Chapter 3: Matter and Energy – Vocabulary Flashcards

Vocabulary flashcards covering key terms from Chapter 3: Matter and Energy, including classifications of matter, states and properties, temperature scales and conversions, energy concepts, calorimetry, and changes of state.

Matter

Anything that has mass and occupies space.

Pure Substance

Matter with a fixed or definite composition; either an element or a compound.

Element

Simplest type of pure substance; made up of atoms; each element is composed of one type of atom.

Atom

The basic unit of an element; the fundamental building block that defines the element.

Compound

Substance that contains atoms from two or more elements; definite composition; can be separated into elements by chemical reaction.

Mixture

Combination of two or more pure substances; can be separated by physical methods; not chemically bonded.

Homogeneous (Mixture/Solution)

Mixture with uniform composition; cannot see distinct parts; examples include air and seawater.

Heterogeneous (Mixture)

Mixture with nonuniform composition; has distinct regions or phases, e.g., oil and water.

Solid

State of matter with definite shape and definite volume; particles fixed and tightly packed.

Liquid

State of matter with definite volume but takes the shape of its container; intermediate particle movement.

Gas

State of matter with no definite shape or volume; particles far apart and move freely.

Physical Property

A characteristic observed or measured without changing the substance’s identity (e.g., color, shape, density).

Physical Change

A change that alters a physical property but does not change the identity of the substance (no new substance formed).

Chemical Property

A characteristic that describes the ability of a substance to undergo a chemical change to form new substances (e.g., flammability, reactivity).

Chemical Change

A process that results in one or more new substances with different properties.

Melting

Solid to liquid; occurs at the melting point, where particles gain enough energy to overcome forces.

Freezing

Liquid to solid; occurs at the freezing point, same temperature as melting point for a given substance.

Vaporization

Liquid to gas; includes evaporation and boiling; boiling point is the temperature at which liquids form gas.

Condensation

Gas to liquid; energy is released as the gas loses energy.

Sublimation

Solid to gas without passing through the liquid phase.

Deposition

Gas to solid; energy is released during the change.

Melting Point

Temperature at which a solid changes to a liquid.

Boiling Point

Temperature at which a liquid becomes a gas at a given pressure (water 100°C at sea level).

Temperature Scales

Systems for measuring temperature: Celsius (°C), Kelvin (K), and Fahrenheit (°F).

Celsius

SI temperature scale; water freezes at 0°C and boils at 100°C at sea level.

Kelvin

SI temperature unit; no degree symbol; 0 K is absolute zero; water freezes at 273 K and boils at 373 K.

Fahrenheit

Temperature scale; water freezes at 32°F and boils at 212°F at sea level.

Absolute Zero

Lowest possible temperature where molecular motion ceases (0 K, -273.15°C).

Specific Heat (c)

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).

Water Specific Heat

Specific heat of water: 4.184 J/g·°C (or 4.184 kJ/kg·K); very high compared to many substances.

q = m × c × ΔT

Heat transferred equals mass times specific heat times the temperature change.

Calorie (cal)

Energy required to raise the temperature of 1 g of water by 1°C; 1 kcal = 1000 cal; 1 Cal (nutrition) = 1 kcal.

Calorie (Cal) / Kilocalorie (kcal)

Nutrition energy unit equal to 1,000 calories; 1 Cal = 1 kcal = 4184 J = 4.184 kJ.

Calorimeter

Device used to measure heat transfer (q) during chemical reactions or food energy; energy values are reported in kcal/g or kJ/g.

Carbohydrate Energy Value

4 kcal/g or 17 kJ/g.

Fat Energy Value

9 kcal/g or 38 kJ/g.

Protein Energy Value

4 kcal/g or 17 kJ/g.

Specific Heat of Substances Table

A chart listing the specific heats of various elements and compounds (e.g., water 4.184 J/g·°C).

Heat of Fusion

Energy required to melt 1 g of a solid at its melting point (e.g., ice ~334 J/g).

Heat of Vaporization

Energy required to convert 1 g of a liquid at its boiling point into gas (water ~2260 J/g; ~540 cal/g).

Heat of Condensation

Energy released when 1 g of a gas changes to a liquid at the boiling point (equal in magnitude to the heat of vaporization for a given substance).

Heating and Cooling Curve

Graph showing temperature vs. time during heating or cooling; horizontal segments indicate phase changes, vertical segments indicate temperature change.