	Structure	Notes
H₂O	8	H–O–H (2 lone pairs)	Bent shape due to lone pairs
CH₄	8	Tetrahedral	4 bonding pairs around C
NH₃	8	Trigonal pyramidal	1 lone pair on N
BF₃	24	Incomplete octet	Only 6e⁻ around B
SF₆	48	Expanded octet (12e⁻)	6 F atoms bonded to S

MCAT Tips

Always verify total e⁻ count matches what’s available.
Formal charges of 0 are favored in stable molecules.
Resonance structures improve stability by delocalizing electrons.
Watch for ions and use brackets with charge labels (e.g., [NH₄]⁺).
Knowing the difference between valid vs. most stable structure can make a difference.

Diagrams

1. Lewis Dot for Water (H₂O):

2. Lewis Dot for Ammonia (NH₃):

3. Lewis Dot for Methane (CH₄):

Concept Check: Try drawing Lewis structures for NO₃⁻, SO₂, and CH₂O. Pay close attention to resonance and formal charge.

Coordinate Covalent Bonds

Definition: A coordinate covalent bond (also called a dative bond) is a type of covalent bond in which both electrons in the shared pair come from the same atom. This differs from a traditional covalent bond where each atom contributes one electron.

Key Features:

Often occurs when a lone pair on one atom is donated to an empty orbital on another.
Common in Lewis acid–base reactions where a base donates a lone pair to an acid.
Once formed, the bond behaves like any other covalent bond.

Examples:

Molecule/Ion	Explanation
NH₄⁺ (ammonium)	N donates a lone pair to bond with H⁺ (which has no electrons).
H₃O⁺ (hydronium)	O donates a lone pair to bond with H⁺.
Metal Complexes	Ligands donate electron pairs to central metal ion (e.g., Fe²⁺).

MCAT Tip:

Watch for positive ions like H⁺ that can’t form bonds on their own, these are frequent candidates for coordinate bonds.

Visualization:

Regular Covalent Bond:

A–B (each contributes 1 e⁻)

Coordinate Bond:

A: → B (both electrons from A)

Why It Matters: Coordinate bonds are crucial in acid-base chemistry, biochemistry (e.g., enzyme-metal interactions), and transition metal complexes. Understanding their formation helps explain stability, charge distribution, and reactivity in molecules.

Formal Charge & Resonance

Formal charge and resonance are essential concepts for evaluating the stability and accuracy of Lewis structures. On the MCAT, these concepts are frequently tested alongside molecular geometry, electronegativity, and reactivity.

Formal charge is a method used to assign electron ownership in a molecule, helping us determine the most stable and likely Lewis structure. It is not a real charge, but rather a formalism that compares the number of electrons an atom “owns” in a molecule to its number of valence electrons.

Formal Charge Formula:

Formal Charge = (Valence electrons) − (Nonbonding electrons) − ½(Bonding electrons)

Key Rules:

Lower formal charges (closer to 0) are generally more stable.
Structures where negative formal charges are placed on the most electronegative atoms are more favorable.
The sum of all formal charges must equal the overall charge of the molecule or ion.

Example: Carbon Dioxide (CO₂)

Valence e⁻: Carbon (4), Oxygen (6 × 2 = 12)
Total = 16 valence electrons
Structure: O = C = O
Oxygen (double bonded): FC = 6 − 4 (lone) − ½×4 (bonding) = 0
Carbon: FC = 4 − 0 − ½×8 = 0 → All atoms have zero formal charge → very stable!

Resonance – Delocalization & Stability

Resonance structures are alternate ways of drawing the same molecule that differ only in the placement of electrons (not atom positions). They represent a more accurate picture of bonding, where electrons are delocalized across multiple atoms.

Resonance Guidelines:

All resonance forms must follow normal bonding rules (valid octets/duets).
The actual molecule is a resonance hybrid, not flipping back and forth.
Structures with minimal formal charges and charge on the correct atoms are most significant.

