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

These revision notes cover Structure 1.1, Introduction to the particulate nature of matter, for IB Diploma Programme Chemistry. They follow the IB’s Diploma Programme Chemistry guide (published February 2023, for first assessment in 2025), understandings Structure 1.1.1 to Structure 1.1.3. Everything here is for SL and HL; the guide has no additional higher level content for this sub-topic. For full explanations and worked examples, use the Structure 1.1 study guide.

Other links: Structure 1.1 practice questions, the Structure 1 revision notes for the wider component, the IB DP Chemistry hub, the printable checklist and the free diagnostics.

Structure 1.1.1 Pure substances and mixtures

Definitions

• Element: one type of atom only; cannot be broken down chemically into simpler substances. Can be atoms (Ar), molecules (N₂, P₄) or a giant lattice (diamond).
• Compound: atoms of different elements chemically bonded in a fixed ratio. Separating it into elements needs a chemical reaction.
• Mixture: two or more elements and/or compounds, not chemically bonded, in no fixed ratio. Separable by physical methods.
• Homogeneous mixture: uniform composition, one phase (e.g. vinegar, bronze, filtered air).
• Heterogeneous mixture: non-uniform, two or more visible phases (e.g. muddy water, a suspension of flour in water).

Must-know distinctions

Element or compound	Mixture
Fixed composition	Variable composition
Sharp melting and boiling points	Melts and boils over a range
Compound’s properties differ from its elements’	Components keep their own properties
Chemical change needed to separate	Physical methods separate it

Quick reminders:

• Nitrogen and hydrogen in a cylinder form a mixture in any ratio, separable by liquefying and boiling off each gas. Ammonia, NH₃, made from them has a fixed 1:3 ratio of N to H atoms, a single boiling point, and properties neither gas has, such as a pungent smell and very high solubility in water.
• An alloy is a mixture, even though it has metallic bonding.
• Diatomic and polyatomic molecules of one element (H₂, O₃, S₈) are elements, not compounds.

Separation methods in one table

You want to…	Use	Key physical property
Get one solid into solution, leaving another	Solvation (dissolving)	Solubility in the solvent
Remove an insoluble solid from a liquid	Filtration	Solid vs dissolved particles
Get a dissolved solid back	Evaporation (or crystallization)	Solvent is volatile, solute is not
Purify an impure solid	Recrystallization	Solubility rises with temperature
Get pure solvent from a solution	Simple distillation	Large difference in boiling point
Separate miscible liquids	Fractional distillation	Different boiling points
Separate and identify small amounts of dissolved substances	Paper chromatography	Distribution between stationary and mobile phases

Method in steps: choosing a separation method

1. Identify what you want to keep: a solid, a liquid, or several dissolved substances.
2. Find the physical property that differs between the parts: solubility, particle size, boiling point or affinity for a solvent.
3. Is one part insoluble? Dissolve, then filter. Is the wanted part dissolved? Evaporate or crystallize.
4. Two liquids that mix? Fractional distillation. A small sample to identify? Chromatography.
5. Several steps are often combined, for example solvation, then filtration, then evaporation.

Method in steps: fractional distillation

1. Heat the mixture in a flask fitted with a fractionating column.
2. Vapour condenses and re-vaporizes many times up the column; the component with the lower boiling point reaches the top first.
3. The thermometer at the top stays steady while one pure component distils; collect each fraction in a separate receiver as the reading changes.

Method in steps: recrystallization

1. Dissolve the impure solid in the minimum of hot solvent.
2. Filter hot to remove insoluble impurities.
3. Cool so the product crystallizes; soluble impurities stay in solution.
4. Filter, wash with a little cold solvent, dry.

Method in steps: paper chromatography

1. Pencil base line; small, concentrated spots on it.
2. Solvent level below the base line; lid on the tank.
3. End the run with the solvent short of the top edge, and mark the furthest point it reached at once.
4. Rf = (base line to spot centre) ÷ (base line to solvent front), so it always lies between 0 and 1.
5. Same Rf as a reference in the same conditions suggests the same substance. A single spot suggests a pure substance (in that solvent).

Worked reminder: the solvent front moves 9.0 cm and a spot moves 2.7 cm, so Rf = 2.7 ÷ 9.0 = 0.30. A second spot at 8.1 cm has Rf = 0.90: it is more attracted to the mobile phase (the solvent) than the first spot is, so it travels further.

Structure 1.1.2 Kinetic molecular theory

	Solid	Liquid	Gas
Spacing	Touching, ordered	Touching, disordered	Far apart
Motion	Vibration only	Vibration, rotation, translation	Rapid, random translation
Shape; volume	Keeps own shape; own volume	Takes container’s shape; own volume	Fills whole container
Forces between particles	Strong	Weaker	Negligible

State symbols

Use (aq) for a solution in water, and (s), (l) or (g) for a solid, liquid or gas. Example: CaCO₃(s) + 2HNO₃(aq) → Ca(NO₃)₂(aq) + H₂O(l) + CO₂(g).

