Module 1: Structure & Bonding

This module, Atomic Structure and Bonding, is aligned with the AAMC’s Official MCAT Content Outline under Category 1A: Structure and Properties of Atoms, forming the foundational basis for all general and organic chemistry reasoning on the MCAT. A deep understanding of atomic theory, periodic trends, and chemical bonding is essential for interpreting reactivity, molecular behavior, and experimental results throughout the exam. Our lessons follow the AAMC emphasis on conceptual mastery, quantitative reasoning, and experimental application, ensuring that every topic covered reflects what is actually tested on MCAT passages and discrete questions. From subatomic particles to molecular orbitals, this module builds the atomic framework upon which the rest of general and organic chemistry depends.

Covalent Bonding: The Foundation of Organic Molecules

Organic chemistry is fundamentally the study of carbon-based compounds, and nearly all of these compounds are held together by covalent bonds. A covalent bond forms when two atoms share a pair of electrons, each contributing one electron to the bond. This sharing allows atoms to achieve a stable electron configuration, typically resembling the noble gas configuration (octet rule).

In contrast to ionic bonding, where one atom donates electrons and another accepts them (forming cations and anions), covalent bonding maintains discrete molecular structures with no full charges on atoms. This type of bonding underlies the unique diversity and complexity of organic molecules.

Why Carbon?

Carbon’s ability to form four stable covalent bonds makes it uniquely suited to serve as the backbone of life. With four valence electrons, carbon can form single, double, or triple bonds and can bond with other carbon atoms indefinitely, creating chains, rings, and complex 3D networks.

Molecular Orbital Theory and Bond Formation

At the quantum level, bonding is explained by the overlap of atomic orbitals, regions in space where electrons are likely to be found. When two atomic orbitals overlap constructively, they form a bonding molecular orbital, which stabilizes the molecule.

There are two major types of covalent bonds:

Bond Type	Description	Strength	Example
Sigma (σ) bond	Formed by head-on overlap of orbitals (s-s, s-p, or p-p)	Strongest type of covalent bond	C–C in ethane
Pi (π) bond	Formed by side-to-side overlap of unhybridized p orbitals	Weaker than sigma bond	C=C in ethene (the second bond)

• A single bond is always a sigma bond
• A double bond consists of one sigma and one pi bond
• A triple bond has one sigma and two pi bonds

MCAT Atomic Structure and Bonding Insight: The presence of pi bonds introduces regions of electron density above and below the molecular plane, making them more chemically reactive (e.g., in addition reactions).

Electronegativity and Bond Polarity

Covalent bonds can be classified based on how evenly electrons are shared:

• Nonpolar covalent bonds: equal sharing (e.g., H–H, C–C)
• Polar covalent bonds: unequal sharing due to differences in electronegativity (e.g., O–H, C=O)
• Ionic bonds (not truly covalent): complete transfer of electrons (e.g., Na⁺Cl⁻)

The greater the difference in electronegativity, the more polar the bond becomes. This influences molecular dipole moments, intermolecular forces, and reactivity — all testable on the MCAT.

Summary of MCAT Atomic Structure and Bonding Key Concepts

Concept	Explanation
Covalent bond	Sharing of electrons between atoms to form a stable molecule
Sigma bond	First bond formed, with head-on overlap of orbitals
Pi bond	Additional bond formed from side-by-side p orbital overlap
Electronegativity	A measure of an atom’s ability to attract shared electrons
Bond polarity	Unequal sharing leads to partial charges and reactivity

Hybridization (sp, sp², sp³)

What Is Hybridization?

Hybridization is a model that explains how atomic orbitals mix to form new, equivalent hybrid orbitals that define the geometry and bonding patterns of molecules. This concept helps explain why methane (CH₄), for example, forms a perfect tetrahedron with four identical C–H bonds — even though carbon’s valence electrons occupy different types of orbitals (2s and 2p) in its ground state.

In atoms that form multiple bonds or have lone pairs, s and p orbitals combine mathematically to create hybrid orbitals — named sp, sp², and sp³ — which better align with observed molecular geometries and bond angles.

MCAT Atomic Structure and Bonding Insight: Hybridization tells you both the shape of a molecule and the number of electron domains (bonding + lone pairs) around an atom.

