Introduction to Matter

Introductory Definitions

In chemistry, matter is defined as anything that has mass and occupies volume. Understanding the basic properties and classifications of matter is essential for studying chemical phenomena.

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

• Weight: The force exerted by gravity on an object’s mass.

• 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 primarily as solids, liquids, or gases.

Composition: Examples include copper (element) and water (compound).

Properties: Characteristics used to describe matter, such as color, density, and reactivity.

Atom: 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 argon).

• Polyatomic elements: Consist of molecules formed by bonding several identical atoms (e.g., O2 for oxygen).

• Diatomic elements: Molecules composed of two atoms of the same element (e.g., H2, N2, O2).

• 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).

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: NaCl (sodium chloride) is formed from sodium and chlorine, which have very different properties from the compound.

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

Examples of chemical reactions:

• Dehydration of sugar:

• Electrolysis of water:

Classifying Matter: Pure Substances and Mixtures

Matter can be classified as either pure substances or mixtures.

• Pure substances: Elements and compounds with uniform composition and properties.

• Mixtures: Combinations of two or more substances not chemically bonded.

Types of Mixtures

• Homogeneous mixture (solution): Particles are microscopic; sample has the same composition and properties throughout (e.g., salt water, air).

• Heterogeneous mixture: Different composition and properties in the same sample; unevenly mixed (e.g., sand in water, salad).

• Alloy: Homogeneous mixture of metals (e.g., 24K gold).

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

Chart for Classifying Matter

Separating Mixtures

Mixtures can be separated by physical means or physical changes. Common methods include:

1. Sorting: Separating based on physical characteristics.

2. Filtration: Separating solids from liquids using a filter.

3. Magnet: Using magnetic properties to separate substances.

4. Chromatography: Separating based on movement through a medium.

5. Density: Separating based on differences in density.

6. Distillation: Separating based on differences in boiling points.

Density and Density Calculations

Density is a physical property defined as mass per unit volume. It is used to characterize substances and predict their behavior in mixtures.

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

• Density formula:

• Density of water:

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

Density Calculation Examples

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

• Another sample of lead occupies 16.2 cm3 of space. If density is known, mass can be calculated:

• A 119.5 g solid cylinder has radius 1.8 cm and height 1.5 cm. Volume:

• A 153 g rectangular solid has edge lengths 8.2 cm, 5.1 cm, and 4.7 cm. Volume:

Properties of Matter

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

• Physical properties: Can be observed without changing the substance chemically (e.g., color, melting point).

• 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, heat content, 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 in 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 in Chemical Changes

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:

Energy Changes in Reactions

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

• Exothermic change: System releases heat (e.g., water freezing, CO2 subliming).

The Mole Concept

Counting Atoms: The Mole

Atoms are extremely small and numerous. Chemists use the mole to count atoms efficiently.

• 1 mole of atoms = atoms (Avogadro's number)

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

Island Diagram for the Mole

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?

Key Equations:

• Number of moles:

• Number of particles:

• Mass from moles:

Additional info: The notes provide foundational concepts for General Chemistry, including classification of matter, separation techniques, density, properties, states of matter, energy changes, and the mole concept. These are essential for understanding chemical reactions, laboratory techniques, and quantitative chemical analysis.