1. States of Matter

Types of States

Matter exists in different physical forms known as states of matter. Each state has distinct properties based on particle arrangement and energy.

• Solid: Definite shape and volume. Particles are closely packed and vibrate in place. Example: Ice.

• Liquid: Definite volume but no definite shape. Particles are less tightly packed and can move past each other. Example: Water.

• Gas: No definite shape or volume. Particles move freely and are far apart. Example: Oxygen.

• Plasma: Ionized gas with free electrons and ions. Example: Lightning, neon signs.

• Bose-Einstein Condensate: State at temperatures near absolute zero where particles behave as a single quantum entity. Example: Laboratory-created condensates.

2. Changes of State

Phase Transitions

Changes between states of matter involve energy transfer and are called phase transitions.

• Fusion (Melting): Solid to liquid.

• Solidification (Freezing): Liquid to solid.

• Evaporation: Liquid to gas.

• Condensation: Gas to liquid.

• Sublimation: Solid to gas.

• Deposition: Gas to solid.

3. Methods for Separating Mixtures

Separation Techniques

Mixtures can be separated into their components using physical methods:

• Filtration: Separates solids from liquids using a porous barrier.

• Distillation: Separates substances based on differences in boiling points.

• Chromatography: Separates components based on movement through a medium.

• Evaporation: Removes a liquid from a solution to leave a solid residue.

• Decantation: Separates liquid from solids by pouring off the liquid.

4. Classification of Matter

Types of Substances

Matter is classified based on composition:

• Pure Substances: Elements and compounds with fixed composition.

• Mixtures: Physical combinations of two or more substances. Can be homogeneous (uniform) or heterogeneous (non-uniform).

5. Physical and Chemical Properties

Types of Properties

• Physical Properties: Observed without changing the substance's identity. Examples: Color, density, melting point.

• Chemical Properties: Describe a substance's ability to undergo chemical changes. Examples: Reactivity, flammability.

Physical and Chemical Changes

• Physical Change: No new substance is formed. Example: Melting ice.

• Chemical Change: New substances are formed. Example: Rusting iron.

6. Subatomic Particles and Atomic Structure

Subatomic Particles

• Proton (p): Positive charge, located in the nucleus.

• Neutron (n): No charge, located in the nucleus.

• Electron (e): Negative charge, located in electron cloud around nucleus.

7. Atomic Models

Development of Atomic Theory

• Dalton's Model: Atoms are indivisible particles.

• Thomson's Model: "Plum pudding" model; electrons in a positive sphere.

• Rutherford's Model: Small, dense, positive nucleus; electrons orbit nucleus.

• Bohr's Model: Electrons in fixed orbits with quantized energy levels.

• Quantum Mechanical Model: Electrons in orbitals; probability-based locations.

8. Atomic Number and Mass Number

Definitions

• Atomic Number (Z): Number of protons in the nucleus; defines the element.

• Mass Number (A): Total number of protons and neutrons in the nucleus.

Isotopes: Atoms of the same element with different numbers of neutrons.

9. Identifying Isotopes and Atomic Structure

Isotope Notation

• Isotopes are written as , where X is the element symbol, A is mass number, Z is atomic number.

• Example: is carbon-14.

10. Organization of the Periodic Table

Structure and Groups

• 7 periods (rows), 18 groups (columns).

• Group Classifications: Alkali metals, alkaline earth metals, halogens, noble gases, transition metals, lanthanides, actinides.

• Metals: Left and center; Nonmetals: Right; Metalloids: Stair-step region.

11. Periodic Trends

Trends Across Periods and Groups

• Atomic Radius: Decreases across a period, increases down a group.

• Ionization Energy: Increases across a period, decreases down a group.

• Electronegativity: Increases across a period, decreases down a group.

12. Electron Configuration Rules

Filling Order and Principles

• Aufbau Principle: Electrons fill orbitals of lowest energy first.

• Pauli Exclusion Principle: No two electrons in an atom can have the same set of quantum numbers.

• Hund's Rule: Electrons occupy degenerate orbitals singly before pairing.

Example: The electron configuration for sodium (Na, atomic number 11) is:

13. Electron Configurations and Valence Electrons

Identifying Valence Electrons

• Valence electrons are those in the outermost shell; they determine chemical reactivity.

• For representative elements, the group number indicates the number of valence electrons.

Identifying an Atom from Electron Configuration

• Count the total number of electrons in the configuration.

• This number equals the atomic number and identifies the element.

• Example: has 11 electrons, which is sodium (Na).