IB DP Chemistry Structure 1.1: Introduction to the particulate nature of matter

This study guide covers Structure 1.1, Introduction to the particulate nature of matter, in IB Diploma Programme Chemistry. It is aligned to the IB’s Diploma Programme Chemistry guide (published February 2023, for first assessment in 2025), and teaches understandings Structure 1.1.1 to Structure 1.1.3. All of this content is for both SL and HL: the guide lists no additional higher level material for this sub-topic, and it allocates 2 teaching hours at both levels. It applies to examinations set under that guide.

For the whole of Structure 1 at a glance, see the Structure 1 overview. The IB DP Chemistry hub lists every resource for the course, the printable checklist lets you tick off each understanding, and the free diagnostics help you find weak spots across your subjects.

What this unit covers

Guide reference	What you must be able to do	Level
Structure 1.1.1	Tell elements, compounds and mixtures apart by their properties; explain homogeneous and heterogeneous mixtures; choose and describe solvation, filtration, recrystallization, evaporation, distillation and paper chromatography	SL and HL
Structure 1.1.2	Use the kinetic molecular theory to describe solids, liquids and gases; write state symbols (s), (l), (g) and (aq); name every change of state	SL and HL
Structure 1.1.3	Link temperature in kelvin to the average kinetic energy of particles; read heating and cooling curves; convert between °C and K	SL and HL

Structure 1.1.1 Elements, compounds and mixtures

The three kinds of matter

• An element contains atoms of just one kind, all with the same atomic number. No chemical reaction can split it into simpler substances. Examples: neon, O₂, sulfur (S₈), copper.
• A compound contains atoms of two or more different elements that are chemically bonded in a fixed ratio. Water is always two hydrogen atoms to one oxygen atom. Getting the elements back out requires a chemical reaction.
• A mixture is made of two or more elements or compounds that are not chemically bonded to each other, in no fixed ratio. Because there are no bonds between the components, physical methods can separate them.

Elements and compounds are both pure substances. A mixture is not.

How their properties differ

Property	Element or compound	Mixture
Composition	Fixed	Variable
Melting and boiling points	Sharp, at one temperature (at a given pressure)	Usually over a range of temperatures
Properties compared with the parts	A compound’s properties are quite different from those of its elements	Each component keeps its own properties
Separation	A compound needs a chemical reaction	Physical methods are enough

Take hydrogen and oxygen. A balloon of both gases is a mixture in any proportions. Ignite it and they react to form water, a compound that is liquid at room temperature and is always H₂O.

Homogeneous and heterogeneous mixtures

• A homogeneous mixture has the same composition throughout and only one phase. You cannot see separate parts. Examples: air, a sugar solution, brass.
• A heterogeneous mixture does not have a uniform composition. You can see (or detect with a microscope) more than one phase. Examples: sand in water, oil floating on vinegar, granite.

Alloys such as brass are treated as mixtures: despite their metallic bonding, the proportions of the metals can vary.

Worked example 1: classifying substances

Classify each of the following: (a) ozone, O₃; (b) magnesium oxide; (c) seawater; (d) chalk powder stirred into water.

1. (a) O₃ contains only oxygen atoms: an element, even though it is molecular.
2. (b) Magnesium and oxygen are chemically bonded in a fixed 1:1 ratio, so magnesium oxide is a compound.
3. (c) Seawater is water with dissolved salts in variable amounts. It is uniform throughout: a homogeneous mixture.
4. (d) Chalk does not dissolve, so solid particles are suspended in the liquid and settle out: a heterogeneous mixture.

Separation methods

Each method uses a difference in a physical property between the components.

Method	Property it uses	What it does
Solvation (dissolving)	Solubility in a chosen solvent	Solvent particles surround the particles of a solute, taking one component into solution and leaving others behind
Filtration	Particle size: solid vs dissolved	Holds back an insoluble solid (residue) while the solution passes through (filtrate)
Evaporation	Volatility	Removes a solvent to leave a dissolved solid behind
Recrystallization	Solubility changing with temperature	Purifies a solid: dissolve in the minimum of hot solvent, cool, collect the crystals; soluble impurities stay in solution
Distillation	Boiling point	Simple distillation recovers a solvent from a solution; fractional distillation separates miscible liquids with different boiling points
Paper chromatography	Distribution between a stationary phase and a mobile phase	Separates and identifies the components of a small sample, such as a mixture of dyes

Worked example 2: purifying a reaction product by recrystallization

A chemist makes benzoic acid. The crude solid contains some sand and a small amount of a coloured impurity that is soluble in water. Benzoic acid is much more soluble in hot water than in cold water. Describe how to obtain pure, dry crystals.

1. Add just enough hot water to dissolve the crude solid (solvation), so that as much product as possible crystallizes later.
2. Filter while hot to remove the sand, so benzoic acid does not crystallize in the funnel.
3. Cool the filtrate slowly. Benzoic acid’s solubility falls and crystals form; the small amount of coloured impurity stays dissolved.
4. Filter to collect the crystals, then wash them with a little cold water to remove the solution left on their surface.
5. Dry the crystals, for example in a warm oven or a desiccator.

A pure sample melts sharply, so a melting point checks purity.

Worked example 3: paper chromatography

Pigments extracted from a flower petal are spotted on a pencil base line on chromatography paper and run in a solvent. The solvent front travels 6.8 cm. Two spots appear, at 2.04 cm and 4.42 cm from the base line. Find each retardation factor, Rf.

1. Rf = (base line to centre of spot) ÷ (base line to solvent front).
2. First spot: 2.04 ÷ 6.8 = 0.30.
3. Second spot: 4.42 ÷ 6.8 = 0.65.
4. Compare with known pigments run in the same solvent and paper. A matching Rf suggests the same substance.

