Chapter 3: Matter and Energy

3.1 Classification of Matter

Matter is the material that makes up all things, defined as anything that has mass and occupies space. Understanding the classification of matter is fundamental in chemistry, as it helps distinguish between different types of substances and their properties.

• Pure Substances: Have a fixed or definite composition. They are further classified as:

  • Elements: Composed of one type of atom (e.g., copper, lead, aluminum).

  • Compounds: Composed of two or more elements combined in a definite ratio (e.g., water, table salt).

• Mixtures: Contain two or more substances physically mixed but not chemically combined. Mixtures can be separated by physical methods and may exist in different proportions.

Examples: Aluminum foil is a pure substance (element), air is a mixture, table salt is a compound.

Homogeneous vs. Heterogeneous Mixtures

• Homogeneous Mixtures: Uniform composition throughout; different parts are not visible (e.g., brass, air, nitrox for scuba tanks).

• Heterogeneous Mixtures: Composition varies; different parts are visible (e.g., hot fudge sundae, peach pie).

Example: Shampoo is a heterogeneous mixture; sugar water is homogeneous.

3.2 States and Properties of Matter

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

• Solids: Definite shape and volume; particles are close together in a fixed arrangement and move slowly.

• Liquids: Indefinite shape but definite volume; particles are close but mobile and move slowly.

• Gases: Indefinite shape and volume; particles are far apart and move rapidly.

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

Physical Change: Change in state or shape without altering the substance's identity (e.g., melting, boiling, chopping wood).

Chemical Properties: Describe the ability of a substance to change into a new substance (e.g., burning, tarnishing).

Chemical Change: Formation of a new substance with new properties (e.g., sugar caramelizing).

3.3 Temperature

Temperature measures how hot or cold an object is, using scales such as Celsius, Fahrenheit, and Kelvin. Understanding temperature conversions is essential in chemistry.

• 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 is 0 K; no negative temperatures; same unit size as Celsius.

Conversion Formulas:

Example: Normal body temperature in kelvins is 310 K.

3.4 Energy

Energy is the ability to do work and makes objects move. It is classified as kinetic or potential energy.

• Kinetic Energy: Energy of motion (e.g., swimming, mowing the lawn).

• Potential Energy: Stored energy based on position or chemical composition (e.g., water at the top of a dam, gasoline in a tank).

Heat: Energy associated with the movement of particles; faster movement means greater heat.

Units of Energy: SI unit is joule (J); also measured in calories (cal) and kilocalories (kcal).

Example: 150 J of energy is equivalent to 36 cal.

3.5 Energy and Nutrition

Nutritionists use calorimeters to measure the energy content of food. Energy values are shown as Calories (Cal) or kilojoules (kJ) on food labels.

• Calorimeter: Device used to measure heat transfer during food combustion.

• Energy Values: Typical values for 1 g of food:

  • Carbohydrate: 4 kcal/g

  • Fat: 9 kcal/g

  • Protein: 4 kcal/g

Example: A cup of whole milk with 13 g carbohydrate, 9 g fat, and 9 g protein contains approximately 170 kcal.

Weight Loss: Occurs when food intake is less than energy output.

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. It varies for different substances and is important in understanding heat transfer.

• Formula:

• Where q is heat (J), m is mass (g), SH is specific heat (J/g°C), and ΔT is temperature change (°C).

Example: If 24.8 g of metal absorbs 275 J and temperature rises, use the formula to calculate specific heat.

3.7 Changes of State

Matter changes state when converted from one form to another, such as melting, freezing, vaporization, condensation, sublimation, and deposition.

• Melting/Freezing: Melting point (mp) and freezing point (fp) are reversible processes.

• Heat of Fusion: Amount of heat needed to melt 1 g of solid or released when 1 g of liquid freezes.

• Sublimation: Solid changes directly to gas; Deposition: Gas changes directly to solid.

• Evaporation/Boiling: Liquid changes to gas at surface (evaporation) or throughout (boiling).

• Heat of Vaporization: Heat absorbed to change 1 g of liquid to gas at boiling point.

Heating and Cooling Curves: Diagonal lines indicate temperature changes; plateaus indicate changes of state.

• Plateau: Constant temperature, change of state.

• Sloped Line: Temperature change.

Combined Heat Calculations: Total heat required for state changes and temperature changes can be calculated by summing individual steps.

• Formula for Heat of Fusion:

• Formula for Heat of Vaporization:

Example: To melt 32.0 g of ice, use the heat of fusion formula.

Summary Tables

Classification of Matter, Comparison of States, Physical and Chemical Properties, Energy Values, and Specific Heats are often summarized in tables for clarity.

Relevant Images

The following image is directly relevant to the chapter and visually reinforces the concept of matter and energy in chemistry:

Additional info: The textbook cover (image_1) visually illustrates the transformation of matter and energy, as effervescent tablets dissolve and release gas, representing both physical and chemical changes discussed in the chapter.