Matter and Its Classification

Introductory Definitions

Matter is anything that has mass and occupies volume. Understanding the basic properties and classification of matter is foundational in chemistry.

• Mass: The amount of matter in an object; measured in units such as 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).

States of Matter

Matter exists in different physical forms called states: solid, liquid, and gas. Each state has distinct properties regarding shape, volume, and particle arrangement.

• Solid: Definite shape and volume; particles are closely packed in a fixed arrangement.

• Liquid: Definite volume but no definite shape; particles are close but can move past one another.

• Gas: No definite shape or volume; particles are far apart and move freely.

Composition of Matter

Matter can be classified based on its composition into pure substances and mixtures.

• Pure Substances: Have a fixed composition and distinct properties. They can be elements or compounds.

• Mixtures: Consist of two or more substances physically combined. Their composition can vary.

Elements and Compounds

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 Ne).

• Polyatomic Elements: Consist of several like atoms bonded together (e.g., P4 for phosphorus).

• Diatomic Elements: Elements that naturally exist as molecules of two atoms (e.g., O2, H2).

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

Compounds

Compounds are pure substances composed of 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).

• Compounds can only be separated into their elements by chemical means.

• Molecules can be altered by physical or chemical means.

Example Equations:

• Dehydration of sugar:

• Electrolysis of water:

Mixtures and Their Classification

Types of Mixtures

Mixtures are combinations of two or more substances that are not chemically bonded. They can be classified as homogeneous or heterogeneous.

• Homogeneous Mixtures (Solutions): Uniform composition and properties throughout; particles are evenly mixed (e.g., saltwater, air).

• Heterogeneous Mixtures: Non-uniform composition; different parts have different properties (e.g., salad, sand in water).

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

• Suspension: A heterogeneous mixture where particles settle over time (e.g., muddy water).

Example: 24K gold is a pure substance (element), while 14K gold is a mixture (alloy).

Chart for Classifying Matter

A flowchart for classifying matter:

• Matter

  • Pure Substance

    • Element

    • Compound

  • Mixture

    • Homogeneous

    • Heterogeneous

Separation of Mixtures

Mixtures can be separated by physical means, which involve physical changes rather than chemical reactions.

• Sorting: Separating based on physical characteristics.

• Filtration: Separating solids from liquids using a filter.

• Magnet: Using a magnet to separate magnetic materials.

• Chromatography: Separating substances based on their movement through a medium.

• Density: Separating based on differences in density.

• Distillation: Separating based on differences in boiling points.

Properties of Matter

Chemical and Physical Properties

Properties of matter can be classified as chemical or physical.

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

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

Extensive and 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 (can be drawn into wire), malleability (can be hammered into shape), brittleness, magnetism.

Density and Density Calculations

Definition and Units

Density is the mass per unit volume of a substance. It is an intensive property and is nearly constant for a given substance at a specific temperature and pressure.

• Formula:

• Typical Units: g/cm3 for solids, g/mL for liquids.

Example Calculations:

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

• Another sample of lead occupies 16.2 cm3. If density is known, mass = density × volume.

**Density of water: 1.00 g/mL at 4°C.

States of Matter and Changes of State

States of matter can change through physical processes such as melting, freezing, condensation, evaporation, and sublimation.

• 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

Diagram Description: A diagram shows particles closely packed in solids, loosely arranged in liquids, and far apart in gases. Arrows indicate transitions between states.

Energy and Chemical Changes

Kinetic Energy and Conservation of Energy

Kinetic energy is the energy of motion. The Law of Conservation of Energy states that energy cannot be created or destroyed, only transformed.

• Example: Combustion of acetylene:

Endothermic and Exothermic Changes

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

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

Diagram Description: Energy diagrams show reactants (R) and products (P) with energy absorbed or released during the reaction.

The Mole and Counting Atoms

The Mole Concept

Atoms and molecules are extremely small, so chemists use the mole to count them. One mole contains Avogadro's number of particles.

• Avogadro's Number: particles per mole.

• For any element, one mole has a mass in grams equal to its atomic mass (from the Periodic Table).

Island Diagram

A diagram shows the relationships between mass, moles, and number of particles (atoms or molecules):

• Mass (g) ↔ Moles (mol) ↔ Particles (atoms/molecules)

• 1 mole = particles

Sample Problems

• How many moles are in atoms of zinc?

• How many atoms are in 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 above provide a comprehensive overview of the foundational concepts in general chemistry, including matter, classification, properties, separation techniques, density, states of matter, energy changes, and the mole concept. These are essential for understanding more advanced topics in chemistry.