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

Chapter 3: Matter and Energy

Classification of Matter

Matter is anything that has mass and occupies space. It can be classified into pure substances and mixtures, each with distinct characteristics.

• Pure Substances: Have a fixed or definite composition. Examples include elements (e.g., copper) and compounds (e.g., water).

• Mixtures: Consist of two or more substances physically combined. Mixtures can be homogeneous (uniform composition, e.g., brass, sugar water) or heterogeneous (non-uniform composition, e.g., peach pie, hot fudge sundae).

Example: Air is a mixture, while helium is a pure substance.

States of Matter

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

• Solids: Definite shape and volume; particles are closely packed and vibrate in fixed positions.

• Liquids: Definite volume but no definite shape; particles are close together but move randomly.

• Gases: Indefinite shape and volume; particles are far apart, move rapidly, and fill the container.

Example: Vitamin tablets are solids, vegetable oil is a liquid, and air in a basketball is a gas.

Physical and Chemical Properties

Properties of matter are classified as physical or chemical, depending on whether the identity of the substance changes.

• Physical Properties: Observed or measured without changing the substance's identity (e.g., color, melting point, density).

• Chemical Properties: Describe the ability of a substance to change into a new substance (e.g., ability to burn, tarnish).

Example: Copper is shiny (physical property); paper can burn (chemical property).

Physical and Chemical Changes

Changes in matter can be physical or chemical.

• Physical Change: Alters the physical appearance or state without changing composition (e.g., melting ice, cutting dough).

• Chemical Change: Produces new substances with different properties (e.g., rusting iron, burning a candle).

Example: Ice melting is a physical change; iron rusting is a chemical change.

Temperature Scales

Temperature is measured using three main scales: Fahrenheit (°F), Celsius (°C), and Kelvin (K).

• Fahrenheit: Water freezes at 32°F and boils at 212°F.

• Celsius: Water freezes at 0°C and boils at 100°C.

• Kelvin: Absolute zero is 0 K; no negative temperatures; 1 K = 1°C.

Conversion Equations:

• To convert °C to °F: $T_F = 1.8(T_C) + 32$

• To convert °F to °C: $T_C = \frac{T_F - 32}{1.8}$

• To convert °C to K: $T_K = T_C + 273$

Example: Normal body temperature (37°C) in kelvins: $T_K = 37 + 273 = 310$ K.

Energy: Kinetic and Potential

Energy is the ability to do work and exists in two main forms:

• Kinetic Energy: Energy of motion (e.g., rollerblading).

• Potential Energy: Stored energy due to position or composition (e.g., gasoline in a tank).

Example: Water behind a dam has potential energy; when released, it becomes kinetic energy.

Heat and Units of Energy

Heat is the energy associated with particle motion. The SI unit is the joule (J), and the calorie (cal) is also used.

• 1 cal = 4.184 J

• 1 kcal (Calorie) = 1000 cal = 4184 J

Example: 150 J of energy is equivalent to $\frac{150}{4.184} \approx 36$ cal.

Specific Heat

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

• Units: J/g°C or cal/g°C

• Equation: $SH = \frac{\text{heat}}{m \times \Delta T}$

Example: If 24.8 g of a metal absorbs 275 J and temperature rises from 20.2°C to 24.5°C, $\Delta T = 24.5 - 20.2 = 4.3$°C $SH = \frac{275}{24.8 \times 4.3}$

Heat Equation

To calculate heat gained or lost:

• $\text{Heat} = m \times \Delta T \times SH$

Example: To heat 255 g of copper from 24°C to 185°C ($SH = 0.385$ J/g°C): $\Delta T = 185 - 24 = 161$°C $\text{Heat} = 255 \times 161 \times 0.385$ J

Changes of State

Matter changes state at constant temperature:

• Melting: Solid to liquid at melting point

• Freezing: Liquid to solid at freezing point

• Sublimation: Solid to gas without liquid phase

• Deposition: Gas to solid

• Evaporation: Liquid to gas at surface

• Boiling: Liquid to gas throughout liquid

• Condensation: Gas to liquid

Example: Dry ice undergoes sublimation; water boils at 100°C.

Heat of Fusion

The heat of fusion is the energy required to convert 1 g of solid to liquid at its melting point.

• For water: 80 cal/g or 334 J/g

• Equation: $\text{Heat} = \text{mass} \times \text{heat of fusion}$

Example: To melt 32.0 g of ice: $\text{Heat} = 32.0 \times 334 = 10,688$ J $= 10.7$ kJ

Energy and Nutrition

Carbohydrates are the primary energy source for the body, followed by fats and proteins when reserves are depleted.

• Energy values on food labels are shown as Calories (kcal) or kilojoules (kJ).

• 1 Calorie (food Calorie) = 1 kcal = 4184 J

Concept Map: Matter and Energy

Matter can be classified as pure substances or mixtures, exists in different states, and undergoes physical or chemical changes. Energy is involved in changes of state and chemical reactions, and is measured in joules or calories.

Study Tips

• Read textbook chapters and highlight challenging concepts.

• Review slide decks and practice example problems.

• Use provided conversion factors and apply dimensional analysis for calculations.

• Pay attention to significant figures and correct unit usage.

Table: Classification of Matter

Type	Example	Description
Element	Copper	Pure substance, one type of atom
Compound	Water	Pure substance, two or more elements chemically combined
Homogeneous Mixture	Sugar water	Uniform composition throughout
Heterogeneous Mixture	Peach pie	Non-uniform composition

Table: Temperature Conversion Equations

Conversion	Equation (LaTeX)
°C to °F	$T_F = 1.8(T_C) + 32$
°F to °C	$T_C = \frac{T_F - 32}{1.8}$
°C to K	$T_K = T_C + 273$

Additional info: These notes expand on the provided slides and text, adding definitions, equations, and examples for clarity and completeness.