Chemistry lessons

248 lessons

• Acids and Bases: Properties, Indicators and Reactions – An acid gives H⁺ ions in water and a base gives OH⁻ ions. Indicators show which one is present by a colour or smell change, and when H⁺ meets OH⁻ they make water, leaving a salt behind.
• Acids, Bases, Salts and Titration – Inorganic compounds fall into four big classes: oxides, acids, bases and salts. They are linked: a metal makes a basic oxide, that makes a base, and a base plus an acid makes a salt and water. A titration uses this neutralisation to measure an unknown concentration: base of known concentration is added from a burette until all the acid is used up (the equivalence point). Then n(acid) = n(base) by the reaction ratio, so C₁V₁ = C₂V₂ for a 1 : 1 reaction.
• Advanced Chemistry: Structure, Properties, Energy and Equilibrium – Advanced chemistry links four big ideas. Structure: atoms join in fixed shapes. Properties: the shape and charge spread decide melting point, solubility and more. Energy: reactions need a push (activation energy) and release or take heat. Equilibrium: reversible reactions settle where forward and backward rates are equal.
• Alcohols – An alcohol is a carbon compound with an –OH group on an sp³ carbon (R–OH). Count the carbons joined to the C–OH carbon to classify it as primary (1°), secondary (2°) or tertiary (3°). Name it as alkan-n-ol, giving –OH the lowest number. Make alcohols from alkenes (acid-catalysed hydration or hydroboration–oxidation), by reducing aldehydes, ketones and acids, or from Grignard reagents. H-bonding gives them high boiling points and water solubility. Reactions break either the O–H bond (acidic: Na, esterification) or the C–O bond (HX, dehydration), and 1°/2° alcohols can be oxidised. Methanol and ethanol are the key commercial alcohols.
• Aldehydes and Ketones: Structure, Preparation and Reactions – Aldehydes and ketones both have a carbonyl group, C=O. In an aldehyde the carbonyl carbon has at least one H on it (−CHO). In a ketone it has two carbon groups (−CO−). The C=O bond is polar: carbon is slightly positive, oxygen slightly negative. So electron-rich particles (nucleophiles) attack the carbon. This one idea explains most of their reactions: addition of HCN, NaHSO₃, alcohols and ammonia derivatives. Aldehydes are easy to oxidise (Tollens and Fehling tests), ketones are not. A hydrogen on the carbon next to C=O (α-H) is slightly acidic, which gives aldol condensation. Aldehydes with no α-H give the Cannizzaro reaction instead.
• Alkali Metals and Sodium Compounds – Alkali metals (lithium, sodium, potassium, rubidium, caesium) are group 1 of the periodic table. Each atom has one outer electron that it loses easily, forming M⁺ ions. They are soft, light and very reactive: they react with water to make hydrogen and an alkali, and they burn in air. Reactivity rises down the group. Sodium gives Na₂O, Na₂O₂, Na₂CO₃ (washing soda) and NaHCO₃ (baking soda), and each metal colours a flame in its own way.
• Alkanes – Alkanes are hydrocarbons with only single C–C bonds. Their general formula is CnH2n+2. We make them by adding hydrogen to alkenes, by the Wurtz reaction, by reducing alkyl halides and by decarboxylation. They are non-polar, burn with a lot of heat and swap H for Cl in light. Around the C–C bond the molecule can twist into staggered and eclipsed shapes called conformations.
• Alkenes – Alkenes are hydrocarbons with one C=C double bond and the formula CnH2n. The double bond is one σ bond plus one π bond, so the molecule is flat around it and cannot twist. That gives cis and trans isomers. We make alkenes by removing small molecules (H₂O, HX, X₂) or by partly adding H₂ to alkynes. Alkenes mainly do addition reactions: H₂, X₂, HX (Markovnikov rule, or anti-Markovnikov with peroxide), water, ozone, KMnO₄ and polymerisation.
• Alkynes – Alkynes have a C≡C triple bond and the formula CnH2n−2. The two triple-bond carbons are sp hybridised, so the molecule is a straight line. A hydrogen on a triple-bond carbon is weakly acidic because an sp carbon holds electrons tightly: strong bases like NaNH₂ remove it. The triple bond can open twice, so alkynes add two molecules of H₂, X₂ or HX. With water (Hg²⁺ catalyst) they give aldehydes or ketones, and three ethyne molecules can join into benzene.
• Allotropes of Carbon: Diamond, Graphite, Graphene and Fullerenes – Allotropes are different forms of the same element in the same state. Carbon has several. In diamond each atom makes 4 strong covalent bonds in a giant 3D network, so it is very hard, has a very high melting point and does not conduct electricity. In graphite each atom bonds to 3 others in flat layers of hexagons; the layers slide (soft, slippery) and one spare electron per atom is free to move, so graphite conducts. Graphene is a single graphite layer: strong, light and an excellent conductor. Fullerenes such as C₆₀ are hollow cages; nanotubes are rolled-up tubes. All of them burn in oxygen to make carbon dioxide.
• Aluminium and Its Amphoteric Oxide and Hydroxide – Aluminium (Al, atomic number 13) is a light, silvery metal with 3 outer electrons, so it forms Al³⁺. In air it grows a thin, tough skin of Al₂O₃ that protects it. Both Al₂O₃ and Al(OH)₃ are amphoteric: they react with acids to make aluminium salts and with strong alkalis to make aluminate salts. The metal itself also gives hydrogen with both dilute acid and hot alkali.
• Amines: Structure, Naming, Preparation, Basicity and Reactions – An amine is ammonia with one, two or three H atoms swapped for carbon groups. The N keeps a lone pair, so amines are bases and nucleophiles. Count the C atoms on N to get 1°, 2° or 3°. Make them by reducing nitro compounds, nitriles or amides, from alkyl halides + NH₃, by Gabriel (1° only) or Hofmann bromamide (one C fewer). Alkyl groups push electrons and raise basicity; the ring in aniline pulls the pair away and lowers it. Key reactions: salts with acids, acylation, carbylamine test (1°), Hinsberg test (1°/2°/3°), nitrous acid and ring substitution in aniline.
• Amorphous Materials: Oxide Glasses and Glass-Ceramics – An amorphous solid has no repeating pattern. Oxide glass is a random net of corner-sharing SiO₄ units (network former). Soda (Na₂O) and lime (CaO) are modifiers: they break links and lower the softening temperature. Glass has no sharp melting point; it softens slowly through the glass transition. A glass-ceramic is glass heated in a controlled way so that millions of tiny crystals grow inside it, giving strength and low thermal expansion.
• Applied Chemistry: How Chemists Test and Use Materials – Applied chemistry uses chemistry to solve real jobs: finding out what is in a sample, checking quality, making useful materials, and cleaning or treating things with living cells. Chemists take a small fair sample, measure it exactly (titration, chromatography), and then use the answer to decide if a product is safe or good.
• Atoms and Molecules – In a chemical reaction mass is neither created nor destroyed (conservation of mass), and a compound always has its elements in the same ratio by mass (constant proportions). Dalton explained both: matter is made of tiny atoms that join in small whole numbers. Atoms join to form molecules; charged atoms or groups are ions. Formulae are written by crossing valencies. Molecular mass (or formula unit mass for ionic compounds) is the sum of the atomic masses in the formula, in u.
• Avogadro's Law: Equal Volumes, Equal Numbers of Molecules – Avogadro's law: at the same temperature and pressure, equal volumes of all gases hold the same number of molecules. So volume is proportional to the number of moles: V ∝ n, or V₁/n₁ = V₂/n₂. At STP (0 °C, 101.3 kPa) one mole of any gas fills 22.4 L (molar volume, Vm). Useful links: n = V ÷ Vm, N = n × 6.02 × 10²³, m = n × M. Equal volumes have different masses, so relative density D = M₁ ÷ M₂ (D by hydrogen = M ÷ 2; by air = M ÷ 29). In gas reactions, volumes are in the same ratio as the equation coefficients (2H₂ + O₂ → 2H₂O: 2 : 1 : 2).
• Baking and Fast Food: How Bread and Cakes Rise – Fast food is food prepared and served quickly using standard recipes, portions and steps, such as burgers, pizza and wraps. Baking cooks food with dry heat in an oven. Flour gives structure through gluten, liquids bind, raising agents (yeast, baking powder, baking soda, steam, air) create gas, sugar sweetens and browns, salt controls yeast and adds taste, and fat makes products tender. In the oven, gas expands, starch and protein set and the surface browns.
• Balancing Redox Reactions: Oxidation Number and Half-Reaction Methods – In a balanced redox equation, three things must match: atoms, charge, and electrons. The electrons lost by the reducing agent must equal the electrons gained by the oxidising agent. The oxidation number method balances the change in numbers first; the half-reaction method splits the reaction into two halves and joins them. O is balanced with H₂O, and H with H⁺ in acid or with OH⁻ in base.
• Basics of Organic Reaction Mechanism – A reaction mechanism is the step-by-step story of which bonds break, which form, and where the electrons move. A covalent bond can break in two ways. In homolysis each atom takes one electron and free radicals form. In heterolysis one atom takes both electrons, giving a carbocation (C with + charge) and a carbanion or an anion. Electron-poor species that seek electrons are electrophiles; electron-rich species that give electrons are nucleophiles. The movement of electrons inside a molecule is controlled by four effects: the inductive effect (pull along σ bonds), the resonance effect (spreading of π electrons), the electromeric effect (a temporary full shift when a reagent attacks) and hyperconjugation (σ C–H electrons spreading into a nearby empty p orbital or π bond). These effects decide how stable an intermediate is and where a reagent attacks.
• Benzene and Aromaticity – Benzene (C₆H₆) is a flat ring of six sp² carbons. Its six π electrons are spread over the whole ring, so all C–C bonds are equal and the ring is extra stable. A ring like this is called aromatic when it is cyclic, planar, fully conjugated and has 4n + 2 π electrons (Hückel rule). Benzene prefers electrophilic substitution (nitration, halogenation, sulphonation, Friedel–Crafts) over addition, because substitution keeps the stable ring. A group already on the ring decides where the next group goes: ortho/para or meta. Benzene and fused-ring hydrocarbons are toxic and can cause cancer.
• Biochemistry: The Chemistry of Life – Plants build glucose by photosynthesis: 6 CO2 + 6 H2O + light gives C6H12O6 + 6 O2. Four families of big molecules (carbohydrates, proteins, fats, nucleic acids) are built from small units. Cells break nutrients down to release energy, with enzymes speeding every step. Proteins are made by reading DNA, copying it to mRNA and joining amino acids on a ribosome.
• Biologically Active Compounds – Our body and our medicines are made of organic compounds that act on living things. Sugars (glucose, sucrose, starch, cellulose) give energy or structure, fats store energy, proteins build and run the body, medicines fit receptors to treat illness, and vitamins in small amounts prevent deficiency diseases.
• Bond Parameters, Resonance and Polarity – A covalent bond is described by four numbers: bond length (distance between nuclei), bond angle (angle between bonds at an atom), bond enthalpy (energy to break 1 mol of bonds) and bond order (number of shared pairs). Higher bond order means a shorter, stronger bond. When one Lewis structure cannot describe a molecule, the real molecule is a resonance hybrid of several structures. Unequal sharing makes a bond polar; the dipole moment μ = q × d measures it, and the shape decides whether bond dipoles cancel.
