Principles of Chemistry: Weeks 1-4 Overview

Molecular Shape

Geometric arrangement of atoms within a molecule.

Atomic Structure

Arrangement of protons, neutrons, and electrons in atoms.

Quantum Theory

Theory describing the behavior of matter and energy at atomic levels.

Wave-Particle Duality

Light exhibits both wave and particle characteristics.

Phase

Starting point of a wave relative to one wavelength.

Speed of Light (c)

Constant speed of electromagnetic radiation in vacuum, 2.998×10^8 m/s.

Electromagnetic Spectrum

Range of all types of electromagnetic radiation.

Photoelectric Effect

Emission of electrons from a material when exposed to light.

Threshold Frequency (ν0)

Minimum frequency required to eject electrons from a metal.

Planck's Constant (h)

Constant relating energy of photons to frequency, 6.626×10^-34 J·s.

Binding Energy (Φ)

Minimum energy needed to remove an electron from an atom.

Photons

Massless packets of energy that make up light.

Lewis Structures

Diagrams showing bonding between atoms in a molecule.

Valence Bond Theory

Theory explaining how atomic orbitals combine to form bonds.

Ionic Bonding

Electrostatic attraction between oppositely charged ions.

Covalent Bonding

Sharing of electron pairs between atoms.

Periodic Table

Tabular arrangement of elements based on atomic number.

Niels Bohr

Developed quantized model of atomic structure.

Quantized Orbits

Electrons occupy specific circular orbits only.

Centrifugal Force

Balances attraction to positive nucleus in orbits.

Photon Absorption

Electrons transition to higher orbits by absorbing photons.

Photon Emission

Electrons transition to lower orbits by emitting photons.

Rydberg Constant (RH)

Value for hydrogen atom, 2.179×10−18 J.

Principal Quantum Number (n)

Indicates energy level of electron orbit.

Emission Spectrum

Specific wavelengths emitted by hydrogen atom.

Lyman Series

Transitions to n=1, occur in ultraviolet region.

Balmer Series

Transitions to n=2, occur in visible light.

Paschen Series

Transitions to n=3, occur in infrared region.

de Broglie Equation

Relates wavelength to momentum of particles.

Wave Properties of Electrons

Electrons exhibit diffraction similar to light.

Heisenberg's Uncertainty Principle

Can't simultaneously know position and momentum of electrons.

Schrödinger Wave Equation

Mathematical description of electron wavefunctions.

Atomic Orbitals (AOs)

Quantized wavefunctions describing electron distributions.

Orbital Quantum Numbers

Describes size and energy of atomic orbitals.

Secondary Quantum Number (l)

Indicates shape of atomic orbital.

Magnetic Quantum Number (ml)

Describes orientation of atomic orbitals.

Spin Quantum Number (ms)

Describes electron spin direction, +½ or −½.

Pauli Exclusion Principle

No two electrons can have identical quantum numbers.

Max Electrons per Shell

Maximum electrons per shell follows 2n² rule.

Shells and Sub-shells

Groups of orbitals with same n and l values.

Wavefunction (ψ)

Mathematical function representing electron's wave properties.

Hamiltonian Operator (Ĥ)

Operator used in Schrödinger's wave equation.

Electron Diffraction

Confirmed wave nature of electrons in 1927.

Energy of nth Orbit (En)

Calculated using En = -RH/n².

Stoichiometry

Calculation of reactants and products in chemical reactions.

Chemical Reactions

Processes where substances transform into new substances.

Chemical Bonding

Attractive forces holding atoms together in compounds.

Wavelength (λ)

Distance between successive peaks of a wave.

Frequency (ν)

Number of wave cycles passing a point per second.

Amplitude

Height of a wave, related to its intensity.