Types of Hybridization

Each type of hybridization corresponds to a specific electron geometry and number of electron domains. Here’s a breakdown:

sp³ Hybridization – Tetrahedral Geometry

• Orbitals mixed: one s + three p orbitals
• Number of domains: 4 (e.g., 4 single bonds or 2 bonds + 2 lone pairs)
• Bond angles: ~109.5°
• Geometry: Tetrahedral (e.g., CH₄), trigonal pyramidal (e.g., NH₃), or bent (e.g., H₂O)
• Example: Methane (CH₄), where each C–H bond is formed by overlap of a C sp³ orbital with an H 1s orbital.

On the MCAT, sp³ hybridized atoms often have single bonds only, and molecules adopt three-dimensional shapes.

sp² Hybridization – Trigonal Planar Geometry

• Orbitals mixed: one s + two p orbitals
• Number of domains: 3
• Bond angles: ~120°
• Geometry: Trigonal planar
• Unhybridized p orbital: One p orbital is left unhybridized → used for π bonding
• Example: Ethene (C₂H₄), where each carbon forms three sigma bonds using sp² orbitals and participates in a double bond with a neighboring carbon via p–p overlap

On the MCAT, sp² indicates a double bond or a carbocation center (which only needs three bonds).

sp Hybridization – Linear Geometry

• Orbitals mixed: one s + one p orbital
• Number of domains: 2
• Bond angles: 180°
• Geometry: Linear
• Unhybridized p orbitals: Two p orbitals remain unhybridized → form two π bonds (i.e., a triple bond)
• Example: Acetylene (C₂H₂), where each carbon is sp hybridized and forms a triple bond (σ + 2π) with the other carbon

sp hybridization is often found in alkynes or nitriles (–C≡N), and implies linear geometry with strong, directional bonding.

Determining Hybridization: Step-by-Step MCAT Strategy

1. Count the number of electron domains (bonds + lone pairs) around the atom:
   • 2 domains → sp
   • 3 domains → sp²
   • 4 domains → sp³
2. Ignore double and triple bonds when counting domains — each counts as one domain.
3. Don’t forget lone pairs! Lone pairs contribute to geometry and hybridization.

Visual Summary Table

Hybridization	Bonded Domains	Geometry	Bond Angle	Example	π Bonds Possible
sp³	4	Tetrahedral	~109.5°	CH₄, NH₃, H₂O	None
sp²	3	Trigonal Planar	~120°	C=C, CH₂=CH₂	1 π bond
sp	2	Linear	~180°	C≡C, HC≡CH	2 π bonds

Hybridization and Molecular Reactivity

The degree of hybridization affects electron density and bond characteristics:

• sp orbitals have more s-character (50%) → closer to nucleus, more electronegative, stronger bonds
• sp² has 33% s-character
• sp³ has 25% s-character → bonds are longer and less strong than sp or sp²

This influences acidity (e.g., terminal alkynes are more acidic than alkenes or alkanes) and reactivity with electrophiles/nucleophiles.

Molecular Geometry and VSEPR Theory

What Is VSEPR Theory?

Valence Shell Electron Pair Repulsion (VSEPR) theory is a simple, yet powerful model used to predict the 3D shape of molecules. It assumes that electron pairs around a central atom repel each other and will orient themselves as far apart as possible to minimize repulsion. These electron pairs include both:

• Bonding pairs (shared between atoms)
• Lone pairs (non-bonding electrons localized on a single atom)

The geometry predicted by VSEPR determines bond angles, molecular shape, and polarity — all of which influence how molecules react, interact, and function biologically.

MCAT Insight: Recognizing VSEPR shapes helps you identify the hybridization, molecular polarity, and biological behavior of key MCAT molecules like water, ammonia, and carbon dioxide.

Electron Domains vs. Molecular Geometry

To determine geometry, first count the total number of electron domains around the central atom (single, double, and triple bonds each count as 1 domain, as do lone pairs).

