Difficulty: Introductory | Prerequisites: None. This is foundational material.
This topic covers the very first concepts in a general chemistry course: what matter is, how it is classified, and how chemists measure and describe it. Everything you learn later (reactions, stoichiometry, thermodynamics) assumes you already speak this language fluently. If you are joining mid-semester, start here. You should be comfortable with basic algebra and unit arithmetic before moving on.
Matter is anything with mass and volume. It comes in three states (solid, liquid, gas), is classified as either a pure substance or a mixture, and is described by physical and chemical properties. Chemists measure matter using metric units, significant figures and scientific notation, and they convert between units with dimensional analysis.
Matter
Anything that has mass and takes up space (volume). In simple terms, everything you can touch, weigh or pour is matter.
Physical property
A characteristic observed or measured without changing the substance's identity (e.g. boiling point, colour, density). Think of it as anything you can describe about a substance while leaving it intact.
Chemical property
Describes how a substance reacts to form new substances. You can only observe a chemical property by performing (or attempting) a chemical change.
Physical change
A change in form or appearance that does not alter the substance's composition. Melting ice is a physical change: the water molecules stay the same.
Chemical change
A change that produces one or more new substances with different compositions. Combustion, rusting and oxidation are chemical changes.
Element
A pure substance that cannot be broken down into simpler substances by any chemical means. Oxygen, gold and carbon are elements.
Compound
A pure substance made of two or more elements chemically combined in a fixed ratio. Water (H2O) and carbon dioxide (CO2) are compounds.
Mixture
Two or more substances physically combined in variable proportions. The components can be separated by physical methods such as filtration or distillation.
Significant figures (sig figs)
All the certain digits in a measurement plus one estimated digit. In simple terms, sig figs tell you how precise a measurement is.
Scientific notation
A way of expressing very large or very small numbers as a coefficient (between 1 and 10) multiplied by a power of 10.
Conversion factor
A ratio equal to 1 used to convert a measurement from one unit to another. You set up fractions so the unwanted unit cancels.
Density
Mass per unit volume, typically in g/mL or g/cm3. It tells you how much stuff is packed into a given space.
Specific gravity
The ratio of a substance's density to the density of water. It has no units and, numerically, equals the density in g/mL.
Matter exists in three primary physical states, each defined by how its particles are arranged and how freely they move.
Solid: Definite volume and definite shape. Particles are closely packed in a regular pattern and vibrate in place. A block of ice keeps its shape regardless of its container.
Liquid: Definite volume but takes the shape of its container. Particles are close together yet slide past one another, which is why liquids flow.
Gas: No definite shape or volume. Particles are far apart, move randomly and rapidly, and fill whatever container they occupy.
The key insight is that moving from solid to liquid to gas, particle spacing increases and the forces holding particles together weaken. This progression matters when you study phase changes and energy later in the course.
Physical properties can be observed or measured without altering the substance. Common examples: boiling point, melting point, colour, odour, state of matter, density.
Physical changes alter a substance's form or appearance but leave its composition untouched. Melting ice and tearing paper are physical changes: the molecules stay the same.
Chemical properties describe how a substance behaves in a reaction. You only discover a chemical property by watching (or attempting) a chemical change.
Chemical changes produce entirely new substances. Burning wood, rusting iron and the tarnishing of silver are all chemical changes. Signs of a chemical change include colour change, gas production, temperature change and formation of a precipitate.
The distinction between physical and chemical is one of the most commonly tested concepts in introductory chemistry. If the composition of the substance is different after the change, it is chemical. If the composition is the same, it is physical.
All matter falls into one of two broad categories.
Pure substances have a constant, uniform composition and cannot be separated into simpler parts by physical methods.
Elements cannot be broken down further by any chemical process. Oxygen (O), gold (Au) and carbon (C) are elements.
Compounds consist of two or more elements chemically bonded in a fixed ratio. Water (H2O) and carbon dioxide (CO2) are compounds. They can be broken down by chemical, but not physical, methods.
Mixtures contain more than one component in variable proportions. They can be separated by physical methods such as filtration, distillation or evaporation.
Homogeneous mixtures (solutions) look uniform throughout. Saltwater and air are examples.
Heterogeneous mixtures have visibly distinct regions. A salad or sand mixed with iron filings are heterogeneous.
When classifying a sample, ask two questions in order: (1) Is the composition uniform? If not, it is a heterogeneous mixture. (2) Can it be separated by physical means? If yes, it is a homogeneous mixture. If no, it is a pure substance, then ask whether it can be broken down chemically to distinguish an element from a compound.
Every measurement in chemistry has two parts: a number and a unit. Without the unit, the number is meaningless.
Metric (SI) base units
Length: metre (m)
Mass: gram (g) (the SI base is the kilogram, but general chemistry courses usually work in grams)
Volume: litre (L)
Time: second (s)
Metric prefixes scale these base units up or down. The ones you will use most: kilo- (k, x 1000), centi- (c, x 0.01), milli- (m, x 0.001). So 1 km = 1000 m, and 1 mL = 0.001 L.
Significant figures rules
All non-zero digits are significant.
Zeros between non-zero digits are significant (e.g. 205 has three sig figs).
Leading zeros are never significant (e.g. 0.0042 has two sig figs).
Trailing zeros after a decimal point are significant (e.g. 3.00 has three sig figs).
Trailing zeros in a whole number without a decimal point are ambiguous (use scientific notation to be clear).
Sig figs in calculations
Multiplication and division: the answer has the same number of sig figs as the factor with the fewest.
Addition and subtraction: the answer has the same number of decimal places as the measurement with the fewest.
Scientific notation expresses a number as a coefficient (between 1 and 10) multiplied by 10 raised to an exponent.
