Aneutere Atomic Structure and Bonding Lecture Notes

These flashcards cover the fundamental principles of atomic structure, electronegativity, Lewis structures, VSEPR, hybridization, and Molecular Orbital (MO) theory as detailed in the lecture transcript.

Atomic Number

The number of protons in an atom. If the atom is neutral, it also equals the number of electrons.

Orbital

A 3D map of electron location that indicates the energy and reactivity of an electron. Each orbital can hold a maximum of $2$ electrons.

Shell

An energy level represented by a coefficient ($1$-$7$). It corresponds to the row number on the periodic table; lower shell numbers are closer to the nucleus.

Sub-shell

The type of orbital within a shell, denoted by letters such as $s$ (which contains $1$ orbital and holds up to $2$ electrons) and $p$ (which contains $3$ orbitals and holds up to $6$ electrons).

Valence electrons

The electrons located in the highest shell number of an atom, regardless of the sub-shell letter.

Electronegativity

The ability of an atom to attract electrons through covalent bonds. It increases as you move from left to right and from down to up on the periodic table (closer to Fluorine).

Electronegativity Mnemonic

$$F > O > N \thickapprox Cl > Br > I \thickapprox S \thickapprox C > H$$

Ionic Association

An association characterized by a large electronegativity difference, usually occurring between a metal and a nonmetal.

Covalent Bonds

A chemical bond characterized by similar electronegativity values, usually between a nonmetal and another nonmetal.

Polar Covalent Bond

A bond between different elements where there is a difference in electronegativity and polarity.

Nonpolar Covalent Bond

A bond formed between two atoms of the same element.

Octet Rule Exceptions

Hydrogen ($H$) needs $2e^-$, Boron ($B$) gets $6e^-$, and Beryllium ($Be$) gets $4e^-$. Additionally, elements in the 3rd row and below can have an expanded octet.

Formal Charge (FC)

The charge of an atom calculated as: $\text{Valence } e^- - (\text{sticks} + \text{dots})$.

VSEPR

Valence Shell Electron Pair Repulsion; a theory stating that electron groups repel each other to determine the shape of a molecule.

Electron Group

A cluster of electrons in the same area, which can be a single bond, double bond, triple bond, or a lone pair.

Linear Geometry

A molecular shape with $2$ electron groups and a bond angle of $180^\text{o}$.

Trigonal Planar Geometry

A molecular shape with $3$ electron groups and a bond angle of $120^\text{o}$.

Tetrahedral Geometry

A molecular shape with $4$ electron groups and a bond angle of $109.5^\text{o}$.

Carbocation

A positively charged carbon atom that has $3$ valence electrons forming $3$ bonds (and typically no lone pairs or hydrogens).

Valence Bond Theory

A theory where a bond is defined by the overlap of two orbitals, resulting in hybridization prior to bonding.

Hybridization

The combination of an atom's own atomic orbitals ($AOs$) to form new hybrid orbitals ($HOs$) for covalent bonding.

$sp^3$ Hybridization

Occurs when $4$ electron groups are present, combining one $s$ orbital and three $p$ orbitals ($s + 3p$) to form a tetrahedral orientation.

$sp^2$ Hybridization

Occurs when $3$ electron groups are present ($s + 2p$), resulting in a trigonal planar orientation with one remaining $p$ orbital perpendicular to the hybrid orbitals.

$sp$ Hybridization

Occurs when $2$ electron groups are present ($s + 1p$), resulting in a linear orientation with two remaining $p$ orbitals.

Sigma ($\sigma$) Bond

A bond formed by the head-to-head overlap of hybrid orbitals or $s$ orbitals.

Pi ($\pi$) Bond

A bond formed by the side-to-side overlap of $p$ orbitals that are perpendicular to the sigma bond.

Molecular Orbital (MO) Theory

A theory using the linear combination of atomic orbitals where all orbitals (not just valence) change energy and shape when bonded.

Bonding MO

A lower energy molecular orbital where the majority of electron density is in the internuclear region, allowing for constructive interference.

Antibonding MO ($^*$)

A higher energy molecular orbital where electron density is outside the internuclear region, causing destructive interference and no bond.

Bond Order (BO)

A value determining the number of bonds, calculated as: \frac{1}{2} (\text{#e- in bonding MO} - \text{#e- in antibonding MO}).

HOMO

Highest Occupied Molecular Orbital; the orbital most likely to donate electrons.

LUMO

Lowest Unoccupied Molecular Orbital; the orbital where electrons are accepted.

Node

A gap in a molecular orbital where there is zero probability of finding an electron.