Introduction to Matter

Basic Definitions

Chemistry is the study of matter, its properties, and the changes it undergoes. Understanding the fundamental concepts of matter is essential for all further study in 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) that make up a sample of matter. Example: Copper (element), Water (compound).

• Atom: The basic building block of matter.

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.

  • Monatomic elements: Consist of unbonded, "like" atoms (e.g., He, Ne).

  • Polyatomic elements: Consist of several "like" atoms bonded together (e.g., O2, S8).

  • 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 that contains two or more different types of atoms chemically bonded together. Compounds have properties different from their constituent elements.

  • Example: Sodium chloride (NaCl) is made from sodium (Na) and chlorine (Cl2), but has very different properties from either element.

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

  • Example reactions:

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

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

Mixtures

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

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

• Heterogeneous mixture: Has different composition and properties in different parts of the sample (e.g., salad, sand in water).

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

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

Chart for Classifying Matter

This chart shows the main categories of matter and their subdivisions.

Properties of Matter

Chemical and Physical 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 of the substance (e.g., melting point, density, color).

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 of Properties

• Electrical conductivity

• Ductility: Ability to be drawn into wire

• Malleability: Ability to be hammered into shape

• Brittleness: Tendency to break or shatter

• Magnetism

States of Matter and Changes of State

States of Matter

• 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

• 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

Separation of Mixtures

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

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

• 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 substances based on differences in density.

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

Density

Definition and Calculation

• Density (d): The mass of a substance per unit volume.

Formula:

$$d = \frac{m}{V}$$

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

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

• Density of water at 4°C: 1.00 g/mL

Density Calculations

• Example 1: A sample of lead (Pb) has mass 22.7 g and volume 2.0 cm3. Find the density. $$d = \frac{22.7\ g}{2.0\ cm^3} = 11.35\ g/cm^3$$

• Example 2: Another sample of lead occupies 16.2 cm3 of space. If the density is 11.35 g/cm3, find the mass. $$m = d \times V = 11.35\ g/cm^3 \times 16.2\ cm^3 = 183.87\ g$$

• Example 3: A solid cylinder has radius 1.8 cm and height 1.5 cm. Find its volume and, given mass, its density. Volume of cylinder: $$V = \pi r^2 h$$

• Example 4: A rectangular solid has edge lengths 8.2 cm, 5.1 cm, and 4.7 cm. Volume: $$V = l \times w \times h$$

Law of Conservation of Energy

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

• Kinetic energy: Energy of motion.

• Potential energy: Stored energy due to position or composition.

Energy Changes in Chemical Reactions

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

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

The Mole Concept

Definition and Use

• Atoms and molecules are counted using the mole, a fundamental unit in chemistry.

• 1 mole = $$6.02 \times 10^{23}$$ 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

This diagram helps visualize conversions between mass, moles, and number of particles.

Sample Problems

• How many moles is 3.79 x 1025 atoms of zinc? $$\text{Moles} = \frac{3.79 \times 10^{25}}{6.02 \times 10^{23}} = 62.96\ mol$$

• How many atoms is 0.68 moles of zinc? $$\text{Atoms} = 0.68 \times 6.02 \times 10^{23} = 4.09 \times 10^{23}$$

• How many grams is 5.69 moles of uranium? (U atomic mass ≈ 238 g/mol) $$\text{Mass} = 5.69 \times 238 = 1353.22\ g$$

• How many grams is 2.65 x 1023 atoms of neon? (Ne atomic mass ≈ 20.18 g/mol) $$\text{Moles} = \frac{2.65 \times 10^{23}}{6.02 \times 10^{23}} = 0.440\ mol$$ $$\text{Mass} = 0.440 \times 20.18 = 8.08\ g$$

• How many atoms is 421 g of promethium? (Pm atomic mass ≈ 145 g/mol) $$\text{Moles} = \frac{421}{145} = 2.90\ mol$$ $$\text{Atoms} = 2.90 \times 6.02 \times 10^{23} = 1.75 \times 10^{24}$$

Summary Table: Elements vs. Compounds

Elements and compounds are both pure substances, but differ in their atomic composition and how they can be separated.

Additional info: Some example calculations and definitions were expanded for clarity and completeness.