Common Resonance Examples:

Molecule	Resonance Forms	Notes
O₃ (ozone)	O–O=O ↔ O=O–O	Lone pairs & double bonds alternate
NO₃⁻	N with 1 double bond, 2 single bonds	Electrons delocalized over 3 oxygens
Benzene (C₆H₆)	Alternating single/double bonds	Classic aromatic resonance; delocalized π e⁻

Summary Table: Formal Charge & Resonance

Feature	Formal Charge	Resonance
Purpose	Estimate electron distribution	Show delocalization of electrons
Involves	e⁻ ownership (bonding vs. lone pair)	Positioning of π electrons/lone pairs
Goal	Identify most stable Lewis structure	Depict blended real structure (hybrid)
Stability rule	Favor 0 or low FC; negatives on EN atoms	All forms must be valid; hybrid is real

MCAT Tips

Resonance increases molecular stability — memorize common examples (e.g., nitrate, carboxylate).
Always calculate formal charges when given multiple Lewis structures.
Resonance hybrids distribute charge/electron density more evenly, reducing reactivity

Quick Practice: Draw all resonance structures for NO₂⁻ and assign formal charges. Which structure is most stable?

VSEPR Geometry & Molecular Shape

Understanding the 3D shape of molecules is essential for predicting their physical and chemical properties, such as polarity, reactivity, and intermolecular interactions. The Valence Shell Electron Pair Repulsion (VSEPR) theory provides a model to determine molecular geometry based on the idea that electron pairs around a central atom repel each other and thus arrange themselves as far apart as possible.

VSEPR Theory: Key Principles

Electron domains are regions of electron density, including:
- Bonding pairs (single, double, or triple bonds count as one domain each)
- Nonbonding (lone) electron pairs
These electron domains repel each other and try to get as far apart as possible around the central atom.
The number of electron domains determines the electron geometry.
The molecular geometry is derived by ignoring lone pairs and focusing only on atom positions.
Lone pairs repel more strongly than bonding pairs, compressing bond angles.

Step-by-Step: How to Use VSEPR to Predict Molecular Geometry

Draw the Lewis Structure
- Determine total valence electrons and construct a complete Lewis diagram.
Count the Electron Domains Around the Central Atom
- Each single, double, or triple bond = 1 domain
- Each lone pair = 1 domain
Determine the Electron Geometry
- Use the number of electron domains to classify the geometry (e.g., 4 domains = tetrahedral).
Determine the Molecular Geometry
- Remove lone pairs from the visual picture.
- Identify the shape created by the atoms only (e.g., 2 bonding + 2 lone pairs = bent).
Estimate Bond Angles and Consider Lone Pair Effects
- Expect bond angles to shrink slightly with lone pairs due to increased repulsion.
Apply to MCAT-Style Interpretation
- Predict polarity, intermolecular forces, and molecular behavior from shape.

Common VSEPR Geometries

Electron Domains	Electron Geometry	Molecular Geometry	Bond Angle	Example
2	Linear	Linear	180°	CO₂
3	Trigonal planar	Trigonal planar	120°	BF₃
3	Trigonal planar	Bent (1 lone pair)	<120°	SO₂
4	Tetrahedral	Tetrahedral	109.5°	CH₄
4	Tetrahedral	Trigonal pyramidal (1 LP)	<109.5°	NH₃
4	Tetrahedral	Bent (2 LPs)	~104.5°	H₂O
5	Trigonal bipyramidal	Trigonal bipyramidal	90°, 120°	PCl₅
5	Trigonal bipyramidal	Seesaw (1 LP)	<90°, <120°	SF₄
6	Octahedral	Octahedral	90°	SF₆
6	Octahedral	Square pyramidal (1 LP)	<90°	BrF₅
6	Octahedral	Square planar (2 LPs)	90°	XeF₄

Important Clarifications

Electron geometry describes all electron regions (bonds and lone pairs).
Molecular geometry describes only the shape formed by atoms.
Lone pairs cause bond angle compression: e.g., CH₄ = 109.5°, NH₃ ≈ 107°, H₂O ≈ 104.5°.
Expanded octets (5 or 6 electron domains) occur in elements from Period 3 and beyond (e.g., P, S, Cl).

In trigonal bipyramidal structures, lone pairs prefer equatorial positions to minimize repulsion.

MCAT Strategy Notes

You must be able to identify molecular shape from a Lewis structure and count electron domains.
Don’t confuse molecular shape with electron geometry. Always distinguish them on the MCAT.
Bond angles are approximate and decrease with more lone pairs.
Polarity depends on both molecular shape and bond polarity. A symmetrical molecule (e.g., CO₂) may be nonpolar even if it contains polar bonds.

MCAT Favorite: Water (H₂O) is a bent polar molecule, explaining its high boiling point and hydrogen bonding ability.

Polarity, Dipole Moments, and Intermolecular Forces

Molecular polarity plays a key role in determining physical properties such as solubility, boiling point, reactivity, and interactions with biological molecules. A solid grasp of polarity and intermolecular forces will help you make sense of molecular behavior in both chemical and biological contexts.

