Chapter 1: Structure and Bonding (Vocabulary)

Vocabulary flashcards covering key concepts from Chapter 1: Structure and Bonding of Organic Chemistry (OpenStax).

Nucleus

The dense, positively charged center of an atom that contains most of the atom’s mass and is made up of protons and neutrons.

Electron

Negatively charged particle that orbits the nucleus and participates in bonding by occupying atomic orbitals.

Electron density

Probability distribution of electrons in space; electron-density surfaces show regions of high electron density.

Orbital

Region in an atom where there is a high probability of finding an electron; labeled by shape (s, p, d, etc.).

s orbital

Spherically shaped orbital that can hold up to 2 electrons.

p orbital

Dumbbell-shaped orbital; there are three mutually perpendicular p orbitals in a subshell.

d orbital

Five cloverleaf-shaped orbitals available in a d subshell.

Node (orbital)

Region within an orbital where the probability of finding an electron is zero.

Aufbau principle

Rule that electrons fill the lowest-energy orbitals first (e.g., 1s before 2s, then 2p, etc.).

Electron configuration

The arrangement of electrons in an atom’s orbitals according to the Aufbau principle.

Shell

A principal energy level that can contain subshells (s, p, d, f) and electrons.

Tetrahedral carbon

Carbon atom bonded to four groups in a tetrahedral geometry, with bond angles near 109.5°.

Valence electrons

Electrons in the outermost shell that participate in bonding.

Covalent bond

A bond formed by sharing a pair of electrons between atoms.

Coordinate covalent bond

a type of covalent bond in which one atom gives both electrons to share with another atom.

Ionic bond

A chemical bond is formed when one atom transfers one or more of its electrons to another atom, creating oppositely charged ions that attract each other.

Lewis structure (a.k.a Electron-dot structures)

Electron-dot representation showing valence electrons around atoms.

Kekulé structure

Line-bond representation of molecules where bonds are shown as lines.

Octet rule

Tendency of main-group elements to have eight electrons in their valence shell in stable structures.

Lone pair

A nonbonding pair of electrons on an atom.

Bond length

Distance between the nuclei of two bonded atoms at minimum energy.

Bond energy

Energy released when a bond forms (or required to break a bond); a measure of bond strength.

Sigma (σ) bond

Bond formed by end-to-end overlap of orbitals along the bond axis; cylindrically symmetric about the bond.

Pi (π) bond

Bond formed by sideways overlap of p orbitals; lies above and below the bond plane.

sp3 hybridization

Mixing of one s and three p orbitals to form four equivalent sp3 hybrids, arranged toward tetrahedral corners; about 109.5°.

sp2 hybridization

Mixing of one s and two p orbitals to form three sp2 hybrids in a plane (120°) with one unhybridized p orbital perpendicular.

sp hybridization

Mixing of one s and one p orbital to form two sp hybrids 180° apart, with two remaining p orbitals; leads to linear geometry.

Ethane (C₂H₆) structure

Molecule in which the C–C bond is formed by σ overlap of two sp3 hybrid orbitals.

Ethylene (C₂H₄)

Molecule with a C=C double bond; σ bond from sp2 overlap and a π bond from unhybridized p orbitals.

Acetylene (C_₂H₂)

Molecule with a C≡C triple bond; one σ bond from sp hybrids and two π bonds from unhybridized p orbitals.

Formaldehyde (CH2O)

Carbonyl compound (CH2O); carbon is sp2-hybridized with a C=O double bond.

Molecular orbital theory

Bonding theory where electrons occupy molecular orbitals that extend over the entire molecule, including bonding and antibonding MOs.

Bonding molecular orbital

Molecular orbital that, when occupied, stabilizes the molecule.

Antibonding molecular orbital

Molecular orbital that, when occupied, destabilizes the molecule.

H₂ molecular orbitals

In H2, a bonding MO is filled and the antibonding MO is empty, leading to a stable molecule.

Condensed structure

A shorthand representation showing connectivity without drawing all bonds (e.g., CH3CH2CH3).

Line-bond rule (interpretation of line drawings)

In line-bond structures, an end of a line represents CH3, a two-way intersection CH2, a three-way intersection CH, and a four-way intersection carbon with no hydrogens.

Double bond

A bond consisting of one σ and one π component; involves sharing two pairs of electrons.

Planar geometry (sp2)

Three sp2 hybrids lie in a plane (120° apart) with one unhybridized p orbital perpendicular, enabling π bonding.

Linear geometry (sp)

Two sp hybrids are 180° apart with two remaining unhybridized p orbitals, giving a linear arrangement.

Atomic Number

The number of protons in the nucleus of an atom, which determines the element's identity and its position in the periodic table.

Mass Number

The total number of protons and neutrons in an atom's nucleus, which gives the atomic mass of an element.

Isotopes

Atoms of the same element with the same number of protons but different numbers of neutrons, resulting in different mass numbers.

Atomic Weight

The weighted average mass of an element's isotopes, measured in atomic mass units (amu), reflecting both the mass and relative abundance of each isotope.

Eletron Shells

The specific regions around an atom's nucleus where electrons are likely to be found, categorized by different energy levels.

Ground-State Electron Configuration

Shows how electrons are arranged in an atom’s orbitals when the atom is in its lowest energy state (most stable condition).

Pauli Exclusion Principle

A fundamental principle in quantum mechanics stating that no two electrons in the same atom can have the same set of quantum numbers, which means only two electrons can fit in one orbital, and they must have opposite spins.

Hund’s Rule

When electrons fill orbitals of the same energy (like in a p, d, or f subshell), they will fill each orbital singly first with the same spin before pairing up.

Molecule

A group of two or more atoms that are chemically bonded together and act as a single unit.

Line-bond structures

A type of chemical structure representation that shows bonds between atoms as lines, with each line representing a pair of shared electrons, and often omitting lone pairs.

Valence Bond (VB) Theory

This theory explains how atoms form bonds by overlapping their atomic orbitals, with shared electrons creating a covalent bond between them.

Polar bonds

Covalent bonds between atoms with differing electronegativities, resulting in a molecule with a dipole moment.