Comprehensive notes: Cambridge IGCSE Chemistry Coursebook — Chapters 1 & 2 (States of Matter and Atomic Structure)

Chapter 1: States of matter

  • Chemistry is a laboratory science with broad real‑world relevance (environment, health, resources).
  • This coursebook provides full syllabus coverage for Cambridge IGCSE Chemistry (0620/0971), including practicals, examinations, and real-world examples.
  • Features explained in the book:
    • Learning intentions, Getting started, Science in context, Experimental skills, Questions, Activity, Key words, Command words, Supplement content, Worked examples, Reflections, Summary, Self‑evaluation, Project, Exam‑style questions, and a glossary and index (as described in the contents).
  • Practical emphasis: safety guidance is included for investigations; teachers are responsible for safety compliance.

1.1 What is matter? States of matter

  • Matter: anything that occupies space (volume) and has mass.
  • States of matter: solid, liquid, gas.
  • Substances can exist in three states depending on temperature and pressure.
  • Key differences between states (Table 1.1 in the book):
    • Solid: definite shape, fixed volume, high density; not fluid.
    • Liquid: definite volume, takes shape of container, fluid.
    • Gas: no fixed volume, fills container, fluid, low density.
  • Fluids: liquids and gases can flow and be poured.
  • All three states expand with increasing temperature; contracting with decreasing temperature (more pronounced in gases).
  • Gas volume is highly affected by pressure: gases are easily compressed; liquids are only slightly compressible; solids are effectively incompressible.
  • The three states differ in particle arrangement and motion (kinetic particle theory).
  • Particles:
    • Solids: tightly packed, vibrate about fixed positions in a lattice.
    • Liquids: less tightly packed, can move past each other.
    • Gases: far apart, move randomly and rapidly.
  • Intermolecular space decreases from gas to liquid to solid; gases have large intermolecular spaces and are easily compressed.
  • The concept of pressure arises from particle collisions with container walls.
  • Diffusion: spreading of particles to fill space; faster in gases, slower in liquids, negligible in solids.
  • Energy changes during state changes: energy is absorbed to overcome intermolecular forces; energy is released when bonds form.

1.2 Changes of state and energy concepts

  • Sublimation: solid to gas (and gas to solid) bypassing the liquid phase (e.g., CO₂ at atmospheric pressure).
  • Evaporation: liquid to gas at the surface; rate increases with surface area and temperature.
  • Boiling: liquid to gas throughout the liquid when vapor pressure equals atmospheric pressure; occurs at a specific boiling point for a pure liquid.
  • Melting and freezing: melting point (m.p.) and freezing point (f.p.) are the same for a pure substance; e.g., water m.p. = 0°C and f.p. = 0°C.
  • Impurities affect melting/boiling points; impure substances melt/boil over a range of temperatures.
  • Pure substances have precise mp and bp; the data can be used for testing purity and identity.
  • Heating/cooling curves: temperature remains constant during melting/freezing (A–B–C regions); energy is absorbed during melting/evaporation and released during freezing/condensation.
  • Energy curves illustrate endothermic (positive ΔH) vs exothermic (negative ΔH) changes.
  • Practical demonstration: the melting point apparatus (melting point tube) for solids; cooling curves can be constructed by heating and monitoring temperature as substance cools.
  • Table and figure references present the relationship between mp, bp, and phase behavior.

1.3 Mixtures, solutions, and diffusion in solutions

  • Mixtures vs compounds: mixtures contain two or more substances not chemically bonded; compounds are chemically bonded combinations with fixed proportions.
  • Solute vs solvent: in a solution, the solid that dissolves is the solute; the liquid that dissolves the solute is the solvent.
  • Solubility and dissolution:
    • Some solids dissolve in liquids to form solutions (soluble); others do not (insoluble).
    • Some liquids mix completely (miscible), e.g., ethanol and water; others do not (immiscible).
  • Salts and solutions: seawater is impure water; impurities alter mp and bp of solutions; salt residues can be observed after evaporation of seawater.
  • Solubility depends on temperature: most solid solubility increases with temperature; crystallisation occurs when a saturated solution is cooled.
  • Diffusion in liquids: ions or molecules diffuse through water; diffusion rate in liquids is slower than in gases.
  • Diffusion of gases: gas molecules diffuse rapidly to fill available space; diffusion rate depends on molecular mass (lighter molecules diffuse faster).
  • Gas diffusion experiments (ammonia vs. hydrogen chloride) illustrate the inverse relationship between molecular mass and diffusion speed.
  • Practical diffusion experiments: silver nitrate and potassium iodide diffusion in water; formation of precipitate indicates diffusion and reaction; safety considerations discussed.
  • Gas diffusion through porous barriers: diffusion through a porous pot shows differences when external gas masses differ (e.g., hydrogen outside vs air outside).
  • Energy interactions during phase changes in diffusion processes are connected to intermolecular forces.

1.4 Experimental skills and activities (Chapter 1 related)

  • Plotting cooling curves requires careful temperature measurement and timing; analysts discuss reliability and improvements.
  • Self-evaluation and reflection tasks are included to develop metacognitive skills in data analysis and graphing.
  • Projects encourage real-world links (e.g., water on Earth and space missions; Goldilocks Zone; exoplanets).