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 present in everyday environments.

• Carbon atoms in graphite represent another type of atomic arrangement.

• Molecular structure determines the properties and behavior of matter.

• Example: Water (H2O) molecules versus carbon atoms in graphite.

Defining Matter

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

• Key Point: Matter includes all substances, regardless of visibility.

• Example: Air is matter, even though it cannot be seen directly.

Atomic and Molecular Composition 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.

• Atom: The fundamental building block of matter.

• Molecule: Two or more atoms bonded together in specific geometric arrangements.

• Example: Water is composed of H2O molecules; aluminum consists of independent aluminum atoms.

Imaging Atoms and Molecules

Advances in microscopy, such as the scanning tunneling microscope (STM), allow scientists to image individual atoms and molecules.

• Example: STM images of nickel atoms and DNA molecules reveal atomic and molecular structure.

States of Matter

Common States of Matter

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

• Solid: Atoms/molecules are closely packed in fixed locations; solids have fixed volume and shape.

• Liquid: Atoms/molecules are close but free to move; liquids have fixed volume but take the shape of their container.

• Gas: Atoms/molecules are far apart and move freely; gases are compressible and take both the shape and volume of their container.

• Example: Water exists as ice (solid), liquid water, and steam (gas).

Types of Solids

Solids can be classified as crystalline or amorphous based on atomic arrangement.

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

• Amorphous solid: Atoms/molecules lack 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 by Composition

Pure Substances and Mixtures

Matter can be classified as a pure substance or a mixture.

• 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: Helium is an element; water (H2O) is a compound.

Mixtures: Homogeneous and Heterogeneous

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

• Heterogeneous mixture: Composition varies from one region to another (e.g., oil and water).

• Example: Air is a homogeneous mixture; seawater is a mixture of salt and water.

Properties and Changes of Matter

Physical and Chemical Properties

• Physical property: A property displayed without changing the composition of the substance (e.g., odor, boiling point).

• Chemical property: A property displayed only through changing the composition (e.g., flammability, rusting).

Physical and Chemical Changes

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

• Chemical change: Matter changes composition, forming new substances (e.g., rusting iron, burning butane).

• Example: Melting ice is a physical change; burning butane is a chemical change.

Conservation Laws

Law of Conservation of Mass

In chemical reactions, mass is conserved: matter is neither created nor destroyed.

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

• Example: Burning butane: 58 g butane + 208 g oxygen → 176 g carbon dioxide + 90 g water (total mass before and after is 266 g).

Law of Conservation of Energy

Energy is conserved in physical and chemical changes: it can be transferred or transformed, but not created or destroyed.

• Example: Energy released or absorbed in chemical reactions.

Energy in Chemistry

Forms 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 1 g of water by 1 °C.

• Nutritional Calorie (Cal): 1 Cal = 1000 cal.

• Kilowatt-hour (kWh): 1 kWh = $3.60 \times 10^6$ J.

Unit	Equivalent
1 cal	4.184 J
1 Cal	1000 cal
1 kWh	$3.60 \times 10^6$ J

Exothermic and Endothermic Reactions

• Exothermic reaction: Releases energy to surroundings.

• Endothermic reaction: Absorbs energy from surroundings.

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): Absolute zero is 0 K; water freezes at 273 K, boils at 373 K.

Scale	Freezing Point	Boiling Point
Fahrenheit	32 °F	212 °F
Celsius	0 °C	100 °C
Kelvin	273 K	373 K

• Conversion formulas:

$K = ^\circ C + 273.15$ $^\circ F = (^\circ C \times \frac{9}{5}) + 32$ $^\circ C = (^\circ F - 32) \times \frac{5}{9}$

Heat Capacity and Specific Heat

Heat capacity is the quantity of heat required to change the temperature of a 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: $q = m \times C \times \Delta T$

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

• Example: To raise the temperature of 2.5 g gallium from 25.0 °C to 29.9 °C (C = 0.372 J/g °C): $q = 2.5 \text{ g} \times 0.372 \text{ J/g °C} \times (29.9 - 25.0) \text{ °C} = 4.6 \text{ J}$

Summary of Key Concepts

• Matter is anything that occupies space and has mass; it exists as solid, liquid, or gas.

• Classification: Pure substances (elements, compounds) and mixtures (homogeneous, heterogeneous).

• Properties: Physical properties do not change composition; chemical properties do.

• Conservation: Mass and energy are conserved in physical and chemical changes.

• Energy: Measured in joules, calories, Calories, and kilowatt-hours; exothermic reactions release energy, endothermic absorb energy.

• Temperature: Related to molecular motion; measured in °F, °C, and K.

• Heat capacity: Determines temperature change upon heat absorption.

Chemical Skills 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.