To convert a large number: move the decimal to the left. 2500 becomes 2.5 x 10^3.
To convert a small number: move the decimal to the right. 0.036 becomes 3.6 x 10^-2.
The exponent equals the number of places you moved the decimal. Left = positive exponent; right = negative.
Conversion factors are ratios equal to 1 that let you switch between units without changing the value of a measurement. The method (dimensional analysis) works by arranging fractions so unwanted units cancel.
Worked example: convert 130 lb to kg.
The conversion factor is 1 kg = 2.205 lb.
Set up: 130 lb x (1 kg / 2.205 lb).
The "lb" units cancel, leaving approximately 59 kg.
Dimensional analysis is the single most useful problem-solving tool in introductory chemistry. If your units do not cancel correctly, the setup is wrong.
Temperature scales
Three scales are used in chemistry: Fahrenheit (°F), Celsius (°C) and Kelvin (K).
Absolute zero: 0 K = -273 °C = -460 °F. This is the lowest possible temperature, where particle motion effectively stops.
Water freezes at 0 °C, 273 K, 32 °F.
Water boils at 100 °C, 373 K, 212 °F.
Kelvin is the SI unit. To convert: K = °C + 273.
Density
Density = mass / volume, typically reported in g/mL or g/cm3. If you know any two of mass, volume and density, you can calculate the third.
Specific gravity compares a substance's density to water's density (1.00 g/mL at roughly 4 °C). Because both densities share the same units, specific gravity is dimensionless. Numerically, it equals the density in g/mL.
Density is how you determine whether an object floats or sinks: if its density is less than the surrounding fluid, it floats.
\text{Density} = \frac{\text{mass}}{\text{volume}}K = \degree C + 273\degree F = \left(\frac{9}{5} \times \degree C\right) + 32\degree C = \frac{5}{9} \times (\degree F - 32)Scientific notation general form:
N \times 10^{n} \quad \text{where } 1 \leq N < 10Density is the principle behind float-or-sink sorting in recycling plants. Different plastics have different densities, and water baths separate them.
Conversion factors and dimensional analysis are how pharmacists calculate drug dosages. A wrong unit conversion in medicine can be life-threatening.
Significant figures matter in engineering tolerances: a bridge designed to 3 sig figs is not the same as one designed to 6.
Students often think boiling is a chemical change. It is not. Boiling is a physical change because the substance (water) stays H2O throughout.
Students confuse "no significant figures" with "zero." Leading zeros are placeholders, not measurements. 0.005 has one sig fig, not three.
Students assume that if something looks different after a change, it must be a chemical change. Tearing paper changes appearance but is physical; the cellulose molecules are unchanged.
Students mix up precision and accuracy. Precision is how close repeated measurements are to each other; accuracy is how close they are to the true value. You can be precise but inaccurate.
⚠️ Expect at least one question asking you to classify a change as physical or chemical. Look for whether new substances form.
⚠️ Sig fig questions appear on nearly every general chemistry exam. Practise until the rules are automatic.
⚠️ Dimensional analysis (conversion factor) problems are a staple. If units cancel to give the unit you want, the setup is correct.
⚠️ Know the three temperature scales and how to convert between them. The Kelvin conversion (K = °C + 273) is the most commonly tested.
⚠️ Be ready to classify a given sample as element, compound, homogeneous mixture or heterogeneous mixture.
True or false: Melting butter is a chemical change.
False. The butter changes state but its composition stays the same.
How many significant figures does 0.00340 have?
Three (the 3, 4 and trailing 0 after the decimal).
Fill in the blank: A substance made of two or more elements chemically bonded in a fixed ratio is called a ______.
Compound.
True or false: 0 K is the same as 0 °C.
False. 0 K = -273 °C.
Convert 2.50 km to metres.
2500 m (or 2.50 x 10^3 m).
Q: A student measures the mass of a sample as 12.110 g and its volume as 4.5 mL. What is the density, reported with the correct number of significant figures?
A: Density = 12.110 / 4.5 = 2.691... The answer should have 2 sig figs (limited by 4.5), so 2.7 g/mL.
Q: Classify each of the following as element, compound or mixture: (a) table salt, (b) oxygen gas, (c) bronze.
A: (a) Table salt (NaCl) is a compound. (b) Oxygen gas (O2) is an element. (c) Bronze is a mixture (an alloy of copper and tin).
Q: Convert 98.6 °F to Celsius and then to Kelvin.
A: °C = 5/9 x (98.6 - 32) = 5/9 x 66.6 = 37.0 °C. K = 37.0 + 273 = 310.0 K.
Q: How many significant figures are in each: (a) 0.0200, (b) 5010, (c) 3.00 x 10^4?
A: (a) Three. (b) Three (the trailing zero is ambiguous without a decimal point, but in most textbook conventions it is not counted). (c) Three.
Q: A block of metal has a mass of 575 g and displaces 65 mL of water. What is its density? Could it be iron (density ~7.87 g/mL)?
A: 575 / 65 = 8.8 g/mL (2 sig figs). This is higher than iron's density, so the block is likely a denser metal.
This material connects directly to stoichiometry, where you will use dimensional analysis constantly to convert between moles, grams and particles. The classification of matter (element vs. compound vs. mixture) comes back when you study solutions and concentration. Temperature and density reappear in gas laws and thermochemistry. Master the measurement and conversion tools here, because every later topic assumes fluency with them.
matter, states of matter, solid liquid gas, physical properties, chemical properties, physical change, chemical change, pure substance, element, compound, mixture, homogeneous, heterogeneous, metric system, SI units, significant figures, sig figs, scientific notation, dimensional analysis, conversion factors, unit conversion, temperature scales, Fahrenheit, Celsius, Kelvin, absolute zero, density, specific gravity, mass, volume, classification of matter