Changes of state

Name	Change	Endothermic or exothermic
Melting	s → l	endothermic
Freezing	l → s	exothermic
Vaporization (evaporation, boiling)	l → g	endothermic
Condensation	g → l	exothermic
Sublimation	s → g	endothermic
Deposition	g → s	exothermic

• Evaporation: surface only, below the boiling point, the most energetic particles escape, so the remaining liquid cools.
• Boiling: throughout the liquid, at a fixed temperature for a given external pressure.
• Evaporation can happen at any temperature at which the liquid exists; boiling happens only once the boiling point is reached.
• Iodine crystals warmed gently in a beaker give a purple vapour that forms crystals again on a cold surface: sublimation, then deposition.

Structure 1.1.3 Temperature and average kinetic energy

Key facts

• Temperature in kelvin is a measure of the average kinetic energy of the particles. Double the kelvin temperature, double the average kinetic energy.
• One kelvin = one Celsius degree in size. Only the zero point differs.
• 0 K is absolute zero. Kelvin temperatures cannot be negative and have no degree sign.

Conversion	Formula
Celsius → kelvin	T/K = T/°C + 273.15
Kelvin → Celsius	T/°C = T/K − 273.15

Worked reminder: a freezer at −18 °C is at −18 + 273.15 = 255.15 K. A temperature change of 30 °C is a change of 30 K: differences are the same on both scales.

Heating curve in brief

• Slope: energy raises average kinetic energy, so temperature rises.
• Plateau: energy overcomes the attractions between particles; average kinetic energy and temperature stay constant; two states coexist.
• Lower plateau = melting point; upper plateau = boiling point.
• Boiling plateau longer than melting plateau (more energy to separate particles fully).
• Cooling curve: same in reverse; a pure liquid freezes at one temperature, a mixture over a range.

Worked reminder: a pure solid heated steadily shows plateaus at 44 °C and 280 °C. Melting point = 44 + 273.15 = 317.15 K; boiling point = 280 + 273.15 = 553.15 K. At 100 °C the substance is a liquid, (l).

Quick self-test

1. Classify: (a) bronze; (b) P₄; (c) glucose; (d) oil and water shaken together.
2. Name the change of state when water vapour forms ice crystals directly on a cold surface.
3. Convert −40 °C to kelvin.
4. Convert 400 K to °C.
5. Which method separates water from a copper(II) sulfate solution, keeping the water?
6. A spot moves 4.8 cm when the solvent front moves 6.4 cm. Calculate Rf.
7. Give the state symbol for sodium chloride dissolved in water.
8. The kelvin temperature of a gas rises from 200 K to 600 K. By what factor does the average kinetic energy change?
9. Explain why the temperature is constant while a pure solid melts.
10. Why is a melting point measured over a range a sign of an impure sample?
11. Heating from 0 °C to 273 °C: by roughly what factor does the average kinetic energy increase?
12. Which method would you use to separate two miscible liquids that boil at 65 °C and 118 °C?

Answers

1. (a) homogeneous mixture (alloy); (b) element; (c) compound; (d) heterogeneous mixture.
2. Deposition.
3. −40 + 273.15 = 233.15 K (233 K).
4. 400 − 273.15 = 126.85 °C (127 °C).
5. Simple distillation.
6. 4.8 ÷ 6.4 = 0.75.
7. NaCl(aq).
8. 600 ÷ 200 = 3: it triples.
9. The energy supplied overcomes the attractions between particles instead of raising their average kinetic energy, so the temperature does not rise until all the solid has melted.
10. A pure substance melts at one sharp temperature. Impurities spread melting over a range (and usually lower it).
11. 273.15 K → 546.15 K, so the factor is 2 (it doubles).
12. Fractional distillation; the liquid boiling at 65 °C distils first.

Where marks are usually lost

• Describing a mixture as “chemically combined” or a compound as “two substances mixed”. Use the words bonded, fixed ratio and physical separation.
• Classifying an alloy as a compound, or O₂ as a compound.
• Mixing up homogeneous and heterogeneous, or giving an example without saying why (one phase vs more than one).
• Writing “dissolve in hot water” for recrystallization without the word minimum, or forgetting the cooling step.
• Drawing a chromatography base line in ink, or placing it below the solvent level.
• Measuring Rf to the end of the paper instead of to the solvent front.
• Using (aq) for a pure liquid, or leaving state symbols off a physical change.
• Saying the temperature rises during a change of state.
• Comparing kinetic energies with Celsius values; always convert to kelvin first.
• Writing “°K” or giving a negative kelvin temperature.

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.