Electron Domains	Electron Geometry	Molecular Geometry	Example	Hybridization	Bond Angle
2	Linear	Linear	CO₂	sp	180°
3	Trigonal planar	Trigonal planar	BF₃	sp²	120°
3	Trigonal planar	Bent (1 lone pair)	SO₂	sp²	<120°
4	Tetrahedral	Tetrahedral	CH₄	sp³	109.5°
4	Tetrahedral	Trigonal pyramidal (1 lone pair)	NH₃	sp³	~107°
4	Tetrahedral	Bent (2 lone pairs)	H₂O	sp³	~104.5°

Note: Lone pairs compress bond angles slightly due to stronger repulsion compared to bonding pairs.

Key MCAT Structure and Bonding Examples and Why They Matter

Water (H₂O): Bent geometry

• 2 bonds, 2 lone pairs → sp³ hybridized
• Bond angle: ~104.5°
• Highly polar → key for hydrogen bonding, solvation, and acid–base reactions

Carbon Dioxide (CO₂): Linear geometry

• 2 double bonds → 2 domains → sp hybridized
• Bond angle: 180°
• Symmetrical → nonpolar despite polar bonds

Ammonia (NH₃): Trigonal pyramidal

• 3 bonds, 1 lone pair → sp³ hybridized
• Bond angle: ~107°
• Polar molecule → relevant to nucleophilicity and hydrogen bonding

Summary and Strategy for MCAT Structure and Bonding

• Step 1: Count electron domains (bonds + lone pairs)
• Step 2: Determine electron geometry
• Step 3: Use number of lone pairs to determine molecular geometry
• Step 4: Use shape and symmetry to assess polarity

Concept	Significance on MCAT
Electron geometry	Based on total domains
Molecular geometry	Actual 3D shape
Lone pairs	Reduce bond angles, affect polarity
Polarity	Affects boiling point, solubility, reactivity

Formal Charge, Resonance, and Stability

1. Formal Charge: Bookkeeping for Electrons

Formal charge (FC) is a method of keeping track of the electrons “owned” by an atom in a Lewis structure. It helps identify the most stable resonance forms, determine reactivity centers, and validate valid Lewis structures on the MCAT.

Formula for Formal Charge:

$$
\text{Formal charge} = \text{valence electrons} – \text{nonbonding electrons} – \frac{1}{2}(\text{bonding electrons})
$$

Or equivalently:

$$
\text{FC} = V – L – \frac{B}{2}
$$

Where:

• V = number of valence electrons
• L = number of lone-pair electrons
• B = number of shared (bonding) electrons

Example: Ammonium Ion (NH₄⁺)

• Nitrogen has 5 valence electrons
• In NH₄⁺, it forms 4 bonds and has 0 lone pairs
• FC = 5 – 0 – 4 = +1

MCAT Tip: Always minimize formal charges and place negative charges on electronegative atoms (like O, Cl) to favor stability.

2. Resonance Structures

Resonance occurs when two or more valid Lewis structures can be drawn for the same molecule or ion by moving only electrons, not atoms. The real structure is a hybrid — a weighted average of all contributing resonance forms.

Rules of Resonance:

• Only pi electrons and lone pairs move; sigma bonds remain fixed
• Atoms do not move
• Total charge and number of electrons must be conserved

Resonance Contributors: Evaluating Stability

Feature	Effect on Stability
Full octets	↑ stability
Fewer formal charges	↑ stability
Negative charge on electronegative atom	↑ stability
Positive charge on less electronegative atom (e.g. C)	↑ stability
Separation of opposite charges	↓ stability

On the MCAT for Structure and Bonding, you should be able to draw all major resonance structures, assess their relative stability, and determine which ones dominate the hybrid.

Example: Acetate Ion (CH₃COO⁻)

The negative charge can be delocalized between the two oxygen atoms:

O⁻ ←→ O⁻

Both forms are equivalent → high resonance stability.

3. Resonance and Delocalization = Stability

What Is Resonance?

Resonance describes a situation where two or more valid Lewis structures can be drawn for a molecule or ion by moving only electrons — not atoms. These different depictions are called resonance structures, and the actual molecule is not any one of them, but rather a resonance hybrid — a weighted average of all contributors.

This electron delocalization distributes charge or bond character across multiple atoms, lowering energy and stabilizing the molecule.

MCAT Atomic Structure and Bonding Tip:

While the MCAT rarely asks you to draw specific resonance structures or compare contributors in detail, it does expect you to understand the concept of resonance. Focus on the big-picture ideas:

• Resonance delocalizes charge, which stabilizes molecules.
• This stabilization can affect reactivity, acidity, and structure.
• Know that resonance hybrids are more stable than any single contributor.

