Complete Course: iGCSE Chemistry & Middle School Chemistry

Solids exhibit a fixed shape and volume; liquids take the shape of their container and have fixed volume; gases have no fixed shape or volume and can be compressed.

Learn how solids melt to liquids and liquids boil or evaporate into gases, while cooling causes condensation and freezing; some substances sublimate directly from solid to gas.

Explore how changes in state between solid, liquid, and gas occur with heat, using melting and boiling points to determine a substance's state, with water and Bromine as examples.

Explore heating and cooling curves, identifying phase changes at melting and boiling points, and how kinetic energy and attractive forces drive transitions between solid, liquid, and gas.

Explore how rising temperature expands gas volume at constant pressure, while in a fixed-volume sealed container, heating increases pressure.

Explore how increasing temperature makes gas particles move faster, expanding volume when allowed or raising pressure when volume is fixed due to more frequent wall collisions.

Diffusion is the process where particles move from high to low concentration, becoming uniform in liquids and gases without energy input, with gases diffusing fastest and solids diffusion being negligible.

Learn how temperature and molecular mass affect diffusion in gases, with lighter molecules diffusing faster, and watch HCl and ammonia form ammonium chloride where their vapors meet.

Define elements as pure substances that cannot be broken down chemically, explain compounds as two or more elements in ratios, and describe mixtures as not chemically bound, water and sand.

Explore protons, neutrons, and electrons; atoms have the nucleus and electron shells. Most mass lies in the nucleus, and atoms stay neutral when protons equal electrons.

Learn how the atomic number, or proton number, equals the number of protons, and how the mass number, or nucleon number, totals protons and neutrons in the atom.

Explore isotopes: atoms of the same element with the same protons but different neutrons, illustrated by chlorine 35 and 37, and how relative atomic mass reflects isotope variation.

Examine isotopes with the same protons and electrons but different neutrons, why chemical properties stay similar, and how relative atomic mass is calculated from isotope abundances using carbon-12.

Learn to write electronic configuration by counting protons and electrons from the atomic number, then fill shells around the nucleus with the 2, 8, 18 rules, as in carbon's 2,4.

Explore how the periodic table organizes elements by atomic number into periods and groups, and how electronic configuration determines shells, valence electrons, and noble gas stability.

Investigate chemical bonding by revisiting group eight noble gases, noting helium's duplet and the stable octet that prompts other elements to gain, lose, or share electrons.

Learn how atoms form ions by losing or gaining electrons to reach a noble gas configuration, then see how cations and anions build ionic compounds held together by electrostatic attraction.

Learn how metals lose electrons to form plus-one to plus-three ions, while nonmetals gain electrons to form negative ions, reaching noble gas configurations; group four and noble gases behave specially.

Explore how ion charges vary, including zinc plus two, silver plus one, copper one or two, iron two or three, and chromium two or three, with roman numerals.

Learn how to write ion formulas with number first followed by the sign and to ignore 1, then name ions by metal versus nonmetal rules, using ide endings for nonmetals.

Learn how to name and write the formulae for ionic compounds by matching cations and anions, using the cross method, and applying brackets for polyatomic ions.

Download the worksheet in this lesson and practice writing the electronic configuration of ions by filling in the blanks, preparing you for the next lesson.

Explore writing ion formulas and names, and determine electronic configurations for ions like hydrogen, oxide ion, phosphate ion, and chloride ion to achieve stable duplet or octet.

Learn to draw dot and cross diagrams for ionic compounds by determining ion types from formula, writing their electronic configurations, and showing charge and electron transfer in sodium chloride.

Explore the structure of ionic compounds as giant lattice crystals, like sodium chloride, with alternating cations and anions in a regular cubic arrangement.

Explore formation and properties of ionic compounds, where metals form cations and nonmetals form anions; they are solids with high melting points and conduct electricity only when aqueous or molten.

Ionic compounds have high melting points due to electrostatic forces between ions, and they conduct electricity when molten or in aqueous solution as the ions move, unlike in solid state.

Explore covalent bonding and covalent molecules, including how sharing electrons achieves stable duplet or octet configurations, and distinguish simple molecules from macromolecules and giant covalent molecules.

Learn to deduce covalent formulas from names, memorize common formulas like water (H2O) and carbon dioxide (CO2), and use mono, di, tri to count atoms.

Draw dot-and-cross diagrams for simple molecules such as hydrogen, chlorine, water, methane, ammonia, and hydrogen chloride to show covalent bonding and electron sharing.

Draw dot-and-cross diagrams for simple molecules, illustrating single, double, and triple bonds in methanol, ethene, oxygen, carbon dioxide, and nitrogen by sharing electrons to form duplet and octet configurations.

