Organic Chemistry Chapter 1: Electrons, Bonds, and Molecular Properties

Vocabulary flashcards reviewing fundamental concepts from Chapter 1 of Organic Chemistry, including atomic structure, formal charge, electronegativity, orbital hybridization, VSEPR theory, molecular polarity, and intermolecular forces.

Organic Chemistry

The study of carbon-containing molecules and their reactions.

Constitutional Isomers

Compounds that have the same molecular formula but different structures because their atoms are bonded together in different ways.

Tetravalent

Describing an atom, such as carbon, that generally forms four bonds.

Trivalent

Describing an atom, such as nitrogen, that generally forms three bonds.

Divalent

Describing an atom, such as oxygen, that generally forms two bonds.

Monovalent

Describing an atom, such as hydrogen or a halogen ($F$, $Cl$, $Br$, or $I$), that generally forms one bond.

Valence Electrons

The electrons located in the outermost shell of an atom, which are primarily involved in chemical bonding.

Formal Charge

A charge calculated by comparing the number of valence electrons an atom owns in a bonding environment to the number of valence electrons it needs to be neutral.

Electronegativity

A measure of how strongly an atom attracts shared electrons within a chemical bond.

Covalent Bond

A chemical bond formed when electrons are shared between two atoms with an electronegativity difference of less than $0.5$.

Polar Covalent Bond

A chemical bond in which electrons are shared unequally between two atoms with an electronegativity difference between $0.5$ and $1.7$.

Ionic Bond

A bond between two atoms differing in electronegativity by more than $1.7$, where electrons are not shared and the more electronegative atom owns the electrons.

Electron Density

A term used to refer to the probability of finding an electron within a specific region of space around an atom or molecule.

Node

A location or plane in an orbital or wavefunction where $\psi = 0$, meaning there is zero probability of finding an electron.

Degenerate Orbitals

Atomic or molecular orbitals that possess equal energy, such as the three $2p$ orbitals.

Sigma Bond (\sigma Bond)

A covalent bond formed by the direct, head-on overlapping of atomic orbitals along the internuclear axis.

Molecular Orbital (MO) Theory

A bonding theory in which atomic orbital wavefunctions mathematically combine via constructive and destructive interference to form molecular orbitals extending over the entire molecule.

Antibonding MO

A molecular orbital formed by destructive interference of atomic orbitals, characterized by a node between the nuclei and a higher energy state.

sp^3 Hybridization

The mixing of one $s$ orbital and three $p$ orbitals to form four equal-energy hybrid orbitals with $25\%$ s-character and $75\%$ p-character.

sp^2 Hybridization

The mixing of one $s$ orbital and two $p$ orbitals to form three degenerate hybrid orbitals with $33\%$ s-character and $67\%$ p-character, leaving one unhybridized $p$ orbital.

Pi Bond (\pi Bond)

A covalent bond formed by side-by-side overlap of unhybridized $p$ orbitals, with electron density spread out above and below the plane of the molecule.

sp Hybridization

The mixing of one $s$ orbital and one $p$ orbital to form two degenerate hybrid orbitals, leaving two unhybridized $p$ orbitals available for $\pi$ bonding.

VSEPR Theory

Valence Shell Electron Pair Repulsion theory; a model predicting molecular geometry based on the principle that valence electron pairs (shared and lone pairs) repel each other.

Steric Number

The sum of the number of $\sigma$ bonds and lone pairs around a central atom, used to predict hybridization and electron group geometry.

Induction

The shifting of electrons within an orbital or bond caused by differences in electronegativity, creating partial charge separation.

Dipole Moment

A quantitative measure of polarity, calculated as the product of the partial charge magnitude and the distance separating the $\delta+$ and $\delta-$ charges.

Hydrogen Bonding

An exceptionally strong type of dipole-dipole attraction occurring between a hydrogen atom bonded to an electronegative atom ($N$, $O$, or $F$) and a lone pair on another electronegative atom.

Protic Solvent

A solvent containing hydrogen atoms bonded to highly electronegative atoms (such as $N$, $O$, or $F$), making it capable of engaging in hydrogen bonding.

Aprotic Solvent

A solvent that lacks hydrogen atoms bonded to highly electronegative atoms and therefore cannot participate in hydrogen bonding.

London Dispersion Forces

Fleeting, attractive intermolecular forces present between all molecules (including nonpolar ones) resulting from induced, transient dipole moments created by constant random electron motion.