• Bonding in Coordination Compounds: VBT, Crystal Field Theory and Metal Carbonyls – Valence bond theory (VBT) says the metal mixes its empty orbitals into hybrid orbitals (like d²sp³ or sp³d²) that accept the ligand electron pairs; it predicts shape and magnetism. Crystal field theory (CFT) treats ligands as negative charges that raise some d orbitals more than others: in an octahedron the 5 d orbitals split into t₂g (lower) and e_g (higher) by Δo. Strong ligands make Δo big and electrons pair (low spin). An electron jumping across Δo by absorbing light gives colour. In metal carbonyls, CO gives a σ pair and takes back π electrons (synergic bonding).
• Buffer Solutions and Solubility Product – A buffer is a solution that keeps its pH nearly the same when a little acid or base is added. An acidic buffer is a weak acid with its salt (CH₃COOH + CH₃COONa); a basic buffer is a weak base with its salt (NH₄OH + NH₄Cl). Its pH is given by the Henderson equation: pH = pKa + log([salt]/[acid]). A sparingly soluble salt in its saturated solution sets up an equilibrium with its ions; the product of ion concentrations (each raised to its coefficient) is the solubility product Ksp. For AB, Ksp = s²; for AB₂ or A₂B, Ksp = 4s³. If the ionic product Q exceeds Ksp, a precipitate forms. A common ion lowers solubility.
• Carbohydrates: Classification, Glucose, Disaccharides and Polysaccharides – Carbohydrates are sugars and the big molecules made from sugars. Chemically they are polyhydroxy aldehydes or ketones (many –OH groups plus one C=O), or things that give these on hydrolysis. We sort them by how many sugar units they have: one (monosaccharide), two to ten (oligosaccharide) or many (polysaccharide). Glucose is an aldose with six carbons; fructose is a ketose. Two units join by a glycosidic (–O–) bond to give disaccharides like sucrose, maltose and lactose. Hundreds join to give starch, cellulose and glycogen.
• Carbon and Its Compounds: Bonding, Hydrocarbons and Naming – Carbon has 4 outer electrons, so it shares electrons (covalent bonds) instead of gaining or losing them. Because it bonds to itself (catenation) and always makes 4 bonds (tetravalency), it forms millions of compounds: chains, branches and rings, saturated or unsaturated. Compounds with the same functional group form a homologous series that differs by –CH₂–, and IUPAC names are built from the number of carbons + a suffix or prefix for the functional group.
• Carboxylic Acids: Carboxyl Group, Acidity and Reactions – A carboxylic acid has the carboxyl group, −COOH: a C=O and an O−H on the same carbon. It can give away H⁺, and the ion left behind (carboxylate, −COO⁻) is stable because its negative charge is shared equally by two oxygen atoms. That is why these acids are stronger than phenols and alcohols (but weaker than mineral acids like HCl). Groups that pull electrons (Cl, F, NO₂) make the acid stronger; groups that push electrons (alkyl) make it weaker. We make carboxylic acids by oxidising alcohols, aldehydes or alkylbenzenes, by hydrolysing nitriles, esters and amides, or from Grignard reagents and CO₂. Their key reactions: salt formation with NaHCO₃ (CO₂ fizz), esterification, acid chloride and anhydride formation, reduction to alcohols, decarboxylation and the HVZ reaction.
• Catalysts: Speeding Up Reactions Without Being Used Up – A catalyst is a substance that makes a chemical reaction faster but is not used up; its mass and chemical nature are the same at the end. It works by giving the reaction a different route with a lower activation energy. Manganese dioxide speeds up the breakdown of hydrogen peroxide into water and oxygen. Enzymes are catalysts in living things.
• Cement: How It Is Made and Why It Hardens – Cement is a grey powder made by heating limestone and clay to about 1450 °C and grinding the result (clinker) with a little gypsum. Mixed with water it sets and hardens, and with sand and stones it makes concrete.
• Ceramic Raw Material Processing – Ceramic ware starts from natural minerals: clay for shape, quartz for strength and feldspar as a flux. They are crushed, ground, sieved, cleaned of iron and then weighed by percent into a batch. Good powder means good ware.
• Ceramics (Pottery): Earthenware and Porcelain – Pottery is made from clay mixed with silica and feldspar. The clay is shaped, dried, fired and often glazed. Earthenware is fired at about 1000 °C: it is porous and orange or buff. Porcelain is fired at about 1300 °C: it is white, dense, ringing and does not soak water.
• Chemical Bonding: Ionic, Covalent and Metallic Bonds – Atoms join together (bond) to become more stable. Only their outer electrons take part. Most atoms are most stable with 8 outer electrons: the octet rule. There are three main ways to reach it. In an ionic bond, a metal gives electrons to a non-metal, making oppositely charged ions that attract. In a covalent bond, two non-metals share pairs of electrons to make molecules. In a metallic bond, metal atoms release outer electrons into a shared 'sea' that holds positive ions together. The difference in electronegativity (how strongly an atom pulls shared electrons) tells us which kind of bond forms. The type of bond explains melting points, whether a substance conducts electricity, and whether it dissolves in water.
• Chemical Cells and Fuel Cells – A chemical cell turns chemical energy into electrical energy. It has two different metals (electrodes) in an electrolyte. The more reactive metal gives away electrons, which flow through the wire to the less reactive metal. The bigger the difference in reactivity, the bigger the voltage; the electrolyte matters too. A battery is two or more cells in series, so their voltages add. Non-rechargeable cells stop when a reactant is used up; rechargeable cells can be reversed by an external current. A hydrogen fuel cell combines hydrogen and oxygen to make water and electricity, as long as fuel is supplied.
• Chemical Change and Energy – In a chemical change atoms swap partners and new substances form. Five ideas explain most of it: acids and bases (H+ meets OH- and makes water), oxidation and reduction (electrons move), heat (exothermic gives heat out, endothermic takes heat in), rate and equilibrium (how fast, and where it settles), and nuclear energy (a nucleus changes, and far more energy comes out).
• Chemical Energy: Heat, Bond Energy and Cells – A chemical reaction breaks old bonds (energy in) and makes new bonds (energy out). If more energy goes out than in, the reaction is exothermic and heat is released (ΔH negative). If more goes in, it is endothermic (ΔH positive). With bond energies, ΔH = energy to break bonds − energy released making bonds. If the electrons of a redox reaction are made to travel through a wire instead of jumping directly, the energy comes out as electricity: that is a galvanic cell, such as Zn–Cu giving 1.1 V. Joining cells in series gives a battery. Primary cells are used once, secondary cells are rechargeable, and fuel cells are fed with fuel.
• Chemical Equilibrium: Kc, Kp, Q and Gibbs Energy – In a closed container a reversible reaction goes both ways. After some time the forward and backward rates become equal, so amounts stop changing, but the reaction does not stop. This is dynamic equilibrium. At equilibrium the ratio of products to reactants (each raised to its coefficient) is a fixed number, the equilibrium constant K. Kc uses concentrations, Kp uses partial pressures, and Kp = Kc(RT)^Δn. Pure solids and liquids are left out of K. The reaction quotient Q tells the direction: Q < K goes forward, Q > K goes backward, Q = K is equilibrium. K and Gibbs energy are linked: ΔG = ΔG° + RT ln Q and ΔG° = −RT ln K.
• Chemical Properties of Carbon Compounds, Ethanol and Ethanoic Acid – Carbon compounds burn in oxygen to give CO₂, water, heat and light (combustion); strong oxidising agents turn alcohols into acids (oxidation); unsaturated compounds add hydrogen with a nickel catalyst (addition); saturated compounds swap hydrogen for chlorine in sunlight (substitution). Ethanol reacts with sodium to give hydrogen and dehydrates to ethene; ethanoic acid is a weak acid that makes esters with alcohols and CO₂ with carbonates.
• Chemical Reactions and Balancing Equations – In a chemical reaction atoms are not made or destroyed; they only change partners. So a chemical equation must have the same number of each kind of atom on both sides. We balance it by changing the numbers in front of formulas, never the formulas themselves.
• Chemical Reactions: Equations, Types, Prediction and Coupled Reactions – A chemical reaction rearranges atoms into new substances; atoms and mass are conserved, so equations must balance. The same reaction can be written as a word equation, a balanced formula equation with state symbols, a full ionic equation or a net ionic equation that leaves out spectator ions. Reactions are sorted into types: synthesis, decomposition, single displacement, double displacement (precipitation, neutralisation, gas-forming), combustion, and by the particle that moves: electrons (redox) or protons (acid–base). Outer (valence) electrons let us predict what forms: metals lose electrons, non-metals gain them, and the numbers lost and gained must match. A reaction that cannot happen alone (ΔG > 0) can be driven by coupling it to a strongly favourable one so that the total ΔG is negative.
• Chemical Technology: From Lab Flask to Factory – A reaction in a small flask behaves differently in a big reactor. When the size grows by L, volume (and heat made) grows by L³ but surface (where heat leaves) grows only by L². So large reactors overheat unless they have cooling and careful feeding. Reactivity also depends on surface area, temperature, concentration and catalysts. Engineers scale up in steps: lab, pilot plant, factory.
• Chemistry and Society – Chemistry studies how atoms join and rearrange. By choosing which atoms to join and how, chemists make new materials, medicines, fertilisers and fuels. Atoms are never lost in a reaction, so every product has a cost and a waste that we must plan for.
• Chemistry Experiments: Prepare, Separate, Test and Measure – Four skills run through every chemistry lab: prepare a substance, separate it from a mixture, test to identify it, and measure exact amounts (as in titration). The same ideas scale up to a factory model with recycling, and to testing water for pH and nitrate.
• Chemistry Frontiers: Catalysts, New Materials and Chemical Plants – Modern chemistry works with physics, biology, engineering and computing. It makes new molecules (often with catalysts), builds new materials and batteries, and scales reactions up in chemical plants where the leftovers are recycled.
• Chemistry in Building and Agriculture – Builders use chemistry to turn limestone into lime, cement, concrete and glass, and to add steel where concrete is weak. Farmers use chemistry to give plants nitrogen, phosphorus and potassium. In both jobs the right amount matters.
• Chemistry in Everyday Life: Medicines, Food, Cleaning and Soil – Chemistry is at work in your home every day. Antacids are weak bases that calm extra stomach acid. Medicines are grouped by what they do: painkillers, antibiotics, antiseptics, antihistamines, tranquilisers. Food chemicals keep food safe and tasty: preservatives, antioxidants and sweeteners. Soap cleans because one end of its molecule loves water and the other loves oil. Farmers add lime to acid soil and fertilisers to feed crops. Some metal ions are needed by the body, others are toxic, which is one reason we recycle.
• Chemistry in Health, Materials, Industry and Society – Chemistry explains the energy in food, why drug dose matters, why we pick steel, plastic, glass or ceramic, and how factories turn raw materials into fertiliser. Every useful product also has a cost, so we weigh benefit against harm and manage the risk.
• Chemistry of Fibres and Dyeing – A fibre is a bundle of long polymer chains made by joining many small monomers. Cotton is cellulose, wool and silk are protein, polyester and nylon are synthetic polymers. Each has different binding sites, so each needs a matching dye class: reactive for cotton, acid for wool, silk and nylon, disperse for polyester, basic for acrylic. Heat, pH, salt and helper chemicals decide how much dye goes in and how firmly it holds.
