Chapter 3: Matter and Energy – GOB Chemistry Study Notes

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 have substances 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, sugar water).

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

Example: Scuba breathing mixtures (nitrox, heliox, trimix) are homogeneous mixtures.

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 together but mobile and move slowly.

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

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 new substances with new properties (e.g., sugar caramelizing, iron rusting).

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 Scale: 0°C (freezing point of water), 100°C (boiling point of water).

• Fahrenheit Scale: 32°F (freezing), 212°F (boiling).

• Kelvin Scale: Absolute zero (0 K), no negative temperatures, same unit size as Celsius.

Conversion Formulas:

• $ \text{Fahrenheit} = 1.8 \times \text{Celsius} + 32 $

• $ \text{Celsius} = \frac{\text{Fahrenheit} - 32}{1.8} $

• $ \text{Kelvin} = \text{Celsius} + 273 $

Example: Normal body temperature in kelvins is 310 K.

3.4 Energy

Energy is the ability to do work and is essential for movement and chemical reactions. It exists in two main forms: kinetic and potential energy.

• Kinetic Energy: Energy of motion (e.g., swimming, water flowing).

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

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

Units of Energy: Joule (J), kilojoule (kJ), calorie (cal), kilocalorie (kcal).

• $ 1 \text{calorie} = \text{energy needed to raise the temperature of 1 g of water by 1°C} $

• $ 1 \text{Calorie} = 1000 \text{calories} = 1 \text{kcal} $

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.

• Typical Energy Values: Carbohydrates, fats, and proteins have different energy values per gram.

Food Type	Energy Value (kcal/g)
Carbohydrate	4
Fat	9
Protein	4

Example: A cup of whole milk contains 13 g carbohydrate, 9.0 g fat, and 9.0 g protein. Total energy: $ (13 \times 4) + (9 \times 9) + (9 \times 4) = 52 + 81 + 36 = 169 $ kcal (rounded to 170 kcal).

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.

• Units: J/g°C (SI), cal/g°C (metric).

• Heat Equation: $ q = m \times c \times \Delta T $ where $ q $ = heat (J), $ m $ = mass (g), $ c $ = specific heat (J/g°C), $ \Delta T $ = temperature change (°C).

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

3.7 Changes of State

Matter changes state when it is converted from one form to another, such as melting, freezing, sublimation, deposition, evaporation, boiling, and condensation. These changes involve energy transfer.

• 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/Deposition: Sublimation is direct change from solid to gas; deposition is gas to solid.

• Evaporation/Boiling/Condensation: Evaporation occurs at the surface; boiling throughout the liquid; condensation is gas to liquid.

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

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

Combined Heat Calculations: To calculate total heat for a substance undergoing multiple changes, sum the heat for each step (warming, phase change).

• Example: To convert 15.0 g of liquid ethanol to gas, calculate heat for warming and vaporization, then add together.

Concept Map

A concept map summarizes the relationships between matter, energy, states, properties, and changes.