Chapter 3 – Matter and Energy: Study Notes

Chapter 3 – Matter and Energy

Section 3.1: Classification of Matter

Matter is the material that makes up all things in the universe. It is anything that has mass and occupies space. Matter can be classified according to its composition into pure substances and mixtures.

• Pure Substances: Matter with a fixed or definite composition. There are two types:

  • Element: Composed of only one type of atom (e.g., copper, gold, helium).

  • Compound: Composed of two or more elements always combined in the same proportion (e.g., water, H2O; hydrogen peroxide, H2O2).

• Mixtures: Matter that consists of two or more substances that are physically mixed but not chemically combined. Mixtures can be separated by physical methods and may contain substances in different proportions.

Homogeneous vs. Heterogeneous Mixtures

• Homogeneous Mixture: The composition is uniform throughout, and the different parts are not visible. Also called a solution (e.g., coffee, brass).

• Heterogeneous Mixture: The composition varies from one part to another, and the different parts are visible (e.g., salad, trail mix).

Practice Example

• Aluminum foil and helium are examples of pure substances.

• Pasta and tomato sauce, and air, are mixtures.

Section 3.2: States and Properties of Matter

Matter exists in three primary states: solid, liquid, and gas. Each state has distinct properties based on the arrangement and movement of particles.

Solids

• Definite shape and volume.

• Particles are close together in a fixed arrangement and move very slowly (vibrate in place).

Liquids

• Indefinite shape but definite volume.

• Take the shape of their container.

• Particles are close together but mobile, moving at moderate speed.

Gases

• Indefinite shape and volume.

• Take the shape and volume of their container.

• Particles are far apart and move very fast.

Physical Properties of Matter

• Characteristics observed or measured without changing the identity of a substance (e.g., shape, state, boiling point, density, color).

Physical Changes

• No change in the identity or composition of the substance.

• Includes changes in state (e.g., melting, boiling) or physical shape (e.g., cutting, grinding).

• Examples: Water boiling, sugar dissolving, copper drawn into wires, paper cut into pieces, pepper ground into flakes.

Chemical Properties and Changes

• Describe the ability of a substance to interact with other substances or to change into a new substance.

• During a chemical change, the original substance is turned into one or more new substances with new composition, chemical, and physical properties.

• Example: Burning methane gas (CH4) in oxygen produces carbon dioxide (CO2) and water (H2O).

Practice Example

• Burning a log: chemical change

• Dissolving sugar in water: physical change

• Melting gold: physical change

• Digesting a hamburger: chemical change

• Fireworks exploding: chemical change

Section 3.3: Temperature

Temperature measures how hot or cold an object is compared to another object. It indicates the direction of heat flow and is measured using a thermometer.

Temperature Scales

• Fahrenheit (°F): 180 degrees between boiling (212°F) and freezing (32°F) points of water.

• Celsius (°C): 100 degrees between boiling (100°C) and freezing (0°C) points of water.

• Kelvin (K): Absolute zero is 0 K (−273.15°C). No negative temperatures. Same size units as Celsius (1 K = 1°C).

Temperature Conversion Equations

• From Celsius to Fahrenheit: $T_F = 1.8T_C + 32$

• From Fahrenheit to Celsius: $T_C = \frac{T_F - 32}{1.8}$

• From Celsius to Kelvin: $T_K = T_C + 273.15$

Practice Example

• Convert 34.8°C to °F: $T_F = 1.8 \times 34.8 + 32 = 94.6\,°F$

• Convert 37.22°C to K: $T_K = 37.22 + 273.15 = 310.4\,K$

Section 3.4: Energy

Energy is the ability to do work and is responsible for making objects move or stop. Work is defined as force applied over a distance.

• Work Equation: $\text{Work} = \text{Force} \times \text{Distance}$

Kinetic and Potential Energy

• Potential Energy: Based on position or chemical composition (e.g., compressed spring, chemical bonds, water at the top of a dam).

• Kinetic Energy: Associated with motion (e.g., working out, car driving, water flowing).

Energy as Heat

• Heat is the energy associated with the movement of particles.

• The faster the particles move, the greater the heat or thermal energy.

• Example: Adding heat to an ice cube increases the motion of H2O molecules, eventually causing melting.

Units of Energy

• SI unit: joule (J); 1 kJ = 1000 J

• Other units: calorie (cal); 1 kcal = 1000 cal

• 1 cal = 4.184 J (exact)

Energy Conversion Example

• How many calories are in 150 J? $\text{cal} = \frac{150\,J}{4.184\,J/\text{cal}} = 35.9\,\text{cal}$

Section 3.5: Energy and Nutrition

On food labels, energy is shown as the nutritional Calorie (Cal), which equals 1000 calories (1 kcal). In other countries, energy may be shown in kilojoules (kJ).

Caloric Food Values

Food Type	kcal/g	kJ/g
Carbohydrate	4	17
Fat	9	38
Protein	4	17

Guide to Calculating Energy from Food

• Multiply the grams of each food type by its energy value (kcal/g).

• Add the energy from all food types to get the total energy.

• Example: A cup of whole milk contains 13 g carbohydrate, 9 g fat, 9 g protein. Total kcal = $(13 \times 4) + (9 \times 9) + (9 \times 4) = 52 + 81 + 36 = 169$ kcal

Additional info: These notes cover the main concepts from Chapter 3, including matter classification, states and properties, physical and chemical changes, temperature and its measurement, energy forms and units, and nutritional energy calculations. Practice problems and examples are included to reinforce understanding.