• Chemistry Review: From the Periodic Law to Electrolysis – This review ties five big ideas together. The periodic law: properties repeat with atomic number. Structure: atoms join by giving or sharing electrons. Rates and energy: faster, more energetic particles react faster, and reactions give out or take in heat. Solutions: a substance spreads out in a solvent. Electrolysis: electric current forces a chemical change and moves ions to the electrodes.
• Chemistry Skills: How Chemists Think – Chemists look at one change on three levels: macro (what we see), meso (groups of particles) and micro (atoms and ions). They draw a process as a system with inputs, the process and outputs, estimate numbers before calculating, link ideas with physics and biology, and judge new solutions for sustainability. Chemical knowledge grows by testing and improving models.
• Chromatography: Separating a Mixture by How Much It Sticks – Chromatography separates the parts of a mixture. It always has two phases. The stationary phase stays still (paper, a silica layer, a packed column). The mobile phase moves over it (a liquid solvent or a gas). Each substance in the mixture sticks to the stationary phase by a different amount. The one that sticks less moves faster and goes further. In paper chromatography and thin-layer chromatography (TLC) we measure the Rf value: distance moved by the spot divided by distance moved by the solvent front. In column chromatography the substance that sticks least comes out of the bottom first. In gas chromatography (GC) each substance leaves the column at its own retention time, and the area of its peak tells how much is there. We use chromatography to test purity, identify substances and separate them.
• Colligative Properties: Counting Particles, Not Their Type – Colligative properties depend only on HOW MANY solute particles are in the solution, not on what they are. There are four: relative lowering of vapour pressure (Δp/p° = x₂), elevation of boiling point (ΔTb = Kb·m), depression of freezing point (ΔTf = Kf·m) and osmotic pressure (π = CRT). We use them to find molar mass. Salts that split into ions (NaCl) or molecules that pair up (acetic acid in benzene) give abnormal molar masses, fixed with the van 't Hoff factor i.
• Colloids and Dispersed Systems – A dispersed system is one substance spread as small bits (the dispersed phase) through another (the dispersion medium). By particle size it is a true solution (below about 1 nm), a colloid (about 1–100 nm) or a suspension (above about 100 nm). Colloid particles are too small to see or to settle, but big enough to scatter light, so a beam passing through shows its path: the Tyndall effect. Colloids are grouped by the states of the two phases (fog, smoke, milk, foam, gel). Colloid particles often carry the same charge and repel each other; adding an electrolyte cancels the charge, so they clump and settle: coagulation.
• Colour and Dyes – A substance looks coloured because its molecules absorb some colours of white light and the rest reaches the eye. The absorbed light lifts an electron across the HOMO-LUMO gap. A longer chain of alternating double bonds makes the gap smaller, so the dye absorbs longer wavelengths. You see the complementary colour. Dyes dissolve; pigments stay as solid grains.
• Combustion – Combustion (burning) is a chemical change in which a fuel reacts fast with oxygen and gives out heat and light. It needs three things: fuel, oxygen and heat (to reach the ignition temperature). Remove any one and the fire goes out. With plenty of air, fuels that contain carbon and hydrogen burn completely to carbon dioxide and water (blue flame). With too little air, burning is incomplete and gives soot and poisonous carbon monoxide (yellow, smoky flame). Carbon dioxide turns limewater milky. The energy released by burning 1 g of a fuel is its calorific value.
• Composition of Air – Air is a mixture of gases. By volume, dry air is about 78% nitrogen, 21% oxygen, 0.93% argon and 0.04% carbon dioxide, with tiny amounts of other noble gases. Water vapour changes from 0 to about 4%. We can show air is a mixture and measure its oxygen by letting iron rust (or phosphorus burn) in a closed tube: water rises about one fifth. Pollutants such as smoke, sulfur dioxide, nitrogen oxides and carbon monoxide make air harmful, and we can reduce them.
• Conductance of Electrolytic Solutions – In a solution, ions carry the current. Resistance R depends on the cell shape through the cell constant G* = l/A, so conductivity κ = G*/R. Molar conductivity Λm = κ × 1000/c tells how well one mole of the solute conducts. On dilution κ falls but Λm rises: a little for strong electrolytes (Λm = Λ°m − A√c) and a lot for weak ones. Kohlrausch's law says Λ°m = sum of the ion values, which gives Λ°m of weak electrolytes and their degree of dissociation.
• Cooking Skills and Techniques – Good cooking is safe, skilful and smart. Keep food out of the danger zone (5–63 °C), cut pieces the same size, know how heat reaches food (conduction, convection, radiation), choose a moist or dry method, use seasonal ingredients, and scale a recipe and its cost for any number of people.
• Coordination Compounds: Werner's Theory, Terms and IUPAC Names – A coordination compound has a central metal atom or ion joined to a fixed number of ions or molecules called ligands. Each ligand gives one pair of electrons to the metal. The metal and its ligands sit together inside square brackets, the coordination sphere. The number of donor atoms bonded to the metal is the coordination number. Werner said a metal shows two kinds of valence: primary (ionisable, outside the bracket) and secondary (fixed, inside the bracket).
• Corrosion, Its Prevention and Alloys – Corrosion is the slow eating away of a metal by air, moisture or chemicals around it. Iron rusts only when both air (oxygen) and water are present, forming hydrated iron oxide. It can be prevented by painting, oiling, galvanising, chrome plating, anodising or making alloys. An alloy is a uniform mixture of a metal with other metals or a non-metal; alloys are usually harder, stronger and more corrosion-resistant than the pure metal.
• Crystalline Ceramic Materials – Crystalline ceramics have ions or atoms in repeating patterns. Silica (SiO₂) is a network of SiO₄ tetrahedra sharing corners. Alumina (Al₂O₃) has aluminium in the gaps between layers of oxygen and is very hard. Clay minerals like kaolinite are stacks of silica and alumina sheets, with water between them that lets wet clay slide and be shaped. Oxide ceramics (MgO, ZrO₂) and non-oxides (SiC, Si₃N₄) serve high-heat and high-wear jobs.
• Culinary Techniques: Stocks, Sauces, Doughs and Creams – Good cooking starts before the stove: choose the right product (quality, label, level of processing), buy in season, follow a recipe card and portion sheet, then do preliminary preparations (washing, cutting, weighing) and base preparations (stocks, sauces, doughs, creams). At the end, judge the dish with your senses: look, smell, taste and texture.
• Design a Simple Oxygen Supplier – A simple oxygen supplier uses 2H₂O₂ → 2H₂O + O₂ with MnO₂ as catalyst. It needs no heating. The gas passes through a water bottle, which cools it and lets you count bubbles, then leaves by an outlet tube. A glowing splint that relights proves it is oxygen. This is a science model, not medical equipment.
• Diazonium Salts: Preparation, Properties and Use in Making Aromatic Compounds – A diazonium salt is Ar–N₂⁺X⁻, e.g. benzenediazonium chloride C₆H₅N₂⁺Cl⁻. It is made by diazotisation: aniline + NaNO₂ + HCl at 273–278 K. It is stable only in the cold and is used at once. On warming, N₂ gas (a great leaving group) escapes. Two kinds of reactions: (1) N₂ is replaced by Cl, Br, CN (Sandmeyer/Gattermann), I (KI), F (HBF₄), H (H₃PO₂) or OH (warm water); (2) N₂ is kept and the ion couples with phenol or aniline to give coloured azo dyes. This lets us put groups on a benzene ring that cannot be added directly.
• Dienes: Two Double Bonds, One Chain – Dienes (alkadienes, general formula CnH2n-2) have two C=C bonds. Bonds can be cumulated, conjugated or isolated. In conjugated dienes such as buta-1,3-diene the pi electrons spread over four carbons, so reagents add at 1,2 or 1,4 positions. Dienes join into polymers, the rubbers.
• Drug Design: From Molecule to Medicine – A medicine is a molecule that fits a target in the body, like a key in a lock. Many drugs started in plants (natural), then chemists changed them in the lab to fit better and cause fewer side effects (synthetic). Nanomedicines are tiny capsules that carry a drug to the right place.
• Dyes, Paints, Flavours and Perfumes – Dyes are coloured substances that stick to fibres or dissolve in water. Pigments are coloured powders used in paints. Paint = pigment + binder + solvent. Food colours must be safe to eat. Aromas, essences and perfumes are light, easily evaporating organic compounds that our nose senses.
• Electrolysis, Batteries, Fuel Cells and Corrosion – In an electrolytic cell an outside source of electricity forces a non-spontaneous reaction: cations are reduced at the cathode (−) and anions oxidised at the anode (+). Faraday's laws link the mass changed to the charge: m = (M/nF) × I × t. Which product forms depends on electrode potentials and overpotential. Batteries are galvanic cells: primary (dry cell, mercury cell) cannot be recharged; secondary (lead storage, Ni–Cd) can. Fuel cells burn H₂ with O₂ to give electricity directly. Corrosion (rusting) is an unwanted galvanic cell on the metal surface.
• Electrolysis: Splitting Compounds with Electricity – Electrolysis uses a direct current to break down an ionic compound that is melted or dissolved (the electrolyte). Positive ions move to the cathode (−) and gain electrons (reduction). Negative ions move to the anode (+) and lose electrons (oxidation). Molten compounds give the metal and the non-metal. In solutions, water also gives H⁺ and OH⁻: hydrogen forms at the cathode unless the metal is less reactive than hydrogen, and oxygen forms at the anode unless a halide is present. Reactive metals like aluminium are extracted by electrolysis.
• Electrolytes and Electrolytic Dissociation – An electrolyte is a substance that splits into free-moving ions when it dissolves in water (or melts), so its solution carries electric current. Non-electrolytes such as sugar stay as neutral molecules and do not conduct. Strong electrolytes split almost fully; weak ones split only a little. The degree of dissociation α = ions-forming units ÷ total units.
• Electronic Configuration and the s, p, d, f Blocks – The periodic table is really a map of electron filling. The period number is the highest shell (n) being used. The block is the subshell that receives the last electron: s-block (groups 1–2, ns¹⁻²), p-block (groups 13–18, ns² np¹⁻⁶), d-block (groups 3–12, (n−1)d¹⁻¹⁰ ns⁰⁻²) and f-block (lanthanoids and actinoids, (n−2)f¹⁻¹⁴ (n−1)d⁰⁻¹ ns²). Group: s-block = number of ns electrons; p-block = 10 + ns + np electrons; d-block = ns + (n−1)d electrons.
• Elements and Their Symbols – An element is a pure substance made of only one kind of atom. Each element has a short symbol: one capital letter, or a capital plus a small letter (O, Cl). Some symbols come from old Latin names (Na, Fe, Cu). Elements are sorted into metals, non-metals and noble gases, and arranged in the periodic table by atomic number.
• Energy and Industry: Sources, CO2 and Life Cycles – Factories need energy. It comes from fossil fuels, nuclear power, hydro, wind and solar. To judge a source we compare CO2 per kWh, reliability, cost and other effects. A life cycle looks at a product from raw material to making, use, disposal and recycling, because energy and waste appear at every stage.
• Enthalpies of Different Types of Reactions – Each kind of change has its own named enthalpy, always for 1 mol. Phase changes: fusion (melting), vaporisation (boiling) and sublimation (ΔsubH = ΔfusH + ΔvapH); all are endothermic. Combustion: 1 mol burns completely in oxygen; always exothermic. Atomisation: 1 mol of a substance breaks fully into gaseous atoms. Bond dissociation enthalpy: 1 mol of a given bond is broken in the gas phase; for polyatomic molecules we use mean bond enthalpy. ΔrH ≈ Σ(bonds broken) − Σ(bonds formed). Lattice enthalpy: 1 mol of an ionic solid is split into gaseous ions (found by a Born–Haber cycle). Enthalpy of solution = lattice enthalpy + hydration enthalpy.
