Chapter1

This set of flashcards covers the fundamental vocabulary from Chapter One of Organic Chemistry, including atomic structure, bonding, Lewis structures, resonance, and molecular geometry.

Protons

Positively charged particles located within the nucleus of an atom.

Neutrons

Uncharged particles located within the nucleus of an atom.

Electrons

Negatively charged particles that compose the electron cloud and can be described as both particles and waves.

Isotopes

Atoms of the same element that have a different number of neutrons, resulting in different masses.

Atomic weight

A weighted average of the mass of all naturally occurring isotopes of an element, such as $12.011\,g/mol$ for carbon.

s orbital

An orbital characterized by a sphere of electron density that is lower in energy than other orbitals in the same shell.

p orbital

A dumbbell-shaped orbital that contains a node at the nucleus where there is zero probability of finding an electron.

Ground state

The lowest energy arrangement of electrons for an atom.

Ionic bonds

Bonds that result from the transfer of electrons from one element to another, typically occurring between elements on the far-left and far-right of the periodic table.

Covalent bonds

Bonds that result from the sharing of electrons between two nuclei.

Tetravalent

The property of an atom, specifically carbon, to form four bonds.

Lone pairs

Nonbonding or unshared electrons that contribute to an atom's octet.

Formal charge

The charge assigned to individual atoms in a Lewis structure, calculated as the number of valence electrons minus the electrons "owned" by the atom.

Isomers

Different molecules that share the same molecular formula.

Constitutional isomers

Isomers that differ in the arrangement of their atoms, such as ethanol and dimethyl ether.

Resonance structures

Two or more valid Lewis structures used to represent a single molecule or ion that differs only in the arrangement of electrons.

Resonance hybrid

The true structure of a molecule that is a composite or weighted average of all its contributing resonance forms.

Delocalization

The spreading out of electron pairs over two or more atoms, which increases the stability of a chemical species.

Curved arrow notation

A convention used to show how electron positions differ between resonance forms; the tail starts at an electron pair and the head points to where it "moves."

VSEPR theory

Valence Shell Electron Pair Repulsion theory, which posits that the most stable geometry keeps groups (atoms and lone pairs) as far apart as possible.

Linear geometry

The shape of an atom surrounded by two groups, resulting in a bond angle of $180^{\circ}$.

Trigonal planar geometry

The shape of an atom surrounded by three groups, resulting in bond angles of $120^{\circ}$.

Tetrahedral geometry

The shape of an atom surrounded by four groups, resulting in bond angles of $109.5^{\circ}$.

Condensed structures

A representation of organic molecules where all atoms are drawn but two-electron bond lines are generally omitted, using parentheses for identical groups.

Skeletal structures

A representation where a carbon atom is assumed at every junction or end of a line, and hydrogens on carbons are omitted to satisfy carbon's tetravalency.

Heteroatoms

Atoms in organic molecules that are not carbon or hydrogen, such as N, O, and halogens.

Sigma ($\sigma$) bond

A cylindrically symmetrical bond that concentrates electron density along the axis connecting two nuclei.

Hybridization

The combination of two or more atomic orbitals to form new hybrid orbitals with the same shape and energy.

sp3 hybrid orbitals

Four hybrid orbitals formed from one $2s$ and three $2p$ orbitals, oriented toward the corners of a tetrahedron.

sp2 hybrid orbitals

Three hybrid orbitals formed from one $2s$ and two $2p$ orbitals, leaving one unhybridized $p$ orbital to form a $\pi$ bond.

Pi ($\pi$) bond

A bond created by the side-by-side overlap of two unhybridized $p$ orbitals.

sp hybrid orbitals

Two hybrid orbitals formed from one $2s$ and one $2p$ orbital, oriented $180^{\circ}$ apart.

Electronegativity

A measure of an atom's attraction for electrons within a chemical bond.

Polar covalent bond

A bond between atoms of different electronegativity values characterized by unequal sharing of electrons and a dipole.

Dipole

A separation of charge within a bond or molecule, often indicated by the symbols $\delta+$ and $\delta-$.