Chapter 3: Matter and Energy – Structured Study Notes

Matter and Energy

Atoms and Molecules in Your Room

All objects and substances in a room are composed of matter. The differences between various kinds of matter arise from the differences in the molecules and atoms that compose them.

• Water molecules are an example of molecules found in liquids.

• Carbon atoms in graphite are an example of atoms arranged in a solid structure.

• Atoms are the fundamental building blocks of matter.

• Molecules are groups of two or more atoms bonded together in specific geometric arrangements.

Example: Water (H2O) consists of two hydrogen atoms and one oxygen atom bonded together.

Defining Matter

Matter is anything that occupies space and has mass. Matter can be visible (e.g., steel, wood, plastic) or invisible without magnification (e.g., air, microscopic dust).

• Visible matter: Steel, water, wood, plastic

• Invisible matter: Air, microscopic dust

Atomic and Molecular Structure of Matter

Matter Is Composed of Atoms and Molecules

Although matter may appear smooth and continuous, it is composed of atoms and molecules at the submicroscopic level.

• Atoms: Submicroscopic particles that are the fundamental building blocks of matter.

• Molecules: Two or more atoms joined in specific geometric arrangements.

Advances in microscopy allow us to image atoms and molecules directly.

Types of Atomic and Molecular Particles

• Independent atomic particles: In substances like aluminum, atoms exist as independent particles.

• Well-defined molecular particles: In substances like rubbing alcohol, atoms bond together in well-defined structures called molecules.

States of Matter

Common States of Matter

Matter exists in three common states: solid, liquid, and gas. The state depends on the arrangement and movement of atoms and molecules.

• Solid: Atoms or molecules are closely packed in fixed locations. Solids have a fixed volume and rigid shape.

• Liquid: Atoms or molecules are close together but free to move around each other. Liquids have a fixed volume but take the shape of their container.

• Gas: Atoms or molecules are far apart and move freely. Gases are compressible and assume the shape and volume of their container.

Types of Solids

• Crystalline solid: Atoms or molecules are arranged in geometric patterns with long-range, repeating order (e.g., salt, diamond).

• Amorphous solid: Atoms or molecules do not have long-range order (e.g., glass, rubber, plastic).

Properties of Solids, Liquids, and Gases

State	Atomic/Molecular Motion	Spacing	Shape	Volume	Compressibility
Solid	Oscillation/vibration about fixed point	Close together	Definite	Definite	Incompressible
Liquid	Free to move relative to one another	Close together	Indefinite	Definite	Incompressible
Gas	Free to move relative to one another	Far apart	Indefinite	Indefinite	Compressible

Classification of Matter

Pure Substances and Mixtures

Matter can be classified according to its composition:

• Pure substance: Composed of only one type of atom or molecule.

• Mixture: Composed of two or more different types of atoms or molecules combined in variable proportions.

Elements and Compounds

• Element: A pure substance that cannot be broken down into simpler substances by chemical means. Elements are listed in the periodic table.

• Compound: A pure substance composed of two or more elements in fixed, definite proportions. Compounds can be decomposed into simpler substances.

Example: Water (H2O) is a compound composed of hydrogen and oxygen in a fixed ratio.

Types of Mixtures

• Homogeneous mixture: Has the same composition throughout (e.g., sweetened tea).

• Heterogeneous mixture: Has a composition that varies from one region to another (e.g., oil and water).

Summary Table: Classification of Matter

Type	Examples
Element	Copper, Helium
Compound	Water, Table Salt, Sugar
Homogeneous Mixture	Sweetened Tea, Air
Heterogeneous Mixture	Oil and Water, Seawater

Properties and Changes of Matter

Physical and Chemical Properties

• Physical property: A property that a substance displays without changing its composition (e.g., odor, boiling point).

• Chemical property: A property that a substance displays only through changing its composition (e.g., flammability, rusting).

Physical and Chemical Changes

• Physical change: Matter changes its appearance but not its composition (e.g., boiling water).

• Chemical change: Matter changes its composition, resulting in new substances (e.g., rusting iron).

Example: Boiling water is a physical change; burning butane is a chemical change.

Law of Conservation of Mass

The law of conservation of mass states that matter is neither created nor destroyed in a chemical reaction. The total mass of reactants equals the total mass of products.

• Equation: $\text{Reactants} \rightarrow \text{Products}$

• Example: Burning 58 g of butane with 208 g of oxygen produces 176 g of carbon dioxide and 90 g of water. Total mass before and after is 266 g.

Energy in Chemistry

Definition and Forms of Energy

Energy is the capacity to do work. The behavior of matter is driven by energy, which is conserved according to the law of conservation of energy.

• Kinetic energy: Energy associated with motion.

• Potential energy: Energy associated with position or composition.

• Electrical energy: Energy from the flow of electrical charge.

• Thermal energy: Energy from random motions of atoms and molecules.

• Chemical energy: A form of potential energy stored in chemical bonds.

Units of Energy

• Joule (J): SI unit of energy.

• Calorie (cal): Amount of energy required to raise the temperature of 1 g of water by 1 °C.

• Calorie (Cal): Nutritional calorie, equal to 1000 cal.

• Kilowatt-hour (kWh): 3.60 × 106 J.

Unit	Equivalent
1 cal	4.184 J
1 Cal	1000 cal
1 kWh	3.60 × 106 J

Energy Changes in Chemical Reactions

• Exothermic reaction: Energy is released.

• Endothermic reaction: Energy is absorbed.

Temperature and Heat

Temperature Scales

• Fahrenheit (°F): Water freezes at 32 °F, boils at 212 °F.

• Celsius (°C): Water freezes at 0 °C, boils at 100 °C.

• Kelvin (K): Water freezes at 273 K, boils at 373 K. Absolute zero is 0 K.

Conversion formulas:

• $K = ^\circ C + 273.15$

• $^\circ F = (^\circ C \times \frac{9}{5}) + 32$

• $^\circ C = \frac{^\circ F - 32}{1.8}$

Heat Capacity and Specific Heat

Heat capacity is the quantity of heat required to change the temperature of a given amount of substance by 1 °C. Specific heat capacity is the heat required to change the temperature of 1 g of a substance by 1 °C.

Substance	Specific Heat Capacity (J/g °C)
Gold	0.128
Silver	0.235
Copper	0.385
Iron	0.449
Aluminum	0.903
Ethanol	2.42
Water	4.184

Equation for heat energy:

• $q = m \times C \times \Delta T$

• Where $q$ is heat (J), $m$ is mass (g), $C$ is specific heat capacity (J/g °C), and $\Delta T$ is temperature change (°C).

Example: To raise the temperature of 2.5 g of gallium from 25.0 °C to 29.9 °C (with $C = 0.372$ J/g °C):

• $q = 2.5 \text{ g} \times 0.372 \text{ J/g °C} \times (29.9 - 25.0) \text{ °C} = 4.6 \text{ J}$

Chapter Review and Learning Objectives

• Classify matter as element, compound, or mixture.

• Distinguish between physical and chemical properties.

• Distinguish between physical and chemical changes.

• Apply the law of conservation of mass.

• Identify and convert among energy units.

• Convert between Fahrenheit, Celsius, and Kelvin temperature scales.

• Perform calculations involving transfer of heat and changes in temperature.