• Enthalpy, Calorimetry and Hess's Law – Enthalpy H = U + pV. At constant pressure the heat taken in or given out equals ΔH; at constant volume it equals ΔU. We measure ΔU in a sealed bomb calorimeter and ΔH in an open cup calorimeter, using q = C ΔT or q = m c ΔT. For gases, ΔH = ΔU + Δn_g RT. The reaction enthalpy ΔrH is negative for exothermic and positive for endothermic reactions, and ΔrH° = ΣΔfH°(products) − ΣΔfH°(reactants). Hess's law: the total enthalpy change is the same whether a reaction happens in one step or many, so thermochemical equations can be added like algebra.
• Environmental Chemistry – Environmental chemistry studies the chemicals in air, water and soil, where they come from and what they do. Dry air is about 78% nitrogen, 21% oxygen, 0.9% argon and 0.04% carbon dioxide. Burning fuels adds pollutants such as SO₂, NOₓ, CO and particulates. SO₂ and NO₂ form acids in rain (pH below 5.6). We measure pollution with quantitative tests such as titration, compare the results with standards, and use laws and cleaner technology to reduce harm.
• Esters: Sweet-Smelling Compounds From Acids and Alcohols – An ester forms when a carboxylic acid reacts with an alcohol: acid + alcohol ⇌ ester + water. The acid loses –OH, the alcohol loses H, and the ester bond –COO– joins the two parts. Concentrated sulfuric acid is the catalyst and also removes water, so more ester forms. Esters smell fruity and are used in flavours, perfumes and solvents. Water splits esters back (hydrolysis); with an alkali this is saponification. Fats and oils are esters of glycerol with three fatty acids; boiling them with NaOH gives soap and glycerol.
• Ethers – An ether has one oxygen joining two carbon groups: R–O–R′. Name it as alkoxyalkane (the smaller group becomes the alkoxy part). Make ethers by dehydrating primary alcohols with conc. H₂SO₄ at 413 K (symmetrical ethers) or by Williamson synthesis: sodium alkoxide + primary alkyl halide (SN2). Ethers cannot H-bond with each other, so their boiling points are low, close to alkanes, but they H-bond with water and small ones dissolve slightly. HI breaks the C–O bond: the smaller group becomes the alkyl iodide (SN2), a tertiary group does so by SN1, and anisole gives phenol + CH₃I. The –OCH₃ group in anisole directs ring substitution to ortho and para.
• Exothermic and Endothermic Reactions – Energy is conserved in every reaction; it only moves between the chemicals and the surroundings. An exothermic reaction transfers energy to the surroundings, so the temperature rises (burning, neutralisation, respiration, hand warmers). An endothermic reaction takes energy in, so the temperature falls (thermal decomposition, citric acid + sodium hydrogencarbonate, photosynthesis, cold packs). A reaction profile shows reactants, products and the activation energy (the minimum energy needed to react). Breaking bonds takes energy in; making bonds gives energy out. ΔH = energy to break bonds − energy released making bonds; negative means exothermic.
• Fertilisers, Soil Acidity and Plant Nutrients – Plants take nitrogen (N), phosphorus (P) and potassium (K) from soil. Fertilisers add them back: N for leaves, P for roots and flowers, K for a strong stem and fruit. In acidic soil nutrients are locked, so lime raises the pH. The scarcest nutrient limits growth, and too much fertiliser burns plants and pollutes water. Inorganic classes are linked: metal to basic oxide to base to salt, and non-metal to acidic oxide to acid to salt.
• Fibre Manufacture and Dyeing Technology – People first spun natural fibres by hand and coloured them with plants. Machines and, later, man-made dyes (1856) and man-made fibres (1930s) built the modern textile industry. A man-made fibre is made by pushing melted or dissolved polymer through the tiny holes of a spinneret. Dyes colour the cloth in a dye bath, and three primary dyes (cyan, magenta, yellow) can be mixed to make almost any colour.
• First Law of Thermodynamics: System, Heat, Work and ΔU – Thermodynamics tracks energy. The part we study is the system; the rest is the surroundings; the wall between them is the boundary. A system can be open, closed or isolated. Internal energy U is the total energy stored in the system. It changes only in two ways: by heat q or by work w. The first law says ΔU = q + w (IUPAC signs: q and w are positive when energy goes INTO the system). U is a state function; q and w are path functions. For expansion against a constant outside pressure, w = −p_ext ΔV; for a reversible isothermal expansion of an ideal gas, w = −2.303 nRT log(V₂/V₁).
• Food Component Analysis – Food analysis finds how much of each component a food has. Dry it to find water, burn it to find ash (minerals), extract fat with a solvent, measure nitrogen and multiply by 6.25 for protein, and take carbohydrate as what is left. Vitamins and some minerals need special tests because they are present in tiny amounts.
• Food Control Experiments – Food labs follow safe habits, measure carefully and write every reading. A chemical test such as titration finds how much acid is in food. A microbe test dilutes the sample and counts colonies on plates to give CFU per gram.
• Food Science: The Chemistry of What We Eat – Food science studies what food is made of and how it changes from farm to plate. Food contains carbohydrates, proteins, fats, water, vitamins and minerals. Cooking changes them: heat unfolds proteins (denaturation), starch swells in water (gelatinisation), and dry heat browns food through the Maillard reaction and caramelisation. Microbes ferment food into bread, curd and pickles, but can also spoil it. Processing, preservation and safe temperatures keep food safe and nutritious for a growing world population.
• Forming and Drying Ceramics – Ceramic powder or clay is first formed into a shape: by hand or wheel (plastic forming), by pressing powder in a die, or by pouring slip into a plaster mold. The wet piece is then dried slowly. Water leaves and the piece shrinks a little. Fast drying cracks it.
• Fuels and Lubricants – Ship engines burn fuel oil: thick, cheap heavy fuel oil (HFO) that must be heated and cleaned, or thinner distillates like marine diesel (MDO) and gas oil (MGO). Important fuel properties are heat value, viscosity, flash point and sulphur. Lubricating oil puts a thin film between moving metal parts. The film cuts friction and wear, carries heat away, cleans, seals and protects from rust. Oil gets thinner when hot and thicker when cold, so temperature matters for both fuel and lube oil.
• Fuels, Furnaces, Firing and Melting – Fuel burns with oxygen to give heat. A furnace (kiln) holds that heat so ware can be fired at 1000 to 1400 °C, where grains sinter and the piece becomes hard. Glass batch needs about 1500 °C to melt. The fuel limits the temperature you can reach.
• Functional Ceramics – Functional ceramics are carefully made ceramics that do a special job. Mechanical ones are very hard and wear-proof (cutting tools, bearings). Electrical ones insulate, store charge or make voltage when squeezed (piezo). Optical ones let light pass or glow (lamp tubes, lasers). Pure powders and tiny, even grains are the secret.
• Functional Groups: Halogen, Hydroxyl and Amine Compounds – A functional group is an atom or small group of atoms that replaces a hydrogen on a carbon chain and decides how the compound behaves. Halogen compounds (R–X) have a polar C–X bond. Hydroxyl compounds (R–OH) form hydrogen bonds, so they boil high; phenols are weakly acidic. Amines (R–NH2) have a lone pair on nitrogen, so they are bases.
• Galvanic Cells and the Nernst Equation – A galvanic cell turns the energy of a redox reaction into electricity. Oxidation happens at the anode (−) and reduction at the cathode (+). Each electrode has a potential measured against the standard hydrogen electrode (0 V). E°cell = E°cathode − E°anode. The Nernst equation, E = E° − (0.059/n) log Q at 298 K, gives the cell voltage at any concentration, and ΔG = −nFE links voltage to energy.
• Gas Laws: How Pressure, Volume and Temperature Are Linked – Gas pressure comes from particles hitting the walls. For a fixed amount of gas: Boyle's law P₁V₁ = P₂V₂ (constant T); Charles's law V₁/T₁ = V₂/T₂ (constant P); pressure (Gay-Lussac's) law P₁/T₁ = P₂/T₂ (constant V). Temperature must be in kelvin. Combined: P₁V₁/T₁ = P₂V₂/T₂. Dalton's law: in a mixture, total pressure = sum of partial pressures, and each partial pressure = mole fraction × total pressure.
• General Chemistry: Atoms, Bonds, Names and Carbon – Everything is made of atoms. Electrons in the outer shell decide how atoms join: by giving and taking (ionic bond) or by sharing (covalent bond). The periodic table sorts atoms by atomic number, simple rules give compounds their names, and carbon, with 4 bonds, builds the huge family of organic molecules.
• Genetic Links Between Substances – Substances are related like a family tree. A metal burns to a basic oxide, the oxide meets water to give a base. A non-metal burns to an acidic oxide, which meets water to give an acid. A base and an acid react to give a salt and water. Organic substances join this web too: they burn to CO2 and water, and plants turn CO2 back into organic matter.
• Glass and Enamel – Common glass is made by melting sand (silica), soda ash and limestone at about 1500 °C, shaping it while soft, then cooling it slowly (annealing). Glass is a solid with no crystal order. Enamel is a thin glass coat fused onto metal at about 800 °C to stop rust and give a shiny, easy-to-clean surface.
• Green Chemistry – Green chemistry means designing chemical products and processes that make less waste, use safer substances, save energy and use raw materials that can be renewed. Key tools are high atom economy, catalysts, safer solvents such as water, renewable feedstocks, and products that break down safely. A life cycle assessment checks the impact of a product from raw material to disposal.
• Group 2, the Alkaline Earth Metals: Trends, Reactions and Uses – Group 2 metals (Be, Mg, Ca, Sr, Ba) have two outer s electrons and form M²⁺ ions. Down the group atomic radius increases, first ionisation energy decreases and melting point generally decreases (Mg is an exception). Reactivity with water increases: M + 2H₂O → M(OH)₂ + H₂, while Mg reacts with steam to give MgO + H₂. Hydroxides become more soluble down the group; sulfates become less soluble, so BaSO₄ is insoluble. Uses: Mg(OH)₂ antacid, Ca(OH)₂ for acidic soil, CaO/CaCO₃ to remove SO₂, BaSO₄ for X-ray meals, BaCl₂ to test for sulfate, Mg to extract titanium.
• Hairstyling and Aesthetics: The Science, Art and Business of the Salon – Hairstyling and aesthetics is the care and styling of hair, skin and nails. You need science (hair structure, pH, chemicals and allergies), art (design principles, colour theory, trends and make-up) and people skills (consultation, record keeping, teamwork, advertising and safe, accurate measuring).
• Haloalkanes – A haloalkane is an alkane in which a hydrogen is replaced by a halogen (R–X). The C–X bond is polar: carbon is slightly positive, so nucleophiles attack it. Haloalkanes react by SN2 (one step, backside attack, inversion; best for 1°), SN1 (two steps, flat carbocation, racemic product; best for 3°) or elimination (strong base in alcohol gives an alkene).
• Haloarenes – A haloarene has a halogen joined straight to a benzene ring (Ar–X), like chlorobenzene C₆H₅Cl. Resonance gives the C–X bond part double-bond character, so haloarenes resist nucleophilic substitution unless –NO₂ groups sit at ortho/para. With electrophiles, the halogen slows the ring slightly but sends new groups to ortho and para positions.