You don’t need to memorize or generate resonance structures on test day — just understand what resonance does and why it matters.

Key Characteristics of Resonance

• Atoms do not move — only electrons (pi bonds and lone pairs) shift between adjacent atoms.
• Total charge remains constant across resonance forms.
• The real molecule exists as a resonance hybrid: not flipping between forms, but an average of them.
• Resonance structures are not isomers — they are alternate depictions of the same species.

Rules for Drawing Resonance Structures

1. Only move electrons: pi electrons and lone pairs
2. Never exceed the octet on second-row elements (C, N, O, F)
3. Maintain the same atom positions and bonding framework
4. Preserve net charge across all structures
5. Use curved arrows to show the flow of electron pairs

Evaluating Resonance Structures: Which Are Most Stable?

Feature	Stability Impact
All atoms have complete octets	↑ Increases stability
Minimal formal charges	↑ Increases stability
Negative charges on electronegative atoms (O, Cl, etc.)	↑ Stability
Positive charges on electropositive atoms (C, H)	↑ Stability
Avoid charge separation when possible	↑ Stability if charges are close together and stabilized

MCAT Atomic Structure and Bonding Tip: The more stable the resonance contributor, the more it contributes to the hybrid.

Delocalization = Resonance in Action

Delocalization refers to the spreading of electrons (especially lone pairs and π electrons) across multiple atoms via overlapping p-orbitals. This lowers energy by distributing electron density, which:

• Stabilizes charges (e.g., in carboxylates)
• Lowers reactivity of intermediates (e.g., allylic carbocations)
• Explains color and UV–Vis absorption in conjugated systems

Examples Where Resonance and Delocalization Matter on the MCAT

Molecule / Ion	Resonance/Delocalization Role
Acetate (CH₃COO⁻)	Negative charge delocalized between two oxygen atoms
Amide (peptide bond)	Partial double bond character → explains rigidity of protein backbone
Benzene	Six π electrons delocalized in cyclic fashion → aromaticity
Allylic carbocation	Positive charge delocalized across two carbon atoms → stabilized
Nitro group (NO₂⁻)	Delocalized negative charge over two oxygen atoms
Phenol	Lone pair on oxygen delocalized into aromatic ring → acidity ↑

Conjugation and Resonance

• Conjugation is a specific form of delocalization where alternating single and double bonds allow π electrons to overlap across three or more atoms.
• Conjugated systems are often colored (e.g., β-carotene) and absorb UV–visible light — useful in biological MCAT contexts (e.g., retinal, heme).
• Many biologically important molecules (nucleotides, heme, chlorophyll) contain conjugated resonance systems.

4. Stability Considerations in Organic Molecules

In organic chemistry, stability is a recurring theme that governs everything from reactivity and acid-base behavior to mechanisms and intermediate lifetimes. The MCAT frequently tests whether you can predict relative stability of molecules, ions, and intermediates, especially in unfamiliar contexts. This section walks through the major principles you must master.

A. Charge Stability: Where Do Charges Want to Live?

Charged species — especially carbocations and anions — are unstable and will rearrange, react, or delocalize to become more stable. Stability of charged species depends on their atomic environment:

Stabilizing Factors:

• Electronegativity: Negative charges prefer electronegative atoms (e.g., O⁻ more stable than C⁻).
• Atom size: Larger atoms spread out negative charge better (e.g., S⁻ more stable than O⁻).
• Delocalization: Resonance spreads charge over multiple atoms, decreasing energy.
• Inductive effects: Electronegative groups nearby can stabilize (+) or (–) charges via electron withdrawal.

Destabilizing Factors:

• Localized charge on the “wrong” atom (e.g., a + charge on oxygen)
• Strain or inability to delocalize or resonate
• Crowding or repulsion from nearby groups

B. Carbocation Stability: Tertiary > Secondary > Primary > Methyl

Carbocations are a high-yield MCAT topic — they are common reaction intermediates, especially in SN1, E1, and rearrangement mechanisms.

Order of Stability:

3∘>2∘>1∘>methyl

Why?

