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 defined as anything that has mass and occupies space. The classification of matter is based on its composition, which determines its properties and behavior.

• Matter can be divided into pure substances and mixtures.

Pure Substances

• A pure substance is a form of matter with a fixed or definite composition.

• There are two types of pure substances:

  • Element: Composed of only one type of atom (e.g., copper, gold). Elements are listed in the periodic table.

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

Mixtures

• A mixture consists of two or more substances that are physically mixed but not chemically combined.

• Mixtures can have substances in different proportions and can be separated by physical methods.

• Types of mixtures:

  • Homogeneous mixture (solution): Composition is uniform throughout; different parts are not visible (e.g., brass, air).

  • Heterogeneous mixture: Composition varies from one part to another; different parts are visible (e.g., water and copper, salad).

Section 3.2: States and Properties of Matter

Matter exists in different physical states, each with distinct properties. The three common states are solids, liquids, and gases.

Solids

• Definite shape and volume

• Particles are close together in a fixed arrangement

• Particles vibrate slowly in place

Liquids

• Indefinite shape but definite volume

• Take the shape of their container

• Particles are close 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 rapidly

Physical Properties

• Observed or measured without changing the identity of the substance

• Include shape, state, boiling and freezing points, density, and color

Physical Changes

• Do not alter the identity or composition of a substance

• Examples: boiling water, dissolving sugar, cutting paper

Chemical Properties and Changes

• Describe the ability of a substance to interact with other substances and form new substances

• A chemical change results in a new substance with different composition and properties

• Examples: burning wood, digesting food, rusting iron

Section 3.3: Temperature

Temperature measures how hot or cold an object is relative to another. It determines the direction of heat flow and is measured using thermometers.

Temperature Scales

• Celsius (°C): 0 °C is the freezing point, 100 °C is the boiling point of water

• Fahrenheit (°F): 32 °F is the freezing point, 212 °F is the boiling point of water

• Kelvin (K): Absolute zero (0 K) is the lowest possible temperature; 1 K = 1 °C

Temperature Conversion Formulas

• Celsius to Fahrenheit:

• Fahrenheit to Celsius:

• Celsius to Kelvin:

Section 3.4: Energy

Energy is the ability to do work or produce heat. It is essential for all physical and chemical processes.

Kinetic and Potential Energy

• Kinetic energy: Energy of motion (e.g., moving car, flowing water)

• Potential energy: Stored energy due to position or composition (e.g., compressed spring, chemical bonds)

Units of Energy

• Joule (J): SI unit of energy

• Calorie (cal): 1 cal = 4.184 J (exact)

• Kilocalorie (kcal): 1 kcal = 1000 cal

Energy Conversion Example

• To convert 150 J to calories:

Section 3.5: Energy and Nutrition

Food provides energy, measured in nutritional Calories (Cal), where 1 Cal = 1 kcal = 1000 cal. Energy values for food types are standardized.

• To calculate total energy from food:

  • Multiply grams of each macronutrient by its energy value

  • Add the results for total energy

Section 3.6: Specific Heat

Specific heat is the amount of heat required to raise the temperature of 1 g of a substance by 1 °C. Each substance has a unique specific heat value.

• Formula: where = heat (J or cal), = mass (g), = temperature change (°C), = specific heat (J/g·°C or cal/g·°C)

Example Calculation

• If 24.8 g of metal absorbs 275 J and its temperature rises from 20.2 °C to 24.5 °C:

Section 3.7: Changes of State

Substances change state (solid, liquid, gas) by absorbing or releasing energy. These changes occur at specific temperatures for each substance.

Melting and Freezing

• Melting point (mp): Temperature at which a solid becomes a liquid

• Freezing point (fp): Temperature at which a liquid becomes a solid

• For water, mp and fp are 0 °C

Heat of Fusion ()

• Amount of heat required to melt 1 g of solid or released when 1 g of liquid freezes

• For water: 80 cal/g or 334 J/g

• Formula:

Vaporization and Condensation

• Boiling point (bp): Temperature at which a liquid becomes a gas

• Condensation point (cp): Temperature at which a gas becomes a liquid

• For water, bp and cp are 100 °C

Heat of Vaporization ()

• Amount of heat absorbed to vaporize 1 g of liquid or released when 1 g of gas condenses

• For water: 540 cal/g or 2260 J/g

• Formula:

Heating and Cooling Curves

• Show temperature changes and phase changes as heat is added or removed

• Plateaus represent phase changes (melting/freezing, boiling/condensing)

• Sloped regions represent temperature changes within a single phase

Summary Table: Heats of Fusion and Vaporization for Water

Example: To melt 32.0 g of ice at 0 °C:

Additional info: The images and diagrams in the slides reinforce the concepts of energy changes, classification of matter, and the relationship between potential and kinetic energy.