Draw the base line in pencil (ink would separate) and keep the spots above the solvent level.

Structure 1.1.2 Kinetic molecular theory and changes of state

The kinetic molecular theory models matter as particles (atoms, molecules or ions) that are always moving. The strength of the attractions between the particles, compared with their kinetic energy, decides the state.

	Solid	Liquid	Gas
Arrangement	Ordered lattice, particles touching	Particles touching but disordered	Widely spaced, no order
Movement	Vibration about a fixed point	Particles slide past each other	Fast, random straight-line motion between collisions
Shape and volume	Fixed shape, fixed volume	Takes the shape of its container, fixed volume	Fills its container
Compressibility	Almost none	Very little	Easy, because of the large spaces
Attraction between particles	Strongest	Intermediate	Very weak

State symbols

Add a state symbol after every formula in an equation:

• (s) solid, (l) liquid, (g) gas
• (aq) aqueous: dissolved in water

Example: Zn(s) + CuSO₄(aq) → ZnSO₄(aq) + Cu(s). Physical changes take state symbols too, for example Br₂(l) → Br₂(g).

Names of the changes of state

Process	What happens	Energy
Melting	a solid becomes a liquid	endothermic
Freezing	a liquid becomes a solid	exothermic
Vaporization	a liquid becomes a gas	endothermic
Condensation	a gas becomes a liquid	exothermic
Sublimation	a solid becomes a gas directly, with no liquid stage	endothermic
Deposition	a gas becomes a solid directly, with no liquid stage	exothermic

Vaporization has two forms. Evaporation happens only at the surface of a liquid and at any temperature: the fastest-moving particles escape. Boiling happens throughout the liquid, at one temperature for a given pressure, with bubbles of vapour forming inside it.

Changes that separate particles need energy to overcome the attractions between them, so they are endothermic; the reverse changes release energy.

Frost forming on a cold windscreen is deposition; solid carbon dioxide turning straight into gas is sublimation.

Structure 1.1.3 Temperature and kinetic energy

Temperature in kelvin

Temperature’s SI unit is the kelvin, symbol K. Temperature in kelvin is a measure of the average kinetic energy of the particles in a sample: if the kelvin temperature doubles, the average kinetic energy doubles. Not every particle has the same energy; the temperature reflects the average.

One kelvin is the same size step as one degree Celsius, so converting only shifts the zero:

• T (K) = T (°C) + 273.15
• T (°C) = T (K) − 273.15

Many questions round this to 273. Kelvin values are written without a degree sign and are never negative. 0 K is absolute zero.

Worked example 4: converting temperatures

Use these values: nitrogen boils at −196 °C; tungsten melts at 3422 °C; gallium melts at 302.9 K.

1. Nitrogen: −196 + 273.15 = 77.15 K (77 K to the nearest kelvin).
2. Tungsten: 3422 + 273.15 = 3695.15 K (3695 K).
3. Gallium: 302.9 − 273.15 = 29.75 °C, so it would melt in your hand.

Worked example 5: does the average kinetic energy double?

A sealed sample of gas is heated from 27 °C to 327 °C. Does the average kinetic energy of the particles double? What if it is heated from 10 °C to 20 °C instead?

1. Convert to kelvin: 27 °C = 300.15 K; 327 °C = 600.15 K.
2. Ratio: 600.15 ÷ 300.15 = 2.00 (to 3 significant figures). So yes, the average kinetic energy doubles (very nearly).
3. For 10 °C → 20 °C: 283.15 K → 293.15 K. Ratio = 293.15 ÷ 283.15 = 1.035.
4. The average kinetic energy rises by only about 3.5%, although the Celsius number doubled. Use kelvin for ratios.

Heating and cooling curves

When a pure solid is heated steadily, the temperature rises, then stays constant while it melts, rises again, then stays constant while it boils.

• Sloping sections: the energy supplied increases the average kinetic energy of the particles, so the temperature rises.
• Flat sections (plateaus): the energy overcomes the attractions between particles, raising their potential energy, not their kinetic energy. No temperature rise happens until every particle has changed state.
• The boiling plateau is usually longer, because separating particles into a gas needs more energy than loosening them into a liquid.

A cooling curve shows the same features in reverse, with plateaus at the condensation and freezing temperatures. A mixture shows no sharp plateau: it melts or freezes over a range of temperatures. That is why a melting point is a quick test of purity.

Worked example 6: reading a heating curve

A pure substance is heated from −20 °C at a steady rate. The temperature rises to 17 °C and stays there for 3 minutes, rises to 118 °C and stays there for 11 minutes, then rises again. Use these data.

1. The first plateau is the melting point: 17 °C = 290.15 K. Solid and liquid are both present along it.
2. The second plateau is the boiling point: 118 °C = 391.15 K. Liquid and gas are both present.
3. At 50 °C the substance is a liquid, so its state symbol would be (l).
4. The boiling plateau is longer because vaporization needs more energy than melting.

Common errors

• Calling O₂ or O₃ a compound because it is a molecule. A molecule of one element is still an element.
• Calling an alloy a compound. It is a mixture, despite the metallic bonding.
• Writing (aq) for a pure liquid such as ethanol. Use (l); (aq) means dissolved in water.
• Confusing evaporation (surface, any temperature) with boiling (throughout the liquid).
• Saying temperature rises during melting. It stays constant on the plateau.
• Calculating a ratio of kinetic energies from Celsius values.

Where next

Next, revise from the revision notes, then try the practice questions.

Official syllabus

International Baccalaureate Organization, Diploma Programme Chemistry guide, first assessment 2025 (published February 2023): Structure 1.1, Introduction to the particulate nature of matter.