• Halogens: The Group 17 Elements – The halogens are fluorine, chlorine, bromine, iodine and astatine, in group 17 of the periodic table. Each atom has 7 outer electrons, so it gains 1 electron to form a halide ion (X⁻). Down the group, atoms get bigger, colours get darker, melting and boiling points rise, and reactivity falls. A more reactive halogen displaces a less reactive one from its salt.
• Heterocyclic Compounds – A heterocyclic compound has a ring that contains at least one atom other than carbon, such as nitrogen, oxygen or sulfur. Pyridine is like benzene with one C–H replaced by N; furan, pyrrole and thiophene are 5-membered rings. Many vitamins, drugs, dyes, DNA bases and blood and plant pigments are heterocycles.
• High-Temperature Reactions – Heat makes atoms vibrate harder, so some jump to new places. This movement (diffusion) lets two solids react at their contact, with a product layer that grows slowly. At the melting point a solid becomes liquid; on cooling, crystals start (nucleation) and grow. In a furnace, oxygen can be added (oxidation) or taken away by carbon or hydrogen (reduction), and this changes the product.
• History of Materials Development – People have named whole eras after their main material: stone, bronze, iron. Each step needed new skills, mostly better control of heat and chemistry. Steel arrived with cheap mass production in the 1800s, plastics and silicon in the 1900s, and now composites and nano-materials are designed atom by atom.
• Household Chemistry: The Chemicals in Your Home – Every home is full of chemicals: vinegar, soap, baking soda, bleach and drain cleaner. Each has a pH that tells if it is an acid or a base. Some substances mix (salt in water) and some do not (oil and water), because like dissolves like. Strong products are diluted with water using C1V1 = C2V2. Some reactions are safe (baking soda + vinegar), but some mixes make toxic gas (bleach + ammonia or acid). Leftover chemicals must go to a hazardous waste drop-off, not down the drain.
• Household Fuels: Change Over Time and Safe Use – Homes moved from dung cakes and wood to coal, kerosene, LPG, piped natural gas and electric or solar cooking. Cleaner fuels make less smoke and put more heat into the pot. Every fuel needs safe use: good ventilation, leak checks for gas, and never burning fuel in a closed room because poisonous carbon monoxide builds up.
• How Structure Decides the Properties of Substances – The way particles are held together decides how a substance behaves. Ionic and giant covalent substances have strong bonds all through, so they melt very high. Small molecules have weak forces between molecules, so they melt and boil low. Metals have free electrons, so they conduct.
• Hydrocarbons: Where They Come From, What They Do, and Rings – Hydrocarbons are compounds of only carbon and hydrogen. We get them from crude oil and natural gas by heating and separating (fractional distillation), and long chains are broken into useful short ones by cracking. They give us fuels and plastics, but burning them makes CO2. If a chain closes into a ring we get a cycloalkane, CnH2n, which has 2 H fewer than the chain.
• Hydrogen Bonding – A hydrogen bond is a weak attraction between an H atom that is joined to a small, very electronegative atom (F, O or N) and a lone pair on another such atom. It is shown with a dotted line: X–H···Y. It is much weaker than a covalent bond (about 10–40 kJ mol⁻¹) but strong enough to raise boiling points, make ice float and hold DNA together. Intermolecular H-bonds join different molecules; intramolecular H-bonds form inside one molecule.
• Hydrogen: The Lightest Element – Hydrogen (H) is element 1: one proton and one electron. It exists as H₂ molecules, a colourless, odourless gas that is the least dense of all gases and hardly dissolves in water. In the lab it is made from a reactive metal and a dilute acid (Zn + 2HCl → ZnCl₂ + H₂) and collected over water. It burns with a squeaky pop to form only water. It reacts with non-metals to give hydrides such as HCl, NH₃ and H₂S, and it reduces some metal oxides. It has three isotopes: protium, deuterium and tritium. Uses: ammonia and fertilisers, margarine, rocket fuel and clean fuel cells.
• Indian Cuisine: Regions, Snacks, Gravies, Sweets and Serving – Indian cooking changes from region to region because of climate, crops and culture: wheat and dairy in the north, rice and coconut in the south, rice and fish in the east, millets and peanuts in the west. Common building blocks are spices, tempering (tadka), a few basic gravies (onion-tomato, white, yellow, green), fried and steamed snacks, and milk-based sweets made from khoya, chhena and sugar syrup. A meal is often served as a thali that balances food groups and the six tastes.
• Industrial Chemistry: How a Chemical Plant Works – A chemical plant turns cheap raw materials into useful products. It prepares the raw materials, reacts them in a reactor under chosen conditions, separates the product and sends unused material back (recycling). Engineers balance speed, yield, cost, safety and the environment.
• Industrial Safety: Risk and Safety in Chemical Production – A hazard is something that can cause harm. Risk is how likely the harm is and how bad it would be. Chemical plants name their risks (fire, poison, pressure, heat), mark them with hazard symbols, and use many layers of protection such as good design, alarms, relief valves and training. Risk is judged as chance times effect and is kept low by adding layers.
• Integrated Rate Equations: Zero and First Order – Integrated rate equations link concentration with time. Zero order: [R] = [R]₀ − kt (straight line), t½ = [R]₀/2k. First order: k = (2.303/t) log([R]₀/[R]), ln[R] falls in a straight line with slope −k, and t½ = 0.693/k, which does not depend on the starting amount.
• Interdisciplinary Project: A Micro Air-Quality Station – A micro air-quality station is a small box that measures the air. A dust sensor counts tiny particles (PM2.5) by watching how they scatter a light beam. Gas and temperature sensors measure other things. A small computer (microcontroller) reads each sensor, turns the signal into numbers, shows them on a screen with a colour band and sounds an alarm when a threshold is crossed. Good stations are placed carefully, calibrated and averaged over time.
• Intermolecular Forces – Intermolecular forces are weak attractions BETWEEN molecules. From weakest to strongest (for similar-size molecules): London dispersion < dipole–dipole < hydrogen bond. Stronger forces mean higher melting and boiling points, higher viscosity and surface tension, and they decide what dissolves in what.
• Ionic Compounds: How Metals and Non-metals React – A metal atom gives its outer electrons to a non-metal atom so that both get a full outer shell (octet). The metal becomes a positive ion (cation), the non-metal a negative ion (anion), and the strong pull between them is an ionic (electrovalent) bond. Ionic compounds are hard crystalline solids with high melting points, dissolve in water, and conduct electricity only when molten or dissolved.
• Ionic Equations and Ion-Exchange Reactions – Salts, strong acids and alkalis split into ions in water: they are electrolytes. When two such solutions mix, a reaction happens only if some ions are removed from the solution as a precipitate, a gas or water. A full ionic equation shows every ion; crossing out the ions that do not change (spectator ions) gives the net ionic equation, for example Ag⁺ + Cl⁻ → AgCl↓. Charges and atoms must balance on both sides.
• Ionic Equilibrium: Acids, Bases, Ionisation and pH – Acids and bases can be defined in three ways: Arrhenius (give H⁺ or OH⁻ in water), Brønsted–Lowry (proton donor or acceptor) and Lewis (electron-pair acceptor or donor). Strong acids and bases ionise fully; weak ones ionise only a little and set up an equilibrium with constant Ka or Kb. Water itself ionises: Kw = [H⁺][OH⁻] = 1.0 × 10⁻¹⁴ at 25 °C. pH = −log[H⁺] and pH + pOH = 14. For a weak electrolyte, α ≈ √(K/C), so dilution increases α (Ostwald's dilution law). For a conjugate pair, Ka × Kb = Kw. Salts of weak acids or weak bases react with water (hydrolysis) and give non-neutral solutions.
• Iron and Its Compounds – Iron (Fe, atomic number 26) is a grey, magnetic transition metal. Its atom can lose 2 or 3 electrons, so it makes two families of compounds: iron(II) with Fe²⁺ and iron(III) with Fe³⁺. Iron reacts with dilute acids to give Fe²⁺ and hydrogen, with chlorine to give FeCl₃, with steam to give Fe₃O₄, and with damp air to give rust. The three oxides are FeO (black, Fe²⁺), Fe₂O₃ (red-brown, Fe³⁺) and Fe₃O₄ (black, magnetic, both ions). The two hydroxides, Fe(OH)₂ (white-green) and Fe(OH)₃ (red-brown), do not dissolve in water; Fe(OH)₂ turns into Fe(OH)₃ in air. An oxidising agent changes Fe²⁺ into Fe³⁺ and a reducing agent changes Fe³⁺ back into Fe²⁺.
• Iron and Steel Smelting – Iron ore is reduced to molten pig iron in a blast furnace by carbon monoxide from coke. Pig iron holds about 4% carbon, so it is brittle. Blowing oxygen in a converter burns most carbon away and gives steel. The steel is then cast into bars or slabs in a continuous caster.
• Isomerism in Coordination Compounds – Isomers have the same formula but a different arrangement of atoms. In structural isomers the bonds themselves are different: ionisation, linkage, coordination and solvate isomers. In stereoisomers the bonds are the same but the positions in space differ: geometrical (cis-trans, fac-mer) and optical (mirror images that cannot overlap).
• IUPAC Nomenclature and Isomerism of Organic Compounds – Millions of organic compounds exist, so every one needs a clear, unique name. The IUPAC system builds a name from three parts: the word root (number of carbons in the main chain), the suffix (type of bond and main functional group) and the prefixes (side groups). You pick the longest chain that holds the main group, number it so the main group and branches get the lowest numbers, and write side groups in alphabetical order. Isomers are compounds with the same molecular formula but different arrangement. Structural isomers differ in how atoms are joined (chain, position, functional group, metamerism). Stereoisomers are joined the same way but differ in 3D arrangement (geometrical cis/trans and optical isomers).
• Kossel-Lewis Approach and the Ionic Bond – Atoms join so that each gets a stable outer shell of 8 electrons (an octet), like a noble gas. Kossel said atoms can give or take electrons to make ions (ionic bond). Lewis said atoms can also share pairs of electrons (covalent bond). Lewis structures show these electrons as dots and lines. Formal charge (V − L − B/2) helps pick the best Lewis structure. Ions pack into a crystal, and the energy released is linked to the lattice enthalpy.
• Lab Activity: Preparing and Testing Oxygen – In the lab, oxygen is made by breaking down hydrogen peroxide with manganese dioxide (a catalyst) at room temperature, or by heating potassium permanganate. The gas goes through a delivery tube and is collected in a gas jar, either over water or by upward displacement of air. Oxygen is tested with a glowing splint: it bursts into flame again. Key safety rules: check the set-up is airtight, wait for steady bubbles, and take the tube out of the water before stopping the heat.
• Lab Safety and Safe Experiments – Dress for safety (goggles, coat, closed shoes, hair tied). Read labels and hazard pictograms. Heat gently with the mouth of the tube pointing away. Add acid to water, waft smells, never taste. Report every accident at once. Plan experiments step by step, assess the risks first and dispose of waste correctly.
• Lab Techniques: Make a Solution, Filter, Mount a Slide, Read Tissues – To make a solution, weigh the solute, dissolve it, then add solvent up to the mark of a flask; concentration = mass of solute ÷ volume of solution (for example 5 g in 100 mL = 50 g/L). Filtering separates an insoluble solid (residue) from a liquid (filtrate) using a funnel and filter paper. A microscope slide is made by putting a thin sample on glass under a coverslip lowered slowly to avoid air bubbles. Tissues are groups of similar cells with one job: epithelial, muscle, nerve and connective in animals. Research methods (observation, experiment, survey, model) answer different kinds of questions.
