Basic Chemistry Concepts: Matter, Atoms, Molecules, and States

Matter

Anything that has mass and takes up space.

Atom

Smallest unit of matter; a pure substance.

Molecule

Two or more atoms chemically bonded together.

Pure Substance

Contains only one type of atom or molecule; has unique chemical and physical properties.

Element

A pure substance made of only one kind of atom; found on the periodic table. Example: Gold (Au).

Single Atom Element

Atoms of an element that do not bond to other atoms. Example: Helium (He).

Diatomic Molecule

Two atoms of the same element bonded together. Examples: H₂, O₂, N₂, Cl₂, Br₂, I₂, F₂.

Allotrope

Substance made of the same atoms but arranged differently. Examples: O₂ and O₃; carbon as diamond or graphite.

Compound

A pure substance composed of two or more different elements chemically bonded in a fixed ratio. Example: H₂O, C₆H₁₂O₆, NaCl.

Chemical Formula

Represents the ratio of elements in a compound. Example: H₂O = 2 Hydrogen : 1 Oxygen.

Separation of Compounds

Can only be separated chemically, not physically.

Mixture

Combination of two or more pure substances not chemically bonded. Example: Salt water.

Intermolecular Forces

Physical forces holding substances together in mixtures.

Properties of Mixtures

No definite chemical composition; each component retains its properties and can be separated physically.

Homogeneous Mixture

Substances evenly distributed. Example: Salt water.

Heterogeneous Mixture

Substances not evenly distributed. Example: Cereal and milk.

Compound vs Mixture

Compounds have definite proportions and properties; mixtures do not.

Crystalline Solid

Has a long-range repeating order in structure.

Amorphous Solid

Lacks long-range order in structure.

Physical Change

Changes physical properties without changing composition. Example: melting, cutting, dissolving.

Chemical Change

Forms new substances with new properties. Reactants → Products. Example: burning, rusting.

Law of Conservation of Mass

Mass and charge are conserved; matter cannot be created or destroyed.

Energy

Ability to do work.

Work

Result of a force applied over a distance.

Kinetic Energy

Energy of motion.

Potential Energy

Stored energy.

Temperature

Measure of the average kinetic energy of a substance.

Heat

Transfer of thermal energy between substances.

Exothermic Reaction

Releases energy; ΔH is negative.

Endothermic Reaction

Absorbs energy; ΔH is positive.

Activation Energy

Energy required to start a chemical reaction.

Activated Complex

Temporary structure formed during the transition from reactants to products.

Heat of Reaction (ΔH)

Energy of products minus energy of reactants.

Celsius to Fahrenheit

°F = (°C × 9/5) + 32

Fahrenheit to Celsius

°C = (°F − 32) × 5/9

Kelvin Conversion

K = °C + 273.15

Heating Curve

Graph showing temperature change as heat is added at a constant rate.

Cooling Curve

Shows temperature as heat is removed; reverse of heating curve.

Flat Region on Heating Curve

Energy breaks intermolecular forces, not increasing kinetic energy.

Slope Region on Heating Curve

Heat increases kinetic energy, raising temperature.

Heat of Fusion (Hf)

Energy needed to melt 1 gram of a substance.

Heat of Vaporization (Hv)

Energy needed to boil 1 gram of a substance.

Specific Heat Capacity

Energy required to raise 1 gram of a substance by 1°C. For water: 4.18 J/g°C.

Melting/Freezing Point

Temperature where solid and liquid coexist; same temperature.

Boiling/Condensation Point

Temperature where liquid and gas coexist; same temperature.

High Specific Heat

Substance that needs a lot of energy to change temperature; example: water.

q = mcΔT

Used when temperature changes. Example: q = (25.6g)(4.18J/g°C)(30.0°C) = 3210 J.

q = mHfus

Used for melting/freezing energy calculations.

q = mHvap

Used for boiling/condensation energy calculations.

Heat Capacity

Heat needed to raise an object's temperature by 1°C (not per gram).

Specific Heat vs Heat Capacity

Specific heat = per gram; heat capacity = whole object.