Bond Polarity

Bond polarity arises from unequal sharing of electrons between atoms in a covalent bond due to differences in electronegativity — the tendency of an atom to attract electrons in a bond.

Nonpolar Covalent Bonds: Electrons are shared equally.
- Electronegativity difference: < 0.5
- Example: H₂, Cl₂, CH₄
Polar Covalent Bonds: Electrons are shared unequally.
- Electronegativity difference: 0.5–1.7
- Example: HCl, H₂O, NH₃
Ionic Bonds: Electrons are fully transferred.
- Electronegativity difference: > 1.7
- Example: NaCl, CaF₂

Dipole Moment Formula:

$$
\mu = q \cdot r
$$

Where:

$$
\text{where:} \
\mu = \text{electric dipole moment (C·m)} \
q = \text{magnitude of the charge (C)} \
r = \text{distance between the charges (m)}
$$

Molecular Polarity

Even if individual bonds are polar, the overall molecular polarity depends on the molecule’s shape (VSEPR geometry) and the orientation of the polar bonds.

Symmetrical molecules (e.g., CO₂, CCl₄) have dipoles that cancel → nonpolar overall.
Asymmetrical molecules (e.g., H₂O, NH₃) have net dipoles → polar.
A molecule with lone pairs on the central atom is often polar due to asymmetry.

Dipole Moment (μ): A vector quantity representing the magnitude and direction of bond polarity.

The greater the dipole moment, the more polar the molecule.
Units: Debye (D)

MCAT Tip: Look at both shape and bond polarities to determine whether a molecule is polar.

Examples: Predicting Molecular Polarity

Molecule	Shape	Bond Polarity	Symmetrical?	Net Dipole?	Polar?
CO₂	Linear	Yes	Yes	No	No
H₂O	Bent	Yes	No	Yes	Yes
CH₄	Tetrahedral	No	Yes	No	No
NH₃	Trigonal pyramidal	Yes	No	Yes	Yes
CCl₄	Tetrahedral	Yes	Yes	No	No

Intermolecular Forces (IMFs)

IMFs are non-covalent interactions between molecules. While much weaker than covalent or ionic bonds, they strongly influence boiling point, solubility, phase behavior, and biological recognition.

Type of IMF	Description	Relative Strength	Example
London Dispersion	Temporary induced dipoles in all molecules	Weakest	CH₄, noble gases
Dipole-Dipole	Interaction between permanent dipoles	Moderate	HCl, acetone
Hydrogen Bonding	H bound to N, O, or F interacts with lone pairs	Strong	H₂O, NH₃, HF
Ion-Dipole	Ion interacts with polar molecule	Strongest	Na⁺ in H₂O, Cl⁻ in NH₃

Hydrogen Bonds require:
1. A hydrogen covalently bonded to N, O, or F
2. A lone pair on a nearby N, O, or F

MCAT Favorite: Hydrogen bonding explains why water has an unusually high boiling point and why DNA base pairs stay together.

Key MCAT Insights

Polarity affects solubility: “Like dissolves like.” Polar molecules dissolve in polar solvents (e.g., H₂O); nonpolar molecules dissolve in nonpolar solvents (e.g., hexane).
Intermolecular forces affect physical properties: Stronger IMFs = higher boiling/melting points.
Hydrogen bonding is crucial in biomolecular structure: DNA base pairing, protein folding, and enzyme-substrate recognition.
Molecules with similar molar masses can have vastly different boiling points due to differences in IMFs.

Hybridization, σ and π Bonds, and Partial Ionic Character

Orbital Hybridization

Definition: Hybridization describes how atomic orbitals mix to form new, degenerate orbitals that participate in bonding.
Types:
- sp: Linear geometry, 180° (e.g., BeCl₂)
- sp²: Trigonal planar, 120° (e.g., BF₃)
- sp³: Tetrahedral, 109.5° (e.g., CH₄)

Hybridization	Geometry	Bond Angle	Example
sp	Linear	180°	CO₂
sp²	Trigonal planar	120°	SO₂
sp³	Tetrahedral	109.5°	CH₄

MCAT Tip: Count electron domains around the central atom to determine hybridization (2 = sp, 3 = sp², 4 = sp³).