• Lab: Make a Simple Cell and Test Rate Factors – Part A: put a zinc plate and a copper plate in dilute acid, join them through a voltmeter and read about 1.0 V. Zinc gives electrons (anode), hydrogen forms at copper (cathode). A different metal pair gives a different voltage, and two identical plates give 0 V. Part B: plot gas collected against time. The slope of the line is the rate. A higher concentration, a higher temperature (roughly doubling per 10 °C), smaller pieces and a catalyst all raise the rate, and a fair test changes only one of them at a time. The final amount of product does not change, only how fast it forms.
• Lab: Prepare a Solution of Given Molarity – To prepare a solution of a known molarity (mol/L) you first calculate the mass of solute from mass = molarity × volume in litres × molar mass. You weigh it, dissolve it in a little water in a beaker, pour it through a funnel into a volumetric flask, rinse the beaker three times into the flask, and then add water until the bottom of the meniscus touches the ring mark, with your eye level with the mark. A stopper and 10 inversions mix the solution. Errors come from reading the mark wrongly, overshooting the mark, losing solute in transfer, using a hot solution or using a wet or dirty flask.
• Lab: Preparing a Salt (NaCl) Solution of a Given Mass Percent – To make a solution with a chosen mass percent, first calculate: salt = total mass x percent / 100, and water = total mass - salt. Weigh the salt on a balance, measure the water in a measuring cylinder (1 g of water is 1 mL), pour both into a beaker and stir until every grain has dissolved.
• Lab: Properties of Iron Compounds (Testing Fe²⁺ and Fe³⁺) – In this practical you identify Fe²⁺ and Fe³⁺ in solution. Sodium hydroxide gives a white-green solid with Fe²⁺ (turning brown in air) and a red-brown solid with Fe³⁺. Potassium thiocyanate (KSCN) turns Fe³⁺ blood red but leaves Fe²⁺ unchanged. Hydrogen peroxide oxidises Fe²⁺ to Fe³⁺ (then KSCN turns red), and iron powder reduces Fe³⁺ to Fe²⁺ (then KSCN gives no red). The lab shows that iron(II) is reducing, iron(III) is oxidising, and each can change into the other.
• Lab: Purify Crude Salt and Change Sulfur Compounds – Crude salt holds Ca²⁺, Mg²⁺ and SO₄²⁻ impurities. We remove them by precipitation: BaCl₂ takes out sulfate as BaSO₄, NaOH takes out magnesium as Mg(OH)₂, and Na₂CO₃ takes out calcium (and the extra barium) as carbonates. Then we filter, neutralise the extra OH⁻ and CO₃²⁻ with dilute HCl, and evaporate to get pure NaCl crystals. In part two we watch sulfur move between oxidation numbers −2 (H₂S), 0 (S), +4 (SO₂) and +6 (H₂SO₄): going up is oxidation, going down is reduction.
• Lab: Trends in a Period and a Group of Elements – In a lab we can see periodic trends with our own eyes. Across period 3 from sodium to chlorine, metal activity falls (sodium fizzes in water, magnesium barely, aluminium not at all), non-metal activity rises (chlorine is the strongest), and oxides change from strong base to strong acid. Down group 1 the metals get more active (Li, Na, K). Down group 17 the halogens get less active, so chlorine pushes bromine and iodine out of their salts, and bromine pushes out iodine.
• Lanthanoids and Actinoids (f-Block) – The f-block has two rows kept below the periodic table: the lanthanoids (Ce to Lu, filling 4f) and the actinoids (Th to Lr, filling 5f). Lanthanoids mostly show +3; actinoids show many more states (up to +7) and are all radioactive. Poor shielding by f-electrons makes the size shrink steadily across the row: the lanthanoid contraction, which makes 4d and 5d elements (like Zr and Hf) almost the same size.
• Laws of Chemical Combination and Dalton's Atomic Theory – Five laws describe how elements combine. Mass is conserved in a reaction. A compound always has the same elements in the same mass ratio. When two elements make more than one compound, the masses of one that join a fixed mass of the other are in a small whole-number ratio. Reacting gases have volumes in simple whole-number ratios, and equal volumes of gases hold equal numbers of molecules. Dalton's atomic theory explains all of these with atoms.
• Le Chatelier's Principle: Factors Affecting Equilibrium – Le Chatelier's principle says: if you disturb a system at equilibrium, it shifts in the direction that reduces the disturbance. Add a reactant or remove a product → shifts forward. Increase pressure (smaller volume) → shifts to the side with fewer gas moles. Raise temperature → shifts in the heat-absorbing (endothermic) direction, and K changes. A catalyst only helps reach equilibrium faster; it does not move it. Inert gas at constant volume does nothing; at constant pressure it acts like lowering pressure. Only temperature changes the value of K.
• Life Cycle Assessment and Recycling – A life cycle assessment (LCA) adds up the environmental impact of a product over its whole life: getting the raw materials, making it, using it and getting rid of it. LCAs let us compare products fairly. Reusing and recycling materials such as metals, glass and plastics cut the use of raw materials, energy and landfill.
• Machining Ceramics with Abrasives and Tools – Fired ceramics are very hard and brittle. Steel tools are too soft, so ceramics are cut with harder abrasives: alumina, silicon carbide and diamond. Grinding removes material fast, lapping and polishing with finer grit make the surface smooth. Too much pressure cracks the part.
• Main-Group and Transition Elements – Inorganic chemistry is the study of elements and their compounds, apart from carbon chains. Main-group elements sit in groups 1, 2 and 13 to 18. Transition elements sit in groups 3 to 12. Elements in one group behave alike. Transition elements are hard metals with coloured ions, several charges and catalyst power.
• Material and Energy Balances – A balance is bookkeeping for a process. Mass in = mass out + mass stored. Energy in = energy out + energy stored. Before you count, put every number in the same units (SI).
• Materials Innovation: From a Lab Idea to a New Product – Chemical innovation turns an idea into a product in stages: idea, lab test, pilot, factory, market. New materials (alloys, bio-based plastics, smart materials) and new food processes are designed from structure and tested for performance, cost and safety. Scale-up from grams to tonnes and the market decide which ideas succeed.
• Materials Science: Why Materials Behave the Way They Do – Materials science links what a material is made of inside (atoms, bonds, crystals, grains) to how it behaves (strong, bendy, brittle, light). The four families are metals, polymers, ceramics and composites. We measure properties with tests: tensile test (stress = force ÷ area, strain = extension ÷ length, Young's modulus = stress ÷ strain), hardness tests and impact tests. Heat treatment and alloying change the inside and so the properties. New materials include composites, nanomaterials, smart materials and biomaterials. Chemistry also explains art materials: pigments, binders, ceramics, glass and metals.
• Matter and Its Measurement – Matter is anything that has mass and takes up space. It can be solid, liquid or gas, and it can be an element, a compound or a mixture. Chemists measure matter in SI units. Every measurement has some doubt, so we write it with the right number of significant figures, use scientific notation for very big or small numbers, and change units with conversion factors.
• Menus and Cooking by Style: Japanese, Western, Chinese and Other – Every food style has its own way of serving and cooking. Japanese: a set of rice, soup and three dishes, light flavours, dashi and soy. Western: courses in order (starter, soup, main, dessert), roasting, baking and sauces. Chinese: shared dishes with rice and soup, fast wok stir-frying. Others such as the Indian thali or Italian meals use local grains and spices.
• Metal-Aqua Ions: Reactions in Aqueous Solution – In water, a metal ion is held by six water ligands: [M(H₂O)₆]²⁺ or [M(H₂O)₆]³⁺. A 3+ ion pulls hard on the O–H bonds, so it releases H⁺ and its solution is more acidic. Adding OH⁻, NH₃ or CO₃²⁻ removes H⁺ step by step until a neutral hydroxide precipitates. Al(OH)₃ dissolves again in excess OH⁻ (amphoteric); Cu(OH)₂ dissolves in excess ammonia by ligand substitution. 2+ ions give carbonates with CO₃²⁻, but 3+ ions give hydroxides and CO₂ gas.
• Metallurgy: From Metallic Bond to Getting Metals Out of Ores – Metals are made of positive ions held in a sea of free electrons. This metallic bond explains why they shine, conduct and bend. Most metals are found in rocks as compounds called ores. Metallurgy is the science of getting a pure metal out of its ore: first concentrate the ore, then turn the compound into metal by heat (pyrometallurgy), by solutions (hydrometallurgy) or by electricity (electrometallurgy), then refine it. The more reactive a metal is, the harder it is to extract and the faster it corrodes.
• Metals and Non-metals: Properties and the Reactivity Series – Metals are usually shiny, hard, malleable, ductile and good conductors; non-metals are usually dull, brittle and poor conductors. Metals form basic oxides, react with water and dilute acids (giving hydrogen) and a more reactive metal pushes a less reactive one out of its salt solution. Listing metals from most to least reactive gives the reactivity series.
• Methods of Cooking: Moist Heat, Dry Heat, Soups, Salads and Eggs – Cooking methods are grouped by what carries the heat. Moist heat (poach, simmer, boil, steam, pressure cook) uses water or steam, so food stays at about 100 °C or below and does not brown. Dry heat (bake, roast, grill) and frying use hot air, metal or oil at 150–250 °C, so food browns. Soups start from stock and are clear or thick. A salad has a base, body, garnish and dressing. Eggs set as their proteins heat: white at about 63 °C, yolk at about 70 °C.
• Mixtures and Their Separation – A mixture has two or more substances mixed without any fixed ratio, and each keeps its own properties. Homogeneous mixtures (solutions) look the same everywhere; heterogeneous ones do not. By particle size we get solutions (< 1 nm), colloids (1–1000 nm) and suspensions (> 1000 nm). Colloids scatter light (Tyndall effect). Concentration tells how much solute is in a solution. We separate mixtures by using a difference in their parts: evaporation, crystallisation, distillation, chromatography, sublimation, centrifugation and coagulation.
• Mole Concept, Molar Mass and Chemical Formulas – Atoms are far too small to count one by one, so chemists count them by weighing. Atomic masses are given in u, where 1 u is one-twelfth the mass of a carbon-12 atom. One mole is 6.022 × 10²³ particles, and its mass in grams equals the formula mass in u. With moles we can find the percentage of each element in a compound, and work back from percentages to the empirical and molecular formulas.
• Molecular Modelling: Seeing Molecules in 3D – A molecular model is a small stand-in for a molecule that you can see and turn around. Ball-and-stick models show bonds and angles, space-filling models show size, and computer models can also show how a reaction happens step by step. From the model you can predict the shape of a molecule and follow a reaction mechanism.
• Molecular Orbital Theory – In molecular orbital (MO) theory, atomic orbitals of the bonded atoms combine into molecular orbitals that belong to the whole molecule. By LCAO, adding two atomic orbitals gives a lower-energy bonding MO; subtracting gives a higher-energy antibonding MO. Electrons fill MOs by the Aufbau, Pauli and Hund rules. Bond order = ½ (Nb − Na). A positive bond order means the molecule exists; unpaired electrons make it paramagnetic. MO theory explains why He₂ does not exist and why O₂ is paramagnetic.