Explore the properties of simple molecules, including low melting and boiling points and poor electrical conductivity, with chlorine as an example, contrasted with ionic compounds.

Simple molecules have melting and boiling points due to weak intermolecular forces. Melting or boiling separates molecules, not covalent bonds; electrons stay localized, so they do not conduct electricity.

Explore giant covalent molecules, where countless covalent bonds create endless structures such as diamond and graphite. Learn why their melting and boiling points are high due to extensive covalent bonding.

Explore diamond’s giant covalent network where carbon forms four bonds. This structure yields extreme hardness, a high melting and boiling point, and no electrical conduction.

Graphite forms a giant covalent structure of carbon with each atom covalently bonded to three others in layered hexagonal sheets held by weak interlayer forces, making it slippery and conductive.

Explore carbon allotropes by comparing graphite and diamond, showing how hexagonal versus tetrahedral structures yield graphite's conductivity and diamond's hardness.

Silicon dioxide forms a three-dimensional tetrahedral network, with silicon–oxygen covalent bonds, making it very hard, with high melting and boiling points, and does not conduct electricity.

Understand metallic bonding as the electrostatic attraction between positively charged metal ions and delocalized valence electrons, which lets metals conduct electricity and heat and be malleable and ductile.

Learn to write formulas for metals by their symbols, diatomic molecules, covalent molecules, and ionic compounds by balancing charges using the periodic table.

Deduce the molecular formula from structures by counting atoms and identifying elements (C, H, O, N), then write subscripts for each element, omitting ones when the count is one.

Learn how to determine the empirical formula by converting a molecular formula to the simplest whole-number ratio, using examples like C2H4O2 to CH2O and P2H4O6 to PH2O3.

Distinguish physical changes, where substances retain their chemical identity (dissolution, distillation, phase changes of water as H2O), from chemical changes that form new compounds like magnesium oxide.

Learn to write word equations that depict chemical reactions by placing reactants on the left, products on the right, and using the arrow to separate them.

Convert word equations to balanced chemical equations by writing formulas for reactants and products, then balance atoms and memorize key formulas for metals, nonmetals, and ionic compounds.

Identify state symbols for solid, liquid, aqueous, and gas, and apply them to writing and balancing chemical equations, with sodium reacting with water forming hydrogen gas and aqueous sodium hydroxide.

Learn to write ionic equations by balancing equations, identifying ionic compounds and acids in aqueous, splitting them into ions, canceling spectator ions, and obtaining the simplest ionic form.

Learn how relative atomic mass and relative molecular mass compare to one twelfth of carbon-12, and apply to water, NaCl (relative formula mass), MgSO4, and NaOH.

Use mass ratios from the balanced equation to find product mass without mole calculations, illustrated by 12 g Mg forming 20 g MgO via 48:32:80.

Explore the mole concept and Avogadro constant, linking moles to particles using 6.02×10^23, and learn formulas to convert between moles and number of particles.

Calculate moles from mass by dividing by molar mass (g per mole); use Mr. for relative mass. For 16 g O2, 0.5 mol, and 2 mol CO2 weighs 88 g.

At RTP, equal moles of gas occupy the same volume, illustrating the ideal gas law. Use 24 dm³/mol to convert volumes and moles, with 48 cm³ becoming 0.048 dm³.

Compute concentration by mass per volume or moles per volume. Use solute, solvent, and solution concepts, and apply molar mass to find the number of moles.

Learn to calculate concentration in moles per dm^3 and grams per dm^3 using molar mass, applying these formulas to NaCl problems like 3 g in 0.1 dm^3.

Convert mass to moles and moles to mass with molar mass, apply n = volume over molar volume for gases only, and learn concentration in solutions in two forms.

Explore stoichiometric ratios in balanced equations to determine mole ratios of reactants and products, then apply them to convert mass to moles and predict product masses.

Learn how stoichiometric ratios become volume ratios for gases at the same temperature and pressure, because the mole ratio equals the volume ratio for gases only.

Apply the percentage purity formula, mass of pure substance over total mass times 100, using mole ratios and molar masses, with sodium chloride and magnesium oxide examples.

Explore how to calculate percentage yield using actual yield over theoretical yield and mole calculations, with real-world deviations due to side reactions when calcium carbonate converts to calcium oxide.

Learn to calculate percentage composition by mass from molar mass, using calcium carbonate and ammonium nitrate as examples, and compute carbon, oxygen, calcium, and nitrogen percentages.

Explain empirical vs molecular formulas and derive empirical formula from mass data, using phosphorus oxide and hydrocarbon examples with mole ratios and molar mass to obtain the molecular formula.