• Hyperconjugation: Alkyl groups donate electron density via σ bonds to stabilize the empty p orbital.
• Inductive effects: Nearby electron-donating groups help stabilize the positive charge.
• Resonance: Allylic and benzylic carbocations are especially stable due to delocalization.

MCAT Atomic Structure and Bonding Tip: If you’re asked to compare two possible carbocation intermediates in a mechanism, always look for resonance stabilization first, then degree (3°, 2°, etc.).

C. Carbanion Stability: Opposite Trends

Carbanions (negatively charged carbon species) show reverse stability trends: methyl>1∘>2∘>3∘

• Why? Electron-donating groups (alkyl chains) destabilize a negative charge on carbon.
• Carbanions prefer to be:
  • On less substituted carbons
  • Adjacent to electron-withdrawing groups
  • Stabilized by resonance (e.g., enolates)

D. Radical Stability: Similar to Carbocations

Radicals are uncharged species with an unpaired electron. Their stability trend mirrors that of carbocations: 3∘>2∘>1∘>methyl

• Radicals are stabilized by hyperconjugation and resonance.
• Allylic and benzylic radicals are very stable.
• MCAT may test this in the context of halogenation (e.g., Br₂ vs Cl₂ selectivity).

E. Resonance Stabilization (Reinforced)

As discussed in the previous section, resonance is a major stabilizing factor for:

• Charged species (e.g., carboxylate anion)
• Intermediates (e.g., allylic cation or radical)
• Functional groups (e.g., amides, conjugated systems)

Key Insight: The more you can spread out a charge (via resonance or induction), the lower the energy and the more stable the structure.

F. Inductive Effects and Electronegative Neighbors

Inductive effects refer to the pulling or pushing of electron density through σ bonds by atoms or groups.

• Electron-withdrawing groups (e.g., –NO₂, –CN, –CF₃) stabilize negative charges and acids.
• Electron-donating groups (e.g., alkyl chains, –OH, –OR) stabilize positive charges.

For example: Trifluoroacetic acid (TFA) is more acidic than acetic acid due to strong –F inductive effects.

G. Aromaticity and Conjugation

Some molecules exhibit extra stability due to aromaticity — a special type of cyclic resonance:

• Aromatic compounds must be cyclic, planar, conjugated, and obey Hückel’s Rule: 4n+24n + 24n+2 π electrons
• Examples: Benzene, pyrrole, imidazole

Aromaticity explains the unusual inertness and low reactivity of many biological and pharmaceutical compounds.

H. Strain and Ring Stability

Cyclic structures may be destabilized by ring strain, which can be:

• Angle strain: Deviation from ideal bond angles (e.g., cyclopropane)
• Torsional strain: Eclipsing interactions in small rings
• Steric strain: Atoms bumping into each other (especially in bulky substituents)

Stability trend (3–6 membered rings):

Cyclohexane>Cyclopentane>Cyclobutane>Cyclopropane

• Chair conformation of cyclohexane minimizes all strain and is the most stable.
• Axial vs. equatorial positioning is important in stereochemistry and reactivity.

MCAT Structure and Bonding Strategy Summary: Predicting Stability

Stability Factor	What to Look For
Resonance	Delocalization of charge or electrons
Inductive effects	Electronegative atoms stabilizing nearby charges
Substitution level	More alkyl groups = better for carbocations, worse for carbanions
Aromaticity	Cyclic conjugation + Hückel’s rule
Ring strain	Smaller rings or eclipsing atoms = less stable
Electronegative atoms	Negative charges → EN atoms; positive charges → less EN atoms

Now that we’ve established how to evaluate molecular structure and stability — including charge distribution, resonance, and inductive effects — we’re ready to apply these insights to predict how organic molecules behave. Up next, we’ll explore acids and bases, nucleophiles and electrophiles, and the principles that govern organic reactivity on the MCAT. Understanding molecular stability is your key to unlocking the logic behind every reaction mechanism that follows.

Acids, Bases, Nucleophiles, and Electrophiles

Conceptual Overview

Understanding how organic molecules behave during chemical reactions starts with identifying what roles they play: are they acting as acids, bases, nucleophiles, or electrophiles? These classifications define how molecules donate or accept electrons, and ultimately determine how and where bonds form or break.