• Moles: Counting Particles by Weighing – A mole is a fixed number of particles: 6.02 × 10²³ (the Avogadro constant). One mole of a substance has a mass in grams equal to its relative formula mass (Mr). So moles = mass ÷ Mr. Balanced equations tell you the ratio of moles that react, and the reactant that runs out first is the limiting reactant.
• Naming Inorganic Compounds: Formulas and Names – Every compound has a formula and a name that follow fixed rules. Ions carry charges: metals form positive ions, non-metals form negative ions. A compound must have zero total charge, so the numbers of ions are chosen to cancel (criss-cross: swap the charge numbers). The metal name comes first, then the non-metal with an -ide ending (sodium chloride). Groups like sulfate, nitrate and hydroxide keep their own names and need brackets when there is more than one. Roman numerals show the charge of metals such as iron(III), and prefixes (mono, di, tri) name molecules like CO2.
• Nanoparticles: Why Tiny Particles Behave Differently – Nanoparticles are 1 to 100 nanometres (nm) across; 1 nm = 1 × 10⁻⁹ m. They are only a few hundred atoms wide, so a large share of their atoms sit on the surface. Their very high surface area to volume ratio makes them react, block light and kill germs differently from the same material in bulk.
• New Materials and Energy in Space Technology – A rocket must survive great heat, be very light yet strong, carry energy-rich fuel and, for satellites, make power from sunlight. So engineers choose materials with a high melting point for heat shields (titanium alloys, ceramics), low density and high strength for the body (aluminium alloys, carbon-fibre composites), and fuels with high energy per kilogram, such as liquid hydrogen (about 120 MJ/kg) with liquid oxygen, which burns to give only steam. Satellites use solar panels, whose power is about sunlight strength × area × efficiency × cos of the tilt angle. In this project you research one material or energy source, compare facts in a table, judge them against a clear need, and present your findings with sources.
• Nitrogen and Its Compounds – Nitrogen gas (N₂) makes up 78% of air. Its triple bond makes it very unreactive, so plants cannot use it directly. Fixation turns N₂ into compounds: lightning and bacteria do it naturally, and the Haber process makes ammonia (NH₃) industrially. Ammonia is a basic gas that forms ammonium salts. At high temperature N₂ and O₂ form NO, which becomes brown NO₂; NO₂ in rain makes nitric acid and acid rain. Nitric acid is a strong acid and oxidising agent used for fertilisers.
• Nitrogen: the Element, Its Compounds and the Nitrogen Cycle – Nitrogen (N₂) makes up 78% of air, but its triple bond is so strong that most living things cannot use it. Fixation (by root-nodule bacteria, lightning and the Haber process) turns it into ammonia. From ammonia we make nitric acid and fertilisers. In nature, nitrogen moves in a cycle: fixation → nitrification → absorption by plants → animals → decay (ammonification) → denitrification back to air.
• NMR Spectroscopy: Carbon-13 and Proton NMR – Nuclear magnetic resonance (NMR) uses a strong magnet and radio waves to see the different environments of ¹H or ¹³C atoms in a molecule. Each environment gives a peak at its own chemical shift δ (ppm), measured from TMS at δ = 0. In ¹³C NMR, the number of peaks = number of carbon environments. In ¹H NMR, the peak area (integration) gives the ratio of H atoms, and the n + 1 rule tells how many H are on the neighbouring carbon.
• Non-metals and Their Reactions – Non-metals are elements like hydrogen, carbon, nitrogen, oxygen, sulfur and chlorine. They are usually dull, brittle and poor conductors. They share electrons with each other (covalent bonds), so hydrogen forms HCl, H₂O, NH₃, CH₄ and H₂S. With metals, chlorine and sulfur take electrons and make ionic salts such as NaCl, MgCl₂, Na₂S and FeS.
• Occurrence and Extraction of Metals – Metals are found in the earth's crust, mostly as compounds called minerals; a mineral from which a metal can be taken out profitably is an ore. Extraction has three stages: enrichment (removing gangue), getting the crude metal (roasting or calcination, then reduction or electrolysis, chosen by the metal's place in the reactivity series), and refining (usually electrolytic).
• Ocean Acidification and Coral Bleaching – The ocean absorbs about a quarter of the CO2 we release. In water, CO2 forms an acid, which lowers pH and removes carbonate that shells and coral skeletons need. Separately, warm water makes corals lose their colour (bleaching). Both threaten reef ecosystems.
• Optical Isomerism: Chiral Centres and Enantiomers – Optical isomers are stereoisomers that are mirror images of each other but cannot be placed on top of each other. They happen when a carbon atom has four different groups: a chiral centre. The two forms, called enantiomers, have the same physical and chemical properties except that they rotate plane-polarised light by the same angle in opposite directions, and they can act differently in the body. A 50 : 50 mix (racemic mixture) has no overall effect on the light.
• Organic Analysis: Identifying Organic Compounds – Chemists identify an unknown organic compound in two ways. Quick test-tube reactions show which functional group is present: bromine water goes colourless with alkenes, acidified dichromate turns green with primary and secondary alcohols and aldehydes, Tollens' reagent gives a silver mirror with aldehydes, and sodium carbonate fizzes with carboxylic acids. Instruments then confirm it: a mass spectrum gives the relative molecular mass from the molecular ion peak (M⁺), and an infrared spectrum shows which bonds are present.
• Organic Compounds: Structure, Formulae and Classification – Organic chemistry is the chemistry of carbon compounds. A carbon atom has 4 outer electrons, so it always makes 4 bonds. This is called tetravalence. With 4 single bonds the carbon is sp³ and its bonds point to the corners of a tetrahedron (109.5°). With one double bond it is sp² and flat (120°). With a triple bond it is sp and straight (180°). We can draw the same molecule in several ways: complete, condensed, bond-line and 3D wedge-dash. Finally we sort organic compounds into open-chain and ring compounds, and into families by their functional group. Members of one family form a homologous series, where each member differs by one CH₂.
• Organic Products in Industry, Environment and the Body – Organic compounds make our plastics, fibres, fuels, medicines and solvents. Some halogen products, like freons (CFCs), reach the stratosphere and destroy ozone because one chlorine atom can break thousands of ozone molecules. In the body, hydrolysis splits big food molecules and oxidation of glucose releases energy. Proteins fold into primary, secondary and tertiary structures, and the shape lets enzymes speed up reactions and hormones carry messages.
• Organic Synthesis – Organic synthesis means making a target molecule from simpler starting materials through a planned series of reactions. Each step changes one functional group and needs the right reagent and conditions (temperature, catalyst, solvent, reflux or distillation). Families of compounds are linked (alkane ↔ alkene → haloalkane → alcohol → aldehyde/ketone → carboxylic acid → ester; haloalkane → amine or nitrile), so we can move between them like stations on a map. Chemists plan backwards from the target (retrosynthesis), prefer short routes with high yield and atom economy, and then carry out the reaction, separate, purify and check the product.
• Overview of the Ceramic Industry – The ceramic industry makes hard, heat-proof, non-metal products from earth materials. The steps are: prepare raw materials, mix, shape, dry and fire in a kiln. Four big families are pottery, glass, refractories and cement. New fine ceramics are used in electronics and machines.
• Oxidation and Reduction: Corrosion and Rancidity – Oxidation is gain of oxygen (or loss of hydrogen); reduction is loss of oxygen (or gain of hydrogen). They always happen together, so these are called redox reactions. The substance that gives oxygen is the oxidising agent; the one that takes it is the reducing agent. In daily life, oxidation causes corrosion (like rusting of iron) and rancidity (fats and oils going stale).
• Oxidation Number: Rules, Calculation and Types of Redox – Oxidation number is a pretend charge. We imagine that shared electrons go fully to the atom that pulls harder. A few simple rules give the number for any atom, and all numbers in a species add up to its charge. If an atom’s oxidation number goes up, it is oxidised; if it goes down, it is reduced. With this we can sort redox reactions into combination, decomposition, displacement and disproportionation.
• Oxides of Carbon: Carbon Dioxide and Carbon Monoxide – Carbon makes two common oxides. CO₂ is heavy, does not burn, dissolves a little in water, and traps heat in the air. CO is a poisonous gas from incomplete burning; it sticks to haemoglobin and, as a reducing agent, takes oxygen from metal oxides.
• Oxides: Acidic, Basic, Amphoteric and Neutral – An oxide is a compound of oxygen with one other element. Metal oxides are usually basic, non-metal oxides are usually acidic, a few (Al₂O₃, ZnO) are amphoteric and a few (CO, N₂O, water) are neutral. A simple test: drop the oxide in water, in acid and in base and see what happens.
• Oxygen – Oxygen (O, atomic number 8) is a colourless, odourless gas that makes up about 21% of air by volume. It exists as O₂ molecules. In the lab it is made by breaking down hydrogen peroxide with a manganese dioxide catalyst (2H₂O₂ → 2H₂O + O₂) or by heating potassium manganate(VII); in industry it comes from fractional distillation of liquid air. It relights a glowing splint. Oxygen supports burning and respiration and reacts with most elements to form oxides: metal oxides are basic, non-metal oxides are mostly acidic. Its other form, ozone (O₃), shields Earth from UV rays.
• Painkillers: How Aspirin Is Made and Checked – Drug molecules are built from functional groups such as -OH, -COOH, ester and amide. Retrosynthesis works backwards from the target. Aspirin is made by acetylating the -OH of salicylic acid with acetic anhydride (an ester forms by addition then elimination, with an acid catalyst). Purity is checked by melting point, an iron(III) chloride test and thin-layer chromatography, where Rf = distance of spot / distance of solvent front.
• Percentage Yield and Atom Economy – Percentage yield compares the product you actually got with the most you could get: actual ÷ theoretical × 100. Atom economy tells you what share of the reactant atoms end up in the product you want: Mr of wanted product ÷ total Mr of reactants × 100. Both help make chemistry cheaper and less wasteful.
• Period 3 Elements and Their Oxides – Across Period 3 (Na, Mg, Al, Si, P, S) the elements burn in oxygen to make oxides. The oxides change from ionic (Na₂O, MgO, Al₂O₃) to giant covalent (SiO₂) to small molecules (P₄O₁₀, SO₂, SO₃). That structure decides their melting point and how they behave with water, acids and bases: basic on the left, amphoteric Al₂O₃ in the middle, acidic on the right.
• Periodic Classification of Elements – Scientists sorted elements so that similar ones sit together. Döbereiner made triads, Newlands found that every 8th element repeats (octaves), Mendeleev arranged elements by atomic mass and left gaps for unknown ones. Moseley showed that atomic number is the real key. Modern periodic law: the properties of elements are a periodic function of their atomic numbers. The modern table has 7 periods and 18 groups. Elements with Z > 100 get temporary IUPAC names built from digit roots (nil, un, bi, tri, quad, pent, hex, sept, oct, enn) plus -ium.
• Periodic Trends in Properties – Two forces decide almost every trend: the pull of the nucleus (effective nuclear charge) and the distance of the outer shell. Across a period the nuclear pull grows while the shell stays the same, so atoms shrink, ionisation enthalpy rises, electron gain enthalpy becomes more negative and electronegativity rises. Down a group a new shell is added, so atoms grow and these values fall. Cations are smaller and anions bigger than their atoms. Valence follows the outer electrons; metallic reactivity is highest at the bottom left and non-metallic reactivity at the top right.
