Introduction to Matter

Definitions and Properties

Chemistry is the study of matter, its properties, composition, and the changes it undergoes. Understanding the basic definitions and classifications of matter is foundational 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) that make up a sample of matter. Example: Copper (element), water (compound).

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

Properties of Matter

• 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 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).

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. Cannot be broken down into simpler substances by chemical means.

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

Types of Elements

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

• Diatomic: Consist of two atoms bonded together (e.g., O2, N2).

• Polyatomic: Consist of more than two atoms bonded together (e.g., P4, S8).

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

Table: Examples of Elements and Molecules

Additional info: Table illustrates the difference between atoms and molecules, and between monatomic, diatomic, and polyatomic species.

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; particles are evenly mixed (e.g., salt water, air).

• Heterogeneous Mixture: 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).

Table: Contrast Between 24K and 14K Gold

Chart for Classifying Matter

Separation of Mixtures

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

• Sorting: Physically separating substances based on observable properties.

• 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 an important physical property that relates the mass of a substance to its volume.

• Density (d): $d = \frac{\text{mass}}{\text{volume}}$

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

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

Example Calculations:

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

2. Another sample of lead occupies 16.2 cm3. If density is 11.35 g/cm3, mass: $m = d \times V = 11.35\ \text{g/cm}^3 \times 16.2\ \text{cm}^3 = 183.87\ \text{g}$

3. A solid cylinder with radius 1.8 cm and height 1.5 cm: Volume $V = \pi r^2 h = \pi \times (1.8)^2 \times 1.5 = 15.3\ \text{cm}^3$ (rounded). If mass is 119.5 g, density: $d = \frac{119.5}{15.3} = 7.81\ \text{g/cm}^3$

4. A rectangular solid with edge lengths 8.2 cm, 5.1 cm, and 4.7 cm: Volume $V = 8.2 \times 5.1 \times 4.7 = 196.5\ \text{cm}^3$. If mass is 153 g, density: $d = \frac{153}{196.5} = 0.78\ \text{g/cm}^3$

Additional info: If the density of an object is less than that of water, it will float; if greater, it will sink.

Physical and Chemical Changes

• Chemical Change: Alters the chemical composition of a substance (e.g., burning, rusting).

• Physical Change: Alters the physical form, not the chemical identity (e.g., melting, boiling).

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: Melting, freezing, vaporization, condensation, sublimation, and deposition.

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: $2 H_2 + O_2 \rightarrow 2 H_2O$

Energy Changes in 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

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

• 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: $1\ \text{mol} = 6.02 \times 10^{23}$ particles (atoms, molecules, ions).

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

Island Diagram

The island diagram is a visual tool for converting between mass, moles, and number of particles:

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

• 1 mol = $6.02 \times 10^{23}$ particles

Sample Problems

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

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

3. How many grams is 5.69 moles of uranium? $\text{Mass} = 5.69 \times \text{atomic mass of U (238.03 g/mol)} = 1,354.4$ g

4. How many grams is $2.65 \times 10^{23}$ atoms of neon? $\text{Moles} = \frac{2.65 \times 10^{23}}{6.02 \times 10^{23}} = 0.440$ mol; $\text{Mass} = 0.440 \times 20.18 = 8.88$ g

5. How many atoms is 421 g of promethium? $\text{Moles} = \frac{421}{145} = 2.90$ mol; $\text{Atoms} = 2.90 \times 6.02 \times 10^{23} = 1.75 \times 10^{24}$ atoms

Summary Table: Classification of Matter

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.