Chemistry

The Octet Rule

The tendency of an atom to have 8 valence electrons in the valence shell in order for atoms to become stable.

Ionic Bonding

When a metal transfers one or more electrons to a non-metal; the metal loses electrons to become a positive ion (cation), and the non-metal gains those electrons to become a negative ion (anion).

Metallic Bonding

Takes place between metal atoms where they pool their valence electrons together, releasing them into a shared, mobile "sea of electrons" that surrounds a grid of positive metal ions.

Covalent Bonding

A chemical link formed when two nonmetal atoms share one or more pairs of electrons; they overlap outer shells so both feel like they have a full complement.

Covalent Prefix 1

Mono-

Covalent Prefix 2

Di-

Covalent Prefix 3

Tri-

Covalent Prefix 4

Tetra-

Covalent Prefix 5

Penta-

Covalent Prefix 6

Hexa-

Covalent Prefix 7

Hepta-

Covalent Prefix 8

Octa-

Covalent Prefix 9

Nona-

Covalent Prefix 10

Deca-

Molecular Formula

A chemical formula that shows the exact number and type of atoms in a molecule.

Electron Dot Diagram

A 2D visual representation that uses dots to represent valence electrons.

Lewis Diagram

A 2D visual representation that uses dots for lone pairs and lines for bonding pairs of electrons.

Structural Formula

A 2D visual representation that uses lines to represent a pair of bonded electrons.

VSEPR Theory (Definition)

Valence Shell Electron Pair Repulsion; describes the way atoms are arranged in a molecule, determined by repulsion between valence electron pairs.

Electronegativity

An atom's ability to attract shared electrons in a covalent bond toward itself; increases moving left-to-right across a period, decreases top-to-bottom down a group.

Nonpolar Covalent Bond (Electronegativity)

Electronegativity difference (ΔEN) of 0 to 0.4; shared electrons are equally distributed (e.g., C-H, O2).

Polar Covalent Bond (Electronegativity)

Electronegativity difference (ΔEN) of 0.4 to 2; shared electrons are pulled closer to the more electronegative atom, creating partial positive (δ+) and partial negative (δ-) charges.

Ionic Bond (Electronegativity)

Electronegativity difference (ΔEN) of 2+; electrons are completely transferred from a metal to a nonmetal.

Bond Polarity vs Molecular Polarity

Bond polarity is the uneven sharing of electrons between two individual atoms; molecular polarity is the overall distribution of charge across the entire 3D shape of a molecule.

Nonpolar Molecule

A molecule with a symmetrical VSEPR shape where bond dipoles cancel out.

Polar Molecule

A molecule with an asymmetrical VSEPR shape resulting in an uneven charge distribution.

VSEPR Shape: 2 Bonded, 0 Lone

Linear

VSEPR Shape: 3 Bonded, 0 Lone

Trigonal Planar

VSEPR Shape: 2 Bonded, 1 Lone

Bent

VSEPR Shape: 4 Bonded, 0 Lone

Tetrahedral

VSEPR Shape: 3 Bonded, 1 Lone

Trigonal Pyramidal

VSEPR Shape: 2 Bonded, 2 Lone

Bent

VSEPR Shape: 5 Bonded, 0 Lone

Trigonal Bipyramidal

VSEPR Shape: 4 Bonded, 1 Lone

Seesaw

VSEPR Shape: 3 Bonded, 2 Lone

T-shaped

VSEPR Shape: 2 Bonded, 3 Lone

Linear

VSEPR Shape: 6 Bonded, 0 Lone

Octahedral

VSEPR Shape: 5 Bonded, 1 Lone

Square Pyramidal

VSEPR Shape: 4 Bonded, 2 Lone

Square Planar

Mass Number

Number of protons + neutrons inside the nucleus. Protons identify the element; protons + neutrons = mass number (e.g., Carbon 6p + 6n = 12C; Bromine 35p + 44n = 79Br, mass of 79u).

Isotopes

Two atoms with the same number of protons but a different number of neutrons (e.g., 6p+6n, 6p+7n, 6p+8n are isotopes of carbon; 7p+7n, 7p+8n are isotopes of nitrogen).

Mole & Avogadro's Number

The unit for measuring the amount of a substance; Avogadro’s number (N_A ≈ 6.022 × 10^23 mol⁻¹) defines how many individual particles are in one mole.

Molar Mass (MR)

Mass of one mole (6.022 x 10^23 molecules) measured in grams per mole (g/mol). Example: 1 mol of 12C weighs 12g (MR = 12 g/mol).

Calculated Relative Molar Mass Example

Natural substances are isotope mixes. Copper: 63Cu (70%) & 65Cu (30%) -> MR(Cu) = (0.7 × 63) + (0.3 × 65) = 63.6 g/mol.

Molecular Mass

The mass of a single molecule, measured in atomic mass units (amu); calculated by adding the atomic masses of all atoms in that molecule.

Formula: Molar Mass (M)

Given chemical formula & periodic table -> Sum of (atomic mass x subscript) for each element (Units: g/mol or amu).

Formula: Moles of Molecule (n)

Given Mass (m) & Molar Mass (M) -> n = m / M (Units: mol).

Formula: Mass (m)

Given Moles (n) & Molar Mass (M) -> m = n × M (Units: g).

Formula: Number of Molecules

Given Moles of Molecule (n) -> Molecules = n × (6.022 × 10^23) (Units: molecules).

Formula: Moles of a Specific Atom

Given Moles of Molecule (n) & Formula -> Moles of Atom = n × (Subscript of that atom) (Units: mol).

Formula: Number of Specific Atoms

Given Moles of Molecule (n) & Formula -> Atoms = n × (Subscript) × (6.022 × 10^23) (Units: atoms).

Formula: Moles from Atom Count

Given Total number of atoms (N) -> n = N / (6.022 × 10^23) (Units: mol).

Formula: Moles from Molecule Count

Given Total number of molecules (N) -> n = N / (6.022 × 10^23) (Units: mol).

Combustion Reactions

Commonly known as burning or exploding; exothermic reactions that release heat or light energy. If there isn't enough oxygen, incomplete combustion occurs, forming carbon monoxide instead of carbon dioxide.

Combustion Energy Formula

Energy = mass x heat of combustion

Chemical Energy (Bond vs Kinetic)

Bond energy is energy stored in bonds (needed to break bonds, released when forming bonds); kinetic energy is due to movement. Occurs when attraction/repulsion between electrons and protons is balanced.

System vs Surroundings

System is the chemical reaction itself; surroundings are the surrounding environment.

Endothermic Reactions

Reactions that absorb heat energy from surroundings; require energy to happen, so products have more energy than reactants (e.g., breaking an icepack releases ammonium nitrate dissolving in water, absorbing heat).

Exothermic Reactions

Reactions that produce heat energy; extra energy is converted to heat and released into surroundings, so products have less energy than reactants (e.g., iron reacting with oxygen in a heat pack to form iron oxide).