Learning objectives
- Use particle motion and spacing to compare solids, liquids, and gases.
- Distinguish elements, compounds, homogeneous mixtures, and heterogeneous mixtures.
- Identify physical properties and changes without confusing evidence with proof of a reaction.
- Choose appropriate SI units, record sensible precision, and calculate density.
Prerequisites: You should be comfortable multiplying, dividing, and reading a simple table. The curriculum overview explains how this lesson leads into atomic structure.
1. A particle model of matter
Matter has mass and occupies space. A useful model says that matter is made of tiny particles in continual motion. The particles are too small to see directly, so the model connects observations at the scale of objects to behavior at the particle scale. Heating usually increases average particle kinetic energy; cooling usually decreases it.
In a solid, particles are close together and mainly vibrate around fixed locations. A solid keeps its shape and volume unless a force changes it. In a liquid, particles remain close but can move past one another, so a liquid keeps nearly the same volume but takes the shape of its container. In a gas, particles are far apart relative to their size and move freely, so a gas spreads to fill its container and is easy to compress. These are patterns, not claims that particles stop moving in a cold solid.
Changes of state such as melting and evaporation alter particle spacing and motion. They do not turn water particles into a new substance. At a fixed pressure, boiling is rapid vaporization throughout a liquid, while evaporation can occur at the surface below the boiling temperature.
2. Pure substances and mixtures
An element contains one kind of atom, such as copper or oxygen. A compound contains atoms of two or more elements chemically combined in a fixed ratio, such as water, H2O. A pure substance has a constant composition and characteristic properties under specified conditions. A sample can be pure even when it contains many particles, because it consists of one chemical substance with a consistent composition. In an ionic compound, that substance is represented by a repeating ratio of ions rather than by separate molecules.
A mixture contains substances together without requiring a fixed chemical ratio. In a homogeneous mixture, composition is uniform on the scale of the sample, as in salt water after the salt dissolves. In a heterogeneous mixture, different parts have visibly or measurably different composition, as in sand stirred into water. Filtration, distillation, and chromatography separate mixtures by physical properties; separation does not by itself make a chemical reaction.
Do not use “natural” or “pure” as synonyms. Natural materials can be mixtures, and a manufactured material can be a pure compound. State what evidence supports a classification and what scale you are using.
3. Properties and changes
A physical property can be observed or measured without changing a substance's chemical identity. Examples include density, color, melting point, and electrical conductivity. A chemical property describes how a substance can change into other substances, such as the ability of iron to react with oxygen.
A physical change can alter state, size, or shape while leaving chemical identities intact. Cutting aluminum, freezing water, and dissolving sugar are commonly treated as physical changes at this level. Dissolving distributes solute particles among solvent particles; it does not automatically mean the solute reacted. A chemical change produces substances with new chemical identities. Rusting iron and reacting vinegar with baking soda are examples. Bubbles, a temperature change, a color change, or a precipitate can be clues, but each can have a physical explanation. For example, boiling produces bubbles without a new substance.
Worked example 1: A student evaporates water from salt water and collects solid salt. The original sample was a homogeneous mixture. Evaporation was a physical change because water particles left as vapor and salt particles remained chemically salt. The mixture was separated, not converted into a new compound.
4. Measuring with units and uncertainty
A measurement includes a number, a unit, and limited precision. The SI base unit for length is the meter (m), for mass the kilogram (kg), and for time the second (s). Chemistry commonly uses grams (g), milliliters (mL), liters (L), centimeters (cm), and degrees Celsius (°C). A milliliter has the same volume as a cubic centimeter: 1 mL = 1 cm3. Temperature differences have the same size in Celsius and kelvin, but kelvin starts at absolute zero; K is written without a degree symbol.
Record every certain digit plus one estimated digit when reading an analog scale. Accuracy means closeness to an accepted value; precision means agreement among repeated measurements. A balance can be precise but inaccurate if it was not zeroed. Do not report more digits than the instrument supports.
Worked example 2: unit conversion and sensible reporting
A graduated cylinder reads 0.250 L. Since 1 L = 1000 mL, multiply by the conversion factor:
0.250 L × (1000 mL ÷ 1 L) = 250 mL
Liters cancel. The three significant digits support 250. mL or 2.50 × 102 mL, depending on classroom convention. Writing only “250 mL” can hide whether the final zero is measured.
5. Density connects mass and volume
Density is an intensive property: it does not depend on how much of a uniform sample you have. It is mass divided by volume:
density = mass ÷ volume or ρ = m ÷ V
Common units are g/mL and g/cm3; 1 cm3 = 1 mL by definition. To find the volume of a rectangular solid, multiply length × width × height. For an irregular object, measure water displacement in a graduated cylinder only when the object is fully submerged, does not dissolve or react in the water, and can be handled safely. Read the bottom of a water meniscus at eye level when that procedure is appropriate. Density can vary with temperature and pressure, especially for gases, so compare values only under stated or suitably similar conditions.
Worked example 3: density
A metal sample has mass 54.0 g. Its dimensions are 2.00 cm × 1.50 cm × 1.00 cm, so V = 3.00 cm3. Therefore:
ρ = 54.0 g ÷ 3.00 cm3 = 18.0 g/cm3
The result has three significant digits because each measured value has three. Density alone can suggest an identity, but it is not sufficient evidence when samples may contain impurities or measurements have uncertainty.
Use m = ρV or V = m/ρ for a missing quantity. Keep units in the calculation so an impossible unit signals an error.
Practice with explained answers
Classify air in a classroom: element, compound, or mixture. Is it homogeneous on an ordinary room scale?
Show reasoning
Air is a mixture of gases, mainly nitrogen and oxygen with smaller amounts of other gases. It is approximately homogeneous in a well-mixed room, although local composition can vary.
A 24.6 g stone raises water from 35.0 mL to 44.0 mL. What is its density?
Show reasoning
The stone volume is 44.0 − 35.0 = 9.0 mL. Density = 24.6 g ÷ 9.0 mL = 2.7 g/mL, limited to two significant digits by the volume difference.
Is melting candle wax automatically a chemical change? What observation would you seek before deciding?
Show reasoning
Melting is a physical change because the wax remains wax while its state changes. Burning wax is chemical. Look for evidence that a new substance forms, while remembering that smoke or heat alone needs interpretation.
Convert 3.20 cm3 to mL.
Show reasoning
The conversion is exact at this scale: 1 cm3 = 1 mL, so the volume is 3.20 mL. Keep the measured three-digit precision.
Sources and next steps
Compare this lesson with OpenStax Chemistry 2e, phases and classification of matter, physical and chemical properties, and measurements and density. For SI definitions and accepted measurement practice, consult the National Institute of Standards and Technology SI units guide. These sources support the definitions and calculations here; this page uses a simplified particle model for introductory learning.