Introduction to Matter

Definitions and Properties

Chemistry is the study of matter, its properties, and the changes it undergoes. Understanding the basic definitions and classifications of matter is fundamental to all of chemistry.

• Matter: Anything that has mass and occupies space (volume).

• Mass: A measure of the amount of matter in an object. Common units: grams (g), kilograms (kg).

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

• State of Matter: The physical form in which matter exists: solid, liquid, or gas.

• Composition: The types of particles (atoms, molecules, ions) that make up a sample of matter. Example: Copper (element), water (compound).

• Properties: Characteristics used to describe matter, such as color, density, melting point, etc.

• Atom: The basic building block of matter; the smallest unit of an element that retains its properties.

Classification of Matter

Elements and Compounds

Matter can be classified as either a pure substance or a mixture. Pure substances include elements and compounds.

• Element: A substance that contains only one type of atom. Elements cannot be broken down into simpler substances by chemical means.

• Monatomic Elements: Consist of single, unbonded atoms (e.g., noble gases like Ne, Ar).

• Polyatomic Elements: Consist of several like atoms bonded together (e.g., O2, P4).

• Diatomic Elements: Elements that naturally exist as molecules of two atoms (e.g., O2, N2).

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

• Compound: A substance composed of two or more different types of atoms chemically bonded together. Compounds have properties different from their constituent elements.

• Example: Sodium (Na) and chlorine (Cl2) are elements; sodium chloride (NaCl) is a compound.

• Atoms can only be altered by nuclear means. Molecules can be altered by chemical means.

Examples of Chemical Reactions:

• Dehydration of sugar:

• Electrolysis of water:

Mixtures

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

• Homogeneous Mixture (Solution): Has uniform composition and properties throughout. Particles are mixed at the molecular level. Example: Saltwater, air.

• Heterogeneous Mixture: Has non-uniform composition; different parts have different properties. Example: 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).

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

Chart for Classifying Matter

Classification Tree:

• Matter

  • Pure Substance

    • Element

    • Compound

  • Mixture

    • Homogeneous

    • Heterogeneous

Separating Mixtures

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

• Sorting: Separating based on physical characteristics (size, color, shape).

• Filtration: Separating solids from liquids using a filter.

• Magnet: Using a magnet to separate magnetic materials from non-magnetic ones.

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

• Density: Separating substances based on differences in density.

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

Density and Density Calculations

Density is a physical property defined as mass per unit volume. It is useful for identifying substances and predicting whether an object will float or sink in a fluid.

• Formula:

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

• Density of water: 1.00 g/mL (at 4°C).

Sample Calculations:

• A sample of lead (Pb) has mass 22.7 g and volume 2.0 cm3. Find density:

• Another sample of lead occupies 16.2 cm3. If density is 11.35 g/cm3, mass is:

• A solid cylinder with radius 1.8 cm and height 1.5 cm:

• A rectangular solid with edge lengths 8.2 cm, 5.1 cm, and 4.7 cm:

Physical and Chemical Properties

• 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 (e.g., color, melting point, density).

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

• Intensive Properties: Do not depend on the amount of substance (e.g., density, boiling point).

Examples:

• Electrical conductivity (intensive, physical)

• Reactivity with water (chemical)

• Heat content (extensive, physical)

• Ductility (intensive, physical): Can be drawn into wire

• Malleability (intensive, physical): Can be hammered into shape

• Brittleness (intensive, physical): Tendency to break

• Magnetism (intensive, physical)

States of Matter and Changes of State

• 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 each other.

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

Changes of State: Transitions between solid, liquid, and gas (e.g., melting, freezing, boiling, condensation, sublimation).

Energy in Chemistry

Kinetic Energy and Conservation of Energy

• Kinetic Energy: The energy of motion.

• Law of Conservation of Energy: Energy cannot be created or destroyed, only transformed from one form to another.

Example: Combustion of acetylene:

Energy Changes: Endothermic and Exothermic

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

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

The Mole and Counting Atoms

Atoms and molecules are extremely small, so chemists use the mole to count them in practical quantities.

• Mole (mol): The amount of substance containing as many entities (atoms, molecules) as there are atoms in 12 g of carbon-12.

• Avogadro's Number: entities per mole.

• 1 mole of atoms = atoms.

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

Island Diagram for Conversions

• Conversions between mass, moles, and number of particles (atoms/molecules) are fundamental in chemistry.

• Key relationships:

  • 1 mole = particles

  • Molar mass (g/mol) from periodic table

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?

General Conversion Equations:

• From moles to particles:

• From particles to moles:

• From moles to mass:

• From mass to moles:

Additional info: These notes provide a foundation for understanding matter, its classification, and basic quantitative relationships in chemistry, which are essential for further study in the subject.