Introductory Definitions

Matter

Matter is defined as anything that has mass and occupies volume. All physical substances are composed of matter, which can exist in different forms and states.

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

• Volume: The amount of space an object occupies. Common units include liters (L), cubic decimeters (dm3), milliliters (mL), and cubic centimeters (cm3).

State of Matter

Matter exists in different states, primarily solid, liquid, and gas, each with distinct physical properties.

• Solid: Definite shape and volume.

• Liquid: Definite volume but takes the shape of its container.

• Gas: No definite shape or volume; expands to fill its container.

Composition

The composition of matter refers to the types of substances present. For example:

• Copper: An element.

• Water: A compound (H2O).

Atom

An atom is the basic building block of matter, consisting of protons, neutrons, and electrons.

Classification of Matter

Elements

Elements are pure substances that contain only one type of atom. They are the simplest form of matter and cannot be broken down by chemical means.

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

• Polyatomic elements: Consist of several "like" 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.

• Example: Sodium chloride (NaCl) is composed of sodium (Na) and chlorine (Cl).

• Atoms can only be altered by nuclear means.

• Molecules can be altered by chemical means.

Example Equations:

• Dehydration of sugar: $\mathrm{C_{12}H_{22}O_{11}(s) \rightarrow 12\ C(s) + 11\ H_2O(g)}$

• Electrolysis of water: $\mathrm{2\ H_2O(l) \rightarrow 2\ H_2(g) + O_2(g)}$

Mixtures

Mixtures are combinations of two or more substances that are not chemically bonded. They can be separated by physical means.

• Homogeneous mixture (solution): Uniform composition and properties throughout (e.g., salt water).

• 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).

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

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

Chart for Classifying Matter

Separating Mixtures

Physical Methods

Mixtures can be separated by physical means, which do not involve changing 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. It is used to characterize substances and predict their behavior in mixtures.

• Formula: $\mathrm{Density = \frac{mass}{volume}}$

• Typical units: g/cm3 for solids, g/mL for fluids.

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

Density Calculations

• To find volume of regular solids, use geometric formulas (e.g., for a cylinder: $\mathrm{V = \pi r^2 h}$).

• For irregular solids, use water displacement.

Example Problems:

• A sample of lead (Pb) has mass 22.7 g and volume 2.0 cm3. Find density: $\mathrm{Density = \frac{22.7\ g}{2.0\ cm^3} = 11.35\ g/cm^3}$

• A solid cylinder with radius 1.8 cm and height 1.5 cm: $\mathrm{V = \pi \times (1.8\ cm)^2 \times 1.5\ cm}$

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 vs. Intensive Properties

• 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, boiling point).

Examples:

• Electrical conductivity

• Ductility: Ability to be drawn into wire

• Malleability: Ability to 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 and Chemical Change

Kinetic Energy

Kinetic energy is the energy of motion. In chemistry, it is important for understanding particle movement and temperature.

Law of Conservation of Energy

Energy cannot be created or destroyed, only transformed from one form to another.

• Example: Combustion of acetylene: $\mathrm{2\ H_2 + O_2 \rightarrow 2\ H_2O}$

Energy Changes

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

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

The Mole Concept

Counting Atoms: The Mole

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

• Avogadro's number: $6.02 \times 10^{23}$ atoms per mole.

• 1 mole of atoms = $6.02 \times 10^{23}$ atoms.

• 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 $3.79 \times 10^{23}$ 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 $2.65 \times 10^{23}$ atoms of neon?

• How many atoms is 421 g of promethium?

Additional info:

• Some example problems and diagrams were inferred to provide context for calculations and concepts.

• Definitions and explanations were expanded for clarity and completeness.