This section bridges the gap between structure and reactivity. You’ll learn to recognize electron-rich vs. electron-poor regions, how molecular properties influence reaction pathways, and how to use simple rules (like pKa or charge) to predict who reacts with whom and why.

These are the battle lines of organic chemistry: understanding who attacks, who defends, and who wins. The MCAT tests this relentlessly, so it’s crucial you don’t just memorize — you develop instincts for molecular behavior.

Acid and Base Definitions (Brønsted–Lowry and Lewis)

There are three major frameworks for defining acids and bases. The MCAT focuses mostly on the first two:

Brønsted–Lowry Definition (Most Common on the MCAT):

• Acid = proton (H⁺) donor
• Base = proton (H⁺) acceptor

Examples:

• HCl donates H⁺ → acid
• OH⁻ accepts H⁺ → base

This framework explains reactions like neutralization (acid + base → salt + water) and ties directly into the pKa scale used to measure acid strength.

Lewis Definition (Broader, MCAT-relevant):

• Lewis Acid = electron pair acceptor
• Lewis Base = electron pair donor

This expands the scope beyond H⁺ transfers and is crucial for understanding:

• Nucleophilic substitution (SN1/SN2)
• Electrophilic addition
• Coordination complexes

Examples:

• BF₃ accepts electrons → Lewis acid
• NH₃ donates a lone pair → Lewis base

MCAT Atomic Structure and Bonding Tip: All Brønsted–Lowry acids/bases are Lewis acids/bases, but not all Lewis acids/bases involve H⁺.

Arrhenius Definition (Low Yield):

• Acid = increases [H⁺] in water
• Base = increases [OH⁻] in water
  Mainly used for aqueous solutions and rarely tested beyond general chemistry.

Acid and Base Strength: Understanding pKa

The strength of an acid or base isn’t just about whether it donates or accepts a proton — it’s about how easily it does so. The MCAT quantifies this behavior using pKa, a logarithmic measure of acid strength.

What Is pKa?

The acid dissociation constant Ka​ reflects the extent to which an acid dissociates in water:

$$
\text{HA} \leftrightarrow \text{H}^+ + \text{A}^-
$$

The pKa is defined as:

$$
\text{p}K_a = -\log K_a
$$

• Low pKa = strong acid (dissociates easily)
• High pKa = weak acid (holds onto its proton)

Rule of Thumb:

• Strong acids: pKa < 0
• Moderate acids: pKa 0–5
• Weak acids: pKa > 5
• Very weak acids (like water or alcohols): pKa ≈ 15–18
• Alkanes/alkenes: pKa > 40 → negligible acidity

MCAT Atomic Structure and Bonding Tip: A lower pKa means the conjugate base is more stable. Stability = strength.

Relative Strength Table

Acid	pKa	Comment
HCl	–7	Very strong acid
H₂SO₄ (1st proton)	–3	Strong diprotic acid
CH₃COOH (acetic acid)	4.75	Common weak acid
NH₄⁺ (ammonium)	9.25	Weak acid, conjugate of NH₃
H₂O	15.7	Very weak acid
CH₃CH₃ (ethane)	~50	Essentially no acidity

Acid–Base Equilibrium and Favorability

A key MCAT Atomic Structure and Bonding Strategy:
Acid–base reactions favor the formation of the weaker acid and base.

So if you’re comparing two acids:

• The one with higher pKa is favored at equilibrium (weaker acid).
• This helps you predict reaction direction.

Summary Table – Acid/Base Definitions

Definition	Acid	Base	Example Acid	Example Base
Arrhenius	Produces H⁺ in water	Produces OH⁻ in water	HCl	NaOH
Brønsted–Lowry	Donates H⁺	Accepts H⁺	NH₄⁺	NH₃
Lewis	Accepts electron pair	Donates electron pair	BF₃	H₂O or NH₃

Nucleophiles and Electrophiles

Organic chemistry revolves around electron movement — and in most MCAT reactions, that movement begins with a nucleophile attacking an electrophile. Think of them as the two opposing forces in every reaction mechanism:

• Nucleophiles = electron-rich species (attackers)
• Electrophiles = electron-poor species (targets)

Nucleophiles

Nucleophiles are species that donate a lone pair or a π-bond to form a new bond. They’re typically:

• Negative ions (e.g., OH⁻, CN⁻)
• Neutral species with lone pairs (e.g., H₂O, NH₃)
• π bonds (e.g., alkenes, alkynes)

“Nucleo-” refers to the nucleus → nucleophiles seek a nucleus (i.e., a positively charged or electron-deficient atom).

