Chapter 3: Matter and Energy – Study Notes for GOB Chemistry

Matter and Energy

Introduction

This chapter explores the fundamental concepts of matter and energy, including the classification of matter, its physical and chemical properties, temperature scales, energy transformations, and changes of state. Understanding these concepts is essential for further study in chemistry, biology, and nutrition.

Classification of Matter

Pure Substances and Mixtures

• Matter is anything that has mass and occupies space.

• Matter is classified by composition into pure substances and mixtures:

  • Pure substances have a fixed or definite composition.

  • Mixtures contain two or more substances physically mixed, not chemically combined, and can be separated by physical methods.

Elements and Compounds

• Elements are pure substances made of only one type of atom (e.g., copper, lead, aluminum).

• Compounds are pure substances composed of two or more elements in a fixed ratio (e.g., water, table salt, hydrogen peroxide).

Types of Mixtures

• Homogeneous mixtures have a uniform composition throughout (e.g., brass, air, sugar water).

• Heterogeneous mixtures have a non-uniform composition; different parts are visible (e.g., hot fudge sundae, peach pie).

States and Properties of Matter

States of Matter

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

• Liquids: Indefinite shape, definite volume; particles are close but can move past each other.

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

Physical and Chemical Properties

• Physical properties can be observed or measured without changing the substance’s identity (e.g., color, melting point, density).

• Chemical properties describe a substance’s ability to change into a new substance (e.g., flammability, reactivity).

Physical and Chemical Changes

• Physical changes alter the state or appearance but not the composition (e.g., melting, boiling, cutting).

• Chemical changes produce new substances with different properties (e.g., burning, rusting, caramelizing sugar).

Temperature

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; no negative values.

Temperature Conversions

• To convert between Celsius and Fahrenheit:

  • $ F = 1.8 \times C + 32 $

  • $ C = \frac{F - 32}{1.8} $

• To convert between Celsius and Kelvin:

  • $ K = C + 273 $

  • $ C = K - 273 $

Energy

Kinetic and Potential Energy

• Kinetic energy is the energy of motion (e.g., moving water, running, rollerblading).

• Potential energy is stored energy due to position or composition (e.g., water at the top of a dam, gasoline in a tank).

Heat and Units of Energy

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

• SI unit: joule (J); 1 kilojoule (kJ) = 1000 J.

• Calorie (cal): Amount of energy needed to raise the temperature of 1 g of water by 1°C.

• 1 Calorie (nutritional Calorie, Cal) = 1000 cal = 1 kcal.

Energy and Nutrition

Calorimetry

• A calorimeter measures the heat released by burning food to determine its energy value.

• Energy values on food labels are given in Calories (Cal) or kilojoules (kJ).

Caloric Values of Nutrients

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

Example: A cup of whole milk with 13 g carbohydrate, 9 g fat, and 9 g protein contains: $ (13 \times 4) + (9 \times 9) + (9 \times 4) = 170 $ kcal (rounded to the tens place).

Specific Heat

Definition and Calculation

• Specific heat (SH) is the amount of heat required to raise the temperature of 1 g of a substance by 1°C.

• Units: J/g·°C or cal/g·°C.

• Heat equation:

$ q = m \times SH \times \Delta T $ where $q$ = heat (J), $m$ = mass (g), $SH$ = specific heat, $\Delta T$ = temperature change (°C).

Changes of State

Melting, Freezing, Sublimation, Deposition

• Melting: Solid to liquid at melting point.

• Freezing: Liquid to solid at freezing point.

• Sublimation: Solid to gas without becoming liquid (e.g., dry ice).

• Deposition: Gas to solid without becoming liquid.

Heat of Fusion and Vaporization

• Heat of fusion: Heat required to melt 1 g of solid or released when 1 g of liquid freezes.

• Heat of vaporization: Heat required to vaporize 1 g of liquid or released when 1 g of gas condenses.

Evaporation, Boiling, and Condensation

• Evaporation: Surface molecules of a liquid gain enough energy to become gas.

• Boiling: Molecules throughout the liquid gain enough energy to become gas.

• Condensation: Gas molecules lose energy and become liquid.

Heating and Cooling Curves

• Heating curve: Shows temperature changes and phase transitions as heat is added.

• Diagonal lines: Temperature changes within a single state.

• Plateaus (horizontal lines): Phase changes at constant temperature.

• Cooling curve: Shows temperature changes and phase transitions as heat is removed.

Summary Table: Physical vs. Chemical Properties and Changes

Type	Properties	Changes
Physical	Color, state, melting point, density	Melting, boiling, cutting
Chemical	Flammability, reactivity	Burning, rusting, tarnishing

Key Equations

• Temperature conversions:

  • $ F = 1.8 \times C + 32 $

  • $ C = \frac{F - 32}{1.8} $

  • $ K = C + 273 $

• Heat calculation:

  • $ q = m \times SH \times \Delta T $

• Energy from food:

  • $ \text{Total kcal} = (\text{g carb} \times 4) + (\text{g protein} \times 4) + (\text{g fat} \times 9) $