SCH3U - Unit 1 - Definitions

Hydrate

A compound that has associated with water

Polyatomic

A group of atoms with an atomic charge

Oxyanion

A polyatomic ion with oxygen

Proton

Positively charged subatomic particle

Atomic Number

The number of protons in an atom

Neutron

A subatomic particle that does not carry a charge, but contributes to an atom’s mass

Electrons

Negatively charges subatomic particle

Isotope

Atoms of an element with the same atomic number but different atomic masses

Half Life

The time it takes for half of an element to decay

Radioisotopes

Isotopes that emit energy in the form of alpha, beta, and gamma rays

Principle Quantum Number “n”

Whole numbers starting at 1 that represent increasing energy levels

Second Quantum Number “l”

The angular momentum number that gives us the shape of an atomic orbital

Third Quantum Number “m_l”

The orientation of an atom

Fourth Quantum Number “m_s”

The two possible electrons spins, either +1/2 or -1/2

Electron Density

The density around the nucleus that electrons occupy

Hund’s Rule

When orbitals of equal energy are being filled, the most stable configuration is the one with the maximum number of unpaired electrons with the same spin

Pauli’s Exclusion Principle

No 2 electrons in the same atom can be described by the same set of quantum numbers

Aufbau Principle

Order of filling electrons with the lowest energy first

Periodic Law

When elements are arranged in order of increasing atomic number and therefore, mass, certain sets of properties recur periodically, defined as trends

Effective Nuclear Charge (Z_eff)

The net positive charge the outer electrons “feel”

Coulomb’s Law

F = (Q₁Q₂)/d² - As the diameter increases, the force of attraction decreases

Atomic/Ionic Radii

The size of the neutral atom

Ionic Radius

The size of an ion

Ionization Energy

A measure of how much energy it takes to remove a valence electron from an atom in the gaseous state

X(g) + energy → X⁺(g) + e^-

Electron Affinity

The energy released when an electron is added

X⁺(g) + e^- → X(g) + energy

Electronegativity

The measure of how strongly an atom pulls electrons towards itself

Ionic Bond

The electrostatic attractive force between the oppositely charged ions produced when a metal atom transfers one or more electrons to a non-metal atom

Formal Charge

Apparent charges on certain atoms in a Lewis structure that arise when atoms have not contributed equal numbers

Resonance Structure

The possible diagrams to represent the location of delocalized electrons

Dipole

The opposite regions of polarity created when there is a charge differential in a molecule

Electron Group

1, 2, or 3 bonds, or lone pairs, attached to the central atom

Electron Geometry

The arrangement of electrons, as described by the bond angle

Molecular Geometry

The arrangement of the molecule, as described by its shap

Melting Point

The temperature at which a compound changes from a solid to a liquid

Boiling Point

The temperature at which a compound changes from a liquid to a gas

Solubility

Whether a compound dissolves or not in water

Conductivity

The ability to allow electric current to flow (the formation of ions in a solution)

Vapour Pressure/Volatility

How easily a compound will vaporize

Intermolecular Forces

The attraction that keeps molecules together

London Dispersion Forces

The temporary dipole created by electron movement, found in all molecules

Dipole-Dipole Forces

The forces between molecules with permanent dipoles, where the negative poles attract to the positive poles

Hydrogen Bonding

A type of dipole-dipole force, which results from a hydrogen atom being bonded to a highly electronegative atom such as nitrogen, oxygen, or fluorine

Ion-Dipole Force

The attraction between an ion and a polar molecule that causes them to dissolve