Common Features of Nucleophiles:

• Lone pairs or π electrons
• Often negatively charged
• Attack the electrophilic carbon in a reaction

Examples:

• OH⁻, RO⁻, CN⁻, NH₃, Cl⁻, Br⁻, I⁻
• Double bonds in alkenes (in electrophilic addition reactions)

Electrophiles

Electrophiles are electron-deficient species that accept electrons to form a new bond. These include:

• Carbocations (e.g., CH₃⁺)
• Partially positive atoms (e.g., δ⁺ carbon in C=O)
• Atoms bonded to electronegative leaving groups (e.g., alkyl halides)

“Electro-” refers to electrons → electrophiles seek electrons from nucleophiles.

Common Features of Electrophiles:

• Partial or full positive charge
• Electron-deficient atoms or carbon centers
• Accept lone pairs or π electrons

Examples:

• C=O carbon in ketones and aldehydes
• Alkyl halides (R–Br, R–Cl)
• Carbocations (R⁺)

Nucleophile vs. Base: What’s the Difference?

Property	Nucleophile	Base
Function	Attacks atoms	Attacks protons (H⁺)
Favors	Substitution (SN1, SN2)	Elimination (E1, E2)
Example	OH⁻ attacking C in CH₃Br	OH⁻ removing a β-H from butane

MCAT Atomic Structure and Bonding Tip: Many species (like OH⁻ or NH₂⁻) can act as both, but the reaction conditions determine which behavior dominates.

Predicting Reactivity Using pKa and Charge

Understanding why molecules behave the way they do in acid-base or nucleophile–electrophile reactions comes down to two central ideas:

1. The stability of the conjugate base or intermediate
2. The driving forces of equilibrium (i.e., pKa)

The MCAT expects you to predict which direction an acid–base reaction will favor, whether a molecule is likely to react, and how charge, electronegativity, resonance, and hybridization affect these behaviors.

How pKa Predicts Direction of Acid–Base Reactions

Acid–base reactions always favor the formation of the weaker acid (i.e., the one with the higher pKa). That means:

$$
\text{Stronger acid (low } pK_a\text{)} + \text{Strong base} \rightarrow \text{Weaker acid (high } pK_a\text{)} + \text{Weaker base}
$$

MCAT Atomic Structure and Bonding Rule: The equilibrium favors the side with the more stable conjugate base.

Example:

Acetylene (HC≡CH) has a pKa of ~25.
Water (H₂O) has a pKa of ~15.7.

So reacting NaNH₂ (conjugate base of NH₃, pKa ~38) with acetylene will deprotonate it:

$$
\mathrm{HC{\equiv}CH + NH_2^- \rightarrow HC{\equiv}C^- + NH_3}
$$

This reaction favors products, because NH₃ is a weaker acid than HC≡CH.

Charge and Reactivity

• Negative charges are often more nucleophilic (more reactive)
• Positive charges are usually more electrophilic (wanting electrons)
• But stabilized charges = lower reactivity
  • e.g., a negative charge resonance-stabilized on an enolate is less nucleophilic than on a regular alkoxide

General Rule: The more unstable or unstabilized the charge, the more reactive the species.

Electronegativity and Hybridization

• More electronegative atoms hold onto negative charges better → weaker bases, more stable
  • (e.g., F⁻ is more stable than O⁻ or N⁻)
• sp-hybridized anions (like acetylide) are more stable than sp² or sp³ anions
  • This explains the higher acidity of alkynes (pKa ~25) over alkenes (~40) or alkanes (~50)

Resonance Stabilization

Resonance allows a charge to delocalize, reducing its reactivity.

• Conjugate bases with resonance are more stable
  • e.g., acetate (CH₃COO⁻) is much more stable than ethoxide (CH₃CH₂O⁻)
  • So acetic acid is stronger than ethanol (pKa ~4.7 vs. ~16)

MCAT Atomic Structure and Bonding Tip: Don’t memorize pKa values. Just compare them and know general trends (e.g., carboxylic acids < alcohols < amines < alkanes).