Ch.3 Matter and Energy – Study Notes

Matter and Energy - CH. 3

Classification of Matter

Matter is anything that occupies space and has mass. It can be classified based on its composition and physical state.

• Pure Substances: Matter with a fixed composition; can be elements or compounds.

• Mixtures: Matter composed of two or more substances physically combined; can be homogeneous (uniform composition) or heterogeneous (non-uniform composition).

Classification Table:

Type	Subtypes	Examples
Pure Substance	Element, Compound	Gold, Water
Mixture	Homogeneous, Heterogeneous	Saltwater, Salad

Example: Saltwater is a homogeneous mixture; salad is a heterogeneous mixture.

States of Matter

Substances exist as solids, liquids, or gases depending on temperature and pressure.

• Solid: Definite shape and volume; particles are closely packed.

• Liquid: Definite volume, indefinite shape; particles can move past each other.

• Gas: Indefinite shape and volume; particles are far apart and move freely.

Key Properties:

• Compressibility: Gases are compressible; solids and liquids are not.

• Viscosity: Resistance to flow; higher in some liquids (e.g., honey) than others (e.g., water).

Example: Helium gas has the lowest viscosity among common substances.

Physical and Chemical Changes

Changes in matter can be classified as physical or chemical.

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

• Chemical Change: Alters composition, forming new substances (e.g., burning wood).

Example: Dissolving sugar in water is a physical change; rusting iron is a chemical change.

Reversible and Irreversible Changes

• Reversible Change: Can be undone (e.g., freezing and melting water).

• Irreversible Change: Cannot be undone (e.g., burning paper).

Phase Changes: Melting, freezing, boiling, and condensation are reversible changes involving energy transfer.

Chemical Properties

Chemical properties describe how a substance reacts with other substances, resulting in a change in composition.

• Examples: Reactivity with acids, flammability, ability to oxidize.

Example: The ability of hydrogen to react with oxygen is a chemical property.

Physical Properties

Physical properties can be observed or measured without changing the substance’s composition.

• Examples: Color, density, melting point, boiling point, electrical conductivity.

Example: Mercury is a silvery liquid at 25°C (physical property).

Temperature and Heat

Temperature measures the average kinetic energy of particles; heat is energy transferred due to temperature difference.

• Thermal Energy: Total kinetic and potential energy of particles in a substance.

• Temperature Conversions:

Conversion	Formula
Celsius to Kelvin	$K = ^\circ C + 273.15$
Celsius to Fahrenheit	$^\circ F = (^\circ C \times \frac{9}{5}) + 32$

Law of Conservation of Mass

Mass is neither created nor destroyed in a chemical reaction; the mass of reactants equals the mass of products.

• Example Equation:

$2H_2 (g) + O_2 (g) \rightarrow 2H_2O (l)$

Example: Mixing 2.0 g of hydrogen with 16.0 g of oxygen yields 18.0 g of water.

Nature of Energy

Energy is the capacity to do work or produce heat. It exists as kinetic (motion) or potential (position or composition) energy.

• Thermochemistry: Study of energy changes in chemical reactions.

• Energy Conversion Factors:

Unit	Equivalent
1 calorie (cal)	4.184 joules (J)
1 kilocalorie (kcal)	1000 calories (cal)
1 kilowatt-hour (kWh)	3.60 × 106 J

First Law of Thermodynamics

Energy cannot be created or destroyed, only transformed between system and surroundings.

• System: The part of the universe being studied.

• Surroundings: Everything else.

Heat and Work: Energy can be transferred as heat (q) or work (w).

$\Delta E = q + w$

Endothermic and Exothermic Reactions

• Endothermic: Absorbs heat from surroundings (e.g., melting ice).

• Exothermic: Releases heat to surroundings (e.g., combustion).

Heat Capacity

Heat capacity is the amount of heat required to change the temperature of a substance by 1 K.

• Specific Heat Capacity (c): Amount of heat required to raise 1 g of a substance by 1°C.

• Molar Heat Capacity (C): Amount of heat required to raise 1 mol of a substance by 1°C.

Formulas:

• Heat: $q = m c \Delta T$

• Molar Heat Capacity: $q = n C \Delta T$

Example: If 15.7 g of silver raises its temperature by 17.2°C when it absorbs 648.5 J, its specific heat capacity is $c = \frac{q}{m \Delta T}$.

Thermal Equilibrium

When two substances at different temperatures are in contact, heat flows from the hotter to the cooler substance until both reach the same temperature.

Summary Table: Specific Heat Capacities

Substance	Specific Heat Capacity (J/g·°C)
Water (l)	4.184
Copper (s)	0.385
Gold (s)	0.129
Ethanol (l)	2.46
Wood	1.76