Chapter 1: Structure Determines Properties – Foundations of Organic Chemistry

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Structure Determines Properties

Introduction

The physical and chemical properties of organic molecules are determined by their atomic structure and bonding. Understanding the arrangement of atoms, electrons, and the types of bonds present is fundamental to predicting reactivity and behavior in organic chemistry.

Cyclohexane and cyclopentane ring structure with ball-and-stick model

Atoms, Electrons, and Orbitals

Particles and Symbols of the Atom

Atoms are composed of three fundamental particles: protons, neutrons, and electrons. The atomic number (Z) is defined by the number of protons in the nucleus and determines the element's identity.

Proton: Positively charged, mass ≈ kg
Neutron: Neutral, mass ≈ kg
Electron: Negatively charged, mass ≈ kg

Diagram of atom showing protons, neutrons, and electrons

Electrons as Waves

Electrons exhibit wavelike behavior, described by the Schrödinger equation. The wavefunction () represents the probability density of finding an electron in a particular region of space, known as an orbital. The energy and shape of an orbital reflect the stability and spatial distribution of electrons.

Quantum Numbers

Atomic orbitals are characterized by four quantum numbers:

Principal quantum number (n): Energy level of the orbital
Orbital quantum number (l): Shape of the orbital
Magnetic quantum number (m): Orientation of the orbital
Spin quantum number (s): Electron spin direction (+1/2 or -1/2)

The Pauli exclusion principle states that no two electrons in an atom can have the same set of quantum numbers.

Quantum numbers for 2p_x orbital

The s Atomic Orbitals

s orbitals are spherically shaped and begin at n = 1. Each shell contains one s orbital, which is lower in energy than other orbitals in the same shell. As n increases, the number of nodes (regions of zero probability) increases.

Spherical s orbital diagram Comparison of 1s and 2s atomic orbitals

The p Atomic Orbitals

p orbitals begin at n = 2 and are dumbbell-shaped. Each np subshell contains three p orbitals (px, py, pz), oriented along the Cartesian axes. All p orbitals have a node at the nucleus and are equal in energy. Hund’s rule states that electrons fill these orbitals singly before pairing.

Shapes of p atomic orbitals

Ionic Bonds

Coulomb’s Law and Ionic Bonding

Ions are atoms or molecules with electric charge. Oppositely charged ions attract each other according to Coulomb’s law, forming ionic bonds. Ionic compounds typically form tightly packed lattices, resulting in high-melting solids that are soluble in polar solvents.

NaCl lattice structure

Ionic Bonding and Electron Transfer

Ionic bonds often arise from the transfer of electrons from a metal to a nonmetal. For example, sodium (Na) transfers an electron to chlorine (Cl), forming Na+ and Cl- ions. The formation of solid NaCl is exothermic due to the release of lattice energy.

Ionic Bonding in Organic Chemistry

While ionic bonds are common in salts, most organic compounds feature covalent bonds, where electrons are shared between nonmetal atoms such as C, N, O, and H.

Covalent Bonds, Lewis Formulas, and the Octet Rule

The Lewis Model of Covalent Bonding

Covalent bonds involve the sharing of electrons between two atoms. Atoms share electrons to achieve a stable electron configuration, often isoelectronic with the nearest noble gas.

Electron sharing in covalent bond

Covalent Bonding in H2

Each hydrogen atom needs one more electron to complete its valence shell. By sharing electrons, both achieve a configuration analogous to helium.

Dot structure for H2 Line structure for H2

Covalent Bonding in F2

Each fluorine atom needs one more electron to complete its octet. By sharing electrons, both achieve a configuration analogous to neon. The remaining electrons are nonbonding pairs (lone pairs).

Lewis Symbols and Structures

Lewis structures use dots to represent electrons and lines to represent shared pairs in covalent bonds. Unshared pairs are shown as dots on the atom’s edges. Electrons are localized on a single atom or between two atoms.

The Octet Rule

Second-row elements achieve stability by sharing electrons until they have eight in their valence shell (octet). When checking the octet rule, shared electrons are counted for both atoms involved.

Electron distribution around atom

Examples of Lewis Structures

Common molecules such as methane, ammonia, water, and hydrogen fluoride illustrate the application of Lewis structures and the octet rule.

Compound	Atom	Number of valence electrons in atom	Atom and sufficient number of hydrogen atoms to complete octet	Lewis formula (Dot)	Lewis formula (Line)
Methane	Carbon	4	H–C–H (four hydrogens)	Dot structure showing electron bonds between each C–H	Line structure showing C bonded to H on its four sides
Ammonia	Nitrogen	5	N with lone pair, bonded to three H	Dot structure showing three electron bonds and lone pair	Line structure showing N bonded to three H, lone pair indicated
Water	Oxygen	6	O with two lone pairs, bonded to two H	Dot structure showing electron bonds and lone pairs	Line structure showing O bonded to two H, lone pairs indicated
Hydrogen fluoride	Fluorine	7	F with three lone pairs, bonded to one H	Dot structure showing electron bond and lone pairs	Line structure showing F bonded to H, lone pairs indicated

Lewis structures table for common molecules

Polar Covalent Bonds, Electronegativity, and Bond Dipoles

Covalent Bonds May be Polarized

Electrons in covalent bonds may not be shared equally, resulting in bond polarization. The electronegativity of atoms determines the degree of polarization.

Electrostatic potential map for H-F

Electronegativity

Electronegativity is the ability of an atom to attract electrons. Highly electronegative atoms hold electrons tightly and tend to gain electrons, while electropositive atoms may lose electrons. Periodic trends in electronegativity are crucial for predicting reactivity.

Polar Covalent Bonds

When two atoms of different electronegativities share electrons, the bond is polar covalent. The greater the difference, the more polarized the bond.

Bond polarity and dipole direction in H-F

Electrostatic Potential Maps

Electrostatic potential maps (EPMs) visualize the spatial distribution of charge in molecules, providing a more accurate representation than simple dipole models.

Electrostatic potential map for H-F

A Selection of Dipole Moments

Bond dipole moments quantify the polarity of bonds. The direction of the dipole is toward the more electronegative atom. Greater differences in electronegativity yield larger dipole moments.

Bond	Dipole moment, D	Bond	Dipole moment, D
H–F	1.7	C–F	1.4
H–Cl	1.1	C–O	0.7
H–Br	0.8	C≡N	2.4
H–I	0.4	C=N	1.4
H–C	0.3	C=N	3.6
H–N	1.3
H–O	1.5

Table of bond dipole moments

Formal Charge

Formal charge is used to determine the electron distribution in Lewis structures. It is calculated by subtracting the group number of the neutral atom from the electron count of the covalently bound atom. Only unshared electrons and half of the bonding electrons are counted.

Formula:

Formal charges greater than ±1 are rare in organic molecules.

Structural Patterns & Formal Charge

Common patterns include oxygen with two bonds (formal charge 0), oxygen with one bond (formal charge -1), and nitrogen with four bonds (formal charge +1). Boron in borohydride (BH4-) has four bonds and a formal charge of -1.

Oxygen (group 6): Two bonds, formal charge 0; one bond, formal charge -1
Nitrogen (group 5): Four bonds, formal charge +1
Boron (group 3): Four bonds, formal charge -1

Additional info: These notes expand on the brief points in the original slides, providing definitions, examples, and context for foundational concepts in organic chemistry. All images included are directly relevant to the adjacent explanations, reinforcing key ideas visually.