• Phase Changes: Latent Heat, Vapour Pressure and Phase Diagrams – A phase change is when matter moves between solid, liquid and gas. During a change the temperature stays the same; the heat used is latent heat (Q = m × L). A liquid boils when its vapour pressure equals the pressure above it, so boiling point falls with lower pressure. A phase diagram maps the state at every temperature and pressure; its lines meet at the triple point and the liquid–gas line ends at the critical point.
• Phase Diagrams of Minerals – A mineral can have the same chemical formula but a different crystal structure, and which one is stable depends on temperature and pressure. A phase diagram is a map of temperature (across) against pressure or depth (up). Each zone shows the stable form. On a line two forms live together, and at the triple point three do. The three forms of Al2SiO5 (andalusite, kyanite, sillimanite) show it well: andalusite at low pressure, kyanite at high pressure, sillimanite at high temperature. If conditions change, the mineral slowly changes too.
• Phenols – A phenol has –OH joined straight to a benzene ring carbon. It is made from chlorobenzene (NaOH, 623 K, 300 atm), benzenesulphonic acid, diazonium salts or cumene. Phenol is more acidic than alcohols and water because the phenoxide ion spreads its negative charge over the ring; electron-pulling groups like –NO₂ (at ortho/para) raise acidity, electron-pushing groups like –CH₃ lower it. The –OH group activates the ring at ortho and para positions: bromine water gives 2,4,6-tribromophenol, dilute HNO₃ gives o- and p-nitrophenol, NaOH + CO₂ gives salicylic acid (Kolbe) and CHCl₃ + NaOH gives salicylaldehyde (Reimer–Tiemann).
• Phosphorus and Its Compounds: Phosphides, Phosphine and Fertilisers – Phosphorus is a non-metal in group 15, just below nitrogen. White phosphorus is P₄ (a strained 3D pyramid) and catches fire in air; red phosphorus is a stable chain. It burns to P₄O₆ or P₄O₁₀, and P₄O₁₀ with water gives phosphoric acid. Metal phosphides such as Ca₃P₂ give poisonous phosphine PH₃ with water. Phosphates from phosphoric acid are the key fertilisers.
• Physical and Chemical Changes – In a physical change, a substance changes its shape, size or state, but no new substance forms; the particles stay the same. In a chemical change, the atoms rearrange to make one or more new substances with new properties. Signs include gas, colour change, heat or light, a precipitate and a new smell. In both kinds, the total mass stays the same.
• Pollution: Air, Water, Land and Noise – Pollution is the release of harmful substances or energy into air, water or land faster than nature can remove them. Main air pollutants are particulates, carbon monoxide, sulfur dioxide, nitrogen oxides and ozone; they cause acid rain and smog. Water pollution by nutrients, sewage and toxins causes eutrophication and biomagnification. We reduce pollution by cutting emissions at the source, cleaning waste before release, and the 3Rs.
• Polyhalogen Compounds – Polyhalogen compounds have more than one halogen on carbon. Dichloromethane is a solvent and paint remover; chloroform is a solvent that turns into poisonous phosgene in air and light; iodoform is a yellow antiseptic; carbon tetrachloride was a solvent and fire extinguisher; freons were coolants that destroy ozone; DDT is an insecticide that builds up in food chains.
• Polymers – A polymer is a giant molecule made by joining many small molecules (monomers) into a long chain. In addition polymerisation, monomers with a C=C double bond open up and link with nothing lost (ethene → poly(ethene)). In condensation polymerisation, two kinds of monomer with reactive groups at both ends join and give off a small molecule such as water at every link (nylon, polyester). Nature makes polymers too: starch, cellulose, proteins, DNA, rubber. Separate chains give thermoplastics that melt and can be recycled; cross-linked chains give thermosets that never melt. Most plastics do not rot, so we must reduce, reuse and recycle them.
• Potassium Dichromate and Potassium Permanganate – K₂Cr₂O₇ (orange, Cr +6) is made from chromite ore; in water it exists as dichromate in acid and yellow chromate in base. KMnO₄ (purple, Mn +7) is made from pyrolusite (MnO₂). Both are strong oxidising agents: acidified dichromate takes 6 electrons per ion (→ green Cr³⁺); permanganate takes 5 in acid (→ Mn²⁺), 3 in neutral/weak base (→ MnO₂) and 1 in strong alkali (→ MnO₄²⁻).
• Practice of Food Chemistry – In real practice you plan the analysis, prepare a fair sample, run every test twice, calculate and check that the parts add up, then write a clear report. Hygiene testing adds checks for safety: pH and acidity, the number of microbes (CFU per gram from a plate count), indicator germs and simple tests for added or harmful substances.
• Preparing and Testing Carbon Dioxide in the Lab – Marble (calcium carbonate) reacts with dilute hydrochloric acid to give carbon dioxide, calcium chloride and water: CaCO3 + 2HCl -> CaCl2 + H2O + CO2. CO2 is colourless, odourless and heavier than air, so it is collected by upward delivery. It turns limewater milky, puts out a burning splint and turns moist blue litmus red.
• Preparing Carbon Dioxide in the Lab – In the lab, CO₂ is made by pouring dilute hydrochloric acid on marble chips (CaCO₃): CaCO₃ + 2HCl → CaCl₂ + H₂O + CO₂. The gas is heavier than air, so it is collected by upward delivery in a jar. A burning splint goes out in it and limewater turns milky.
• Principles of Food Processing and Raw Materials – Raw foods such as grain, milk, fruit and meat contain water, starch, protein, fat and living germs. Food processing uses a few simple ideas to make them safe, tasty and long-lasting: change size (cut, grind), use heat (kill germs, cook), remove water (dry), cool, and mix. Each raw material behaves differently, so we choose the method to match it.
• Processing Textile Products: Dyeing and Finishing – Cloth straight from the loom (grey cloth) is processed in steps: prepare (wash, bleach), dye or print, finish with machines (calendering, raising, shearing, heat setting, sanforising) and treat the surface (mercerising, water repellent, flame retardant, anti-crease). Each step changes colour, feel, size or safety of the cloth. Colour can be added to fibre, yarn, fabric or the finished garment.
• Proteins and Enzymes: Amino Acids, Peptide Bond, Structure and Denaturation – Proteins are long chains of amino acids. Each amino acid has an amino group (–NH₂), an acid group (–COOH), a hydrogen and a side group R on one carbon. The –COOH of one amino acid joins the –NH₂ of the next, losing water, to make a peptide bond (–CO–NH–). The order of amino acids is the primary structure; coils and sheets are the secondary structure; the full 3D fold is the tertiary structure; several chains together make the quaternary structure. Heat, acid or alcohol can undo the folds (denaturation). Enzymes are mostly proteins that speed up reactions in the body.
• Purification and Analysis of Organic Compounds – A compound made in a lab or taken from a plant is never pure at first. We purify it by using a difference in properties: solubility (crystallisation), boiling point (distillation and its types), sublimation, solubility in two liquids (differential extraction) or how strongly it sticks to a surface (chromatography). A pure solid has a sharp melting point. Next we find which elements are present (qualitative analysis): carbon and hydrogen by heating with copper(II) oxide, and N, S, halogens and P by Lassaigne's sodium fusion test. Finally we find how much of each element is present (quantitative analysis): Liebig's method for C and H, Dumas and Kjeldahl methods for N, Carius method for halogens and S, and oxygen by difference. From the percentages we can get the empirical formula.
• Quantum Mechanical Model: Quantum Numbers, Orbitals and Configuration – Moving electrons also behave like waves: λ = h/mv (de Broglie). Because of this we cannot know exact position and momentum together: Δx·Δp ≥ h/4π (Heisenberg). So we drop sharp orbits and use orbitals: regions where the electron is likely to be, from the Schrödinger equation (ψ² gives probability). Four quantum numbers describe each electron: n (shell, size, energy), l (subshell, shape: 0 to n−1), mₗ (orientation: −l to +l) and mₛ (spin: +½ or −½). s is spherical, p is dumbbell, d is mostly four-lobed. Electrons fill by the Aufbau (n + l) rule, Pauli principle (max 2, opposite spins) and Hund's rule (single first). Half-filled and fully filled subshells are extra stable, so Cr is 3d⁵4s¹ and Cu is 3d¹⁰4s¹.
• Raoult's Law: Vapour Pressure, Ideal and Non-Ideal Solutions – In a closed jar, some molecules of a liquid escape and push on the walls: that push is the vapour pressure. Raoult's law says each volatile part of a solution gives a vapour pressure equal to its own pure vapour pressure times its mole fraction: pA = xA·pA°. Solutions that obey this at every mix are ideal (ΔmixH = 0, ΔmixV = 0). Others deviate: positive deviation (weaker A–B pull, more vapour, forms minimum-boiling azeotropes) or negative deviation (stronger A–B pull, less vapour, forms maximum-boiling azeotropes).
• Rate of a Chemical Reaction – The rate of a reaction tells how fast a reactant is used up or a product is made, per unit time. Rate = −Δ[R]/Δt = +Δ[P]/Δt (unit mol L⁻¹ s⁻¹). The rate law, rate = k[A]^x[B]^y, is found by experiment; x + y is the order. Molecularity is the number of particles that collide in one elementary step.
• Rates of Reaction – The rate of a reaction is how fast reactants are used up or products are made. Rate = change in amount ÷ time. Particles must collide with at least the activation energy to react. More frequent, more energetic collisions mean a faster rate. Raising concentration (or gas pressure), temperature or surface area, or adding a catalyst, makes reactions faster. A catalyst gives a path with lower activation energy and is not used up.
• Redox Reactions and Electrode Processes – If we keep the oxidation half and the reduction half in separate beakers and join them with a wire and a salt bridge, the electrons travel through the wire. This makes a cell. Each metal in its ion solution is an electrode with its own potential. The hydrogen electrode is set at 0.00 V. Listing electrodes by standard potential gives the electrochemical series, and E°cell = E°(cathode) − E°(anode).
• Redox Reactions: From Oxygen to Electron Transfer – Oxidation first meant adding oxygen or removing hydrogen. Reduction meant the opposite. Today we use a bigger idea: oxidation is losing electrons and reduction is gaining electrons. Both always happen together, so we call them redox reactions. A more active metal gives electrons to the ion of a less active metal.
• Refractories – Refractories are heat-proof ceramic materials that keep their shape and strength at very high temperature. They line furnaces, kilns and ladles for steel, cement, glass and ceramics. Main types are fireclay, high-alumina, silica and magnesia bricks, plus unshaped castables. Choose the brick by the heat and the slag (acidic or basic).
• s- and p-Block Elements: Configuration and Trends – s-block = groups 1 and 2 (outer configuration ns¹ or ns²); p-block = groups 13 to 18 (ns² np¹⁻⁶). Down a group, atoms get bigger and ionisation enthalpy falls; across a period, atoms get smaller and ionisation enthalpy rises overall. Small ions have large hydration enthalpy. The first element of each group behaves differently (small size, high charge density, no d orbitals) and often resembles the element diagonally below it.
• Safety, Waste Treatment and Recycling in the Ceramic Industry – A ceramic factory makes dust, smoke and dirty water, and it also makes scrap. Masks, ventilation and filters protect workers and air. Settling tanks clean water so it can be reused. Broken ware is crushed to grog and mixed back into new clay, which saves raw material and cuts waste.
• Salt from Sea Water and the Ocean's Resources – Sea water holds about 3.5% dissolved salt. In salt pans, sun and wind evaporate only the water. The brine gets stronger until the salt cannot stay dissolved and crystallises. Raw salt is purified by dissolving, filtering and evaporating. The sea also gives food, minerals, