Homors Chem Unit 2

Dalton's Atomic Theory

All matter is made of atoms; atoms of the same element are identical; atoms combine in simple ratios; atoms aren't created or destroyed in reactions.

Heisenberg's Uncertainty Principle

It's impossible to know both the exact position and momentum of an electron at the same time.

Valence Electrons

Electrons in the outermost energy level of an atom; responsible for bonding and reactivity.

Einstein's Dual-Wave Theory of Electromagnetic Radiation

Light behaves as both a wave and a particle (photon).

Frequency and Wavelength

Inversely related

Frequency and Energy

Directly related

Photoelectric Effect

When light of sufficient frequency hits a metal surface, electrons are ejected; shows light behaves as particles (photons).

Photon

A particle of light carrying a specific amount of energy (E = hν).

How a Line-Emission Spectrum Is Produced

When excited electrons fall back to lower levels, they emit light at specific wavelengths unique to each element.

Ground State vs. Excited State

Ground state = lowest energy level; excited state = higher energy level after absorbing energy.

Law of Definite Proportions

A compound always contains the same elements in the same ratio by mass.

Law of Multiple Proportions

When elements form more than one compound, the ratios of their masses are small whole numbers.

Max Planck

Proposed that energy is quantized and emitted in small packets called quanta.

Quanta (Quantum)

A discrete amount or packet of energy.

Bohr's Model of the Atom

Electrons orbit the nucleus in fixed energy levels; energy is absorbed or emitted when electrons jump between levels.

Principal Quantum Number (n)

Indicates the main energy level or distance from the nucleus.

Orbital

A region in space where an electron is most likely found.

Names of Orbital Shapes/Sublevels

s (sphere), p (dumbbell), d (cloverleaf), f (complex).

Number of Orbitals in Each Sublevels

s = 1, p = 3, d = 5, f = 7.

Number of Electrons in Each Orbital

Each orbital holds 2 electrons with opposite spins.

Number of Orbitals in Each Energy Level

n=1 → 1; n=2 → 4; n=3 → 9; n=4 → 16.

Electron Configuration

Describes how electrons are arranged in orbitals (ex: 1s² 2s² 2p⁴).

Noble Gas Notation

Uses the previous noble gas to shorten configuration (ex: Na → [Ne] 3s¹).

Orbital Notation

Uses boxes and arrows to show electrons and their spins (↑↓).

Hund's Rule

Electrons fill orbitals singly before pairing up.

Lewis Dot Structures

Use dots around an element's symbol to represent valence electrons.