Introductory Definitions and Classification of Matter

Matter

Matter is defined as anything that has mass and occupies volume. All physical objects and substances are composed of matter.

• Mass: The quantity of matter in an object, typically measured in grams (g) or kilograms (kg).

• Volume: The amount of space an object occupies, measured in units such as liters (L), cubic decimeters (dm3), milliliters (mL), or cubic centimeters (cm3).

State of matter: Matter exists primarily in three states: solid, liquid, and gas.

• Composition: The types of substances present in matter, e.g., copper (element), water (compound).

• Atom: The basic building block of matter, consisting of protons, neutrons, and electrons.

Elements and Compounds

Elements

Elements are pure substances that contain only one type of atom. They cannot be broken down into simpler substances by chemical means.

• Monatomic elements: Consist of single, unbonded atoms (e.g., noble gases like helium).

• Polyatomic elements: Consist of molecules made of several identical atoms bonded together.

• Diatomic elements: Molecules composed of two atoms of the same element (e.g., O2, N2).

• Allotropes: Different forms of the same element in the same state of matter (e.g., oxygen as O2 and O3; carbon as diamond and graphite).

Compounds

Compounds are substances that contain two or more different types of atoms chemically bonded together. Their properties differ from those of their constituent elements.

• Examples: NaCl (sodium chloride), H2O (water).

• Atoms can only be altered by chemical means; molecules can be altered by physical means.

Example Equations:

• Dehydration of sugar:

• Electrolysis of water:

Classifying Matter

Pure Substances vs. Mixtures

• Pure substances: Have a fixed composition and distinct properties. Includes elements and compounds.

• Mixtures: Combinations of two or more substances not chemically bonded. Can be separated by physical means.

Types of Mixtures

• Homogeneous mixture (solution): Uniform composition throughout; particles are microscopic and evenly mixed (e.g., saltwater).

• Heterogeneous mixture: Non-uniform composition; different properties in different parts of the sample (e.g., sand in water).

• Alloy: Homogeneous mixture of metals (e.g., brass, bronze).

• Suspension: Heterogeneous mixture that settles over time (e.g., muddy water).

Contrast Example: 24K gold (pure substance) vs. 14K gold (mixture/alloy).

Chart for Classifying Matter

Separating Mixtures

Physical Methods

Mixtures can be separated by physical means, which do not change the chemical identity of the substances.

• Sorting: Separating based on physical characteristics.

• Filtration: Separating solids from liquids using a filter.

• Magnet: Using magnetic properties to separate substances.

• Chromatography: Separating based on movement through a medium.

• Density: Separating based on differences in density.

• Distillation: Separating based on differences in boiling points.

Density and Density Calculations

Density

Density is a physical property defined as mass per unit volume.

• Formula:

• Units: g/cm3 for solids, g/mL for fluids.

• The density of a liquid or solid is nearly constant, regardless of sample size.

• Density of water: Approximately 1.00 g/mL at room temperature.

Density Calculations

• Given mass and volume, calculate density using the formula above.

• Given density and volume, calculate mass:

• Given mass and density, calculate volume:

Example: A sample of lead (Pb) has mass 22.7 g and volume 2.0 cm3. Density =

Properties of Matter

Chemical vs. Physical Properties

• Chemical properties: Describe how a substance reacts with other substances (e.g., reactivity with water).

• Physical properties: Can be observed without changing the chemical identity (e.g., melting point, density).

• Extensive properties: Depend on the amount of substance present (e.g., mass, volume).

• Intensive properties: Do not depend on the amount of substance (e.g., density, color).

Examples: Electrical conductivity, ductility (can be drawn into wire), malleability (can be hammered into shape), brittleness, magnetism.

States of Matter and Changes of State

States of Matter

• Solid: Particles are closely packed in a fixed arrangement.

• Liquid: Particles are close but can move past each other.

• Gas: Particles are far apart and move freely.

Changes of State

• Melting: Solid to liquid

• Freezing: Liquid to solid

• Evaporation/Boiling: Liquid to gas

• Condensation: Gas to liquid

• Sublimation: Solid to gas

• Deposition: Gas to solid

Energy Changes in Chemical Processes

Kinetic Energy and Conservation of Energy

• Kinetic energy: Energy of motion.

• Law of Conservation of Energy: Energy cannot be created or destroyed, only transformed.

Example: Combustion of acetylene:

Endothermic vs. Exothermic Changes

• Endothermic change: System absorbs heat (e.g., water boiling, steam condensing).

• Exothermic change: System releases heat (e.g., water freezing, ice melting).

The Mole Concept

Counting Atoms: The Mole

Atoms are extremely small, so chemists use the mole to count large numbers of atoms.

• 1 mole = particles (Avogadro's number).

• For any element, 1 mole has a mass in grams equal to its atomic mass from the periodic table.

Island Diagram

Sample Problems

• How many moles is atoms of zinc?

• How many atoms is 0.68 moles of zinc?

• How many grams is 5.69 moles of uranium?

• How many grams is atoms of neon?

• How many atoms is 421 g of promethium?

Additional info: The notes cover foundational topics in general chemistry, including matter, classification, properties, mixtures, separation techniques, density, states of matter, energy changes, and the mole concept. These are essential for understanding chemical reactions, laboratory techniques, and quantitative chemistry.