Sigma (σ) and Pi (π) Bonds

σ bonds: Formed by head-on orbital overlap. Every single bond contains one sigma bond.
π bonds: Formed by sideways overlap of p orbitals. Found in double and triple bonds.
- Double bond = 1 σ + 1 π
- Triple bond = 1 σ + 2 π

Bond Type	Orbital Overlap	Relative Strength	Found In
Sigma (σ)	Head-on (s or p orbitals)	Stronger	All single bonds
Pi (π)	Sideways (p orbitals)	Weaker	Double, triple bonds

MCAT Insight: Pi bonds prevent rotation around double/triple bonds, important in understanding molecular shape and isomerism.

Comparative Case Study

Compound	Formula	IMF Type(s)	Boiling Point (°C)	Polar?
Ethanol	C₂H₅OH	Dispersion, Dipole, Hydrogen Bonding	78.5	Yes
Dimethyl ether	C₂H₆O	Dispersion, Dipole	-24.8	Yes

Though both have the same molecular formula (C₂H₆O), ethanol can hydrogen bond, dramatically increasing its boiling point.

Partial Ionic Character

Concept: Even “covalent” bonds can have some ionic character if there’s a significant difference in electronegativity.
The greater the electronegativity difference, the more polar the bond behaves.
Water (H₂O), for instance, has significant partial charges on O and H due to high ΔEN.

MCAT Application: Polar covalent bonds may exhibit partial ionic character, explaining phenomena like hydrogen bonding and dipole interactions.

Solubility Logic Table

Solute	Solvent	IMF Match?	Soluble?
NaCl	H₂O	Yes	Yes
Hexane	H₂O	No	No
Ethanol	H₂O	Yes	Yes

Metallic Bonds & Network Solids

Chemical bonding isn’t limited to ionic and covalent types. Some substances, particularly metals and certain crystalline solids, exhibit distinct bonding frameworks. On the MCAT, understanding the structure and properties that result from metallic bonding and network covalent bonding is essential for predicting conductivity, melting points, and hardness.

Metallic Bonds

Metallic bonding occurs between metal atoms that collectively share a “sea of electrons.” These delocalized electrons are free to move throughout the solid structure, giving metals their characteristic properties.

Key Features of Metallic Bonds

Delocalized electrons: Electrons are not bound to individual atoms.
Malleability and ductility: Metal atoms can slide past one another without breaking bonds.
High electrical and thermal conductivity: Due to free-moving electrons.
Luster: Electrons reflect light easily, giving metals a shiny appearance.

Property	Explanation
Conductivity	Free electrons carry charge and thermal energy
Malleability & Ductility	Atoms shift without breaking metallic bonds
Shiny (Luster)	Free electrons reflect light across the surface
Insolubility	Most metals are insoluble in water or organic solvents

Examples: Copper (Cu), Iron (Fe), Aluminum (Al)

Network Covalent Solids

These are structures where atoms are held together by a continuous network of covalent bonds, rather than discrete molecules. This results in materials with very high melting points and hardness.

Characteristics

Very high melting and boiling points: Due to extended covalent bonding throughout the solid.
Hard and brittle: Bond breakage requires breaking a continuous lattice.
Poor conductors: No free electrons (exception: graphite).

Substance	Bonding Type	Electrical Conductivity	Notes
Diamond	Network covalent	None	3D tetrahedral lattice of carbon
Graphite	Network covalent	Yes (along planes)	Layers held by dispersion forces
Quartz (SiO₂)	Network covalent	None	Silicon and oxygen lattice

MCAT Tip: Remember graphite as an exception due to delocalized electrons in planes.

Summary Table of Bond Types & Properties

This overview brings together key features of different types of chemical bonds, helping you compare and contrast their characteristic, a common MCAT expectation.

Bond Type	Particle Types Involved	Electron Behavior	Physical Properties	Examples
Ionic	Metal + Nonmetal	Electrons transferred	High melting point, soluble in water, conducts in solution	NaCl, MgCl₂
Covalent (Polar)	Nonmetals	Electrons shared unequally	Medium melting point, soluble, may conduct if polar	H₂O, NH₃
Covalent (Nonpolar)	Nonmetals	Electrons shared equally	Low melting point, not soluble, non-conductive	CH₄, O₂
Metallic	Metals	Delocalized electrons	Malleable, ductile, conductive, shiny	Cu, Fe, Al
Network Covalent	Nonmetals in 2D or 3D lattices	Electrons fixed in bonds	Very high melting point, hard, brittle, poor conductivity	Diamond, Quartz

MCAT Tip: Focus on how electron behavior influences bulk properties — conductivity, solubility, hardness, melting/boiling points.