Chemistry Chapters 3-5 Exam Review

Flashcards covering key vocabulary, concepts, and principles from Chapters 3, 4, and 5 related to atomic structure, bonding, and molecular geometry for exam preparation.

Electromagnetic Radiation

Concepts of energy (E), wavelength, and frequency and their relationship to different regions of the electromagnetic spectrum.

Quantum Numbers

Numbers assigned to orbitals to describe their sizes, energies, and orientations within an atom.

Aufbau Principle

A principle used to determine the electron configuration of an atom or ion by filling orbitals from lowest energy to highest.

Hund's Rule

A rule stating that electrons will occupy separate orbitals within a subshell with parallel spins before pairing up.

Electronic Configurations

The distribution of electrons of an atom or molecule in atomic or molecular orbitals.

Orbital Diagrams

Visual representations of electron configurations, showing electrons in individual orbitals with their spins.

Ionization Energy

The energy required to remove an electron from a gaseous atom or ion.

Electron Affinity

The energy change that occurs when an electron is added to a gaseous atom to form a negative ion.

Lewis Structures

Diagrams that show the bonding between atoms of a molecule and the lone pairs of electrons that may exist in the molecule.

Resonance Structures

Two or more Lewis structures that collectively describe the delocalized bonding in a molecule or polyatomic ion, differing only in the placement of electrons.

Formal Charge

The hypothetical charge an atom would have if all bonds were perfectly covalent and electrons were shared equally.

Extended Octet

Molecules or ions where the central atom has more than eight valence electrons, typically seen with elements in Period 3 and beyond.

Bond Order

A measure of the number of electron pairs shared between two atoms (e.g., single=1, double=2, triple=3), relating to bond energy and bond length.

Odd Electron Species

Molecules or ions that contain an odd number of valence electrons, also known as radical species.

Ionic Bonds

Chemical bonds formed by the electrostatic attraction between oppositely charged ions, formed by the transfer of electrons.

Covalent Bonds

Chemical bonds formed by the sharing of electron pairs between atoms.

Metallic Bonds

Chemical bonds formed by the electrostatic attraction between positively charged metal ions and a 'sea' of delocalized electrons.

Ionic Lattice Energy

The energy required to separate one mole of an ionic compound into its gaseous ions, indicating the strength of ion-ion attractions.

Bond Polarity

A measure of the unequal sharing of electrons in a covalent bond, determined by the difference in electronegativity between the bonded atoms.

Electronegativity (X)

A measure of the tendency of an atom to attract a bonding pair of electrons.

Valence Shell Electron Pair Repulsion (VSEPR) Theory

A model used to predict the 3D molecular geometry of a molecule based on minimizing the repulsion between electron pairs in the valence shell of its central atom.

Molecular Geometry

The 3D arrangement of atoms in a molecule, determined by the VSEPR theory.

Steric Number

The sum of the number of atoms bonded to the central atom and the number of lone pairs on the central atom, used in VSEPR theory.

Electron-Pair Geometry

The 3D arrangement of all electron groups (bonding and lone pairs) around a central atom.

Bond Dipoles

A separation of charge within a bond creating a partial positive and partial negative end, represented by an arrow pointing towards the more electronegative atom.

Molecular Polarity

Describes whether a molecule has a net dipole moment, depending on the presence of polar bonds and the molecule's overall 3D shape.

Green House Effect

The process by which certain gases in the Earth's atmosphere trap heat, caused by molecules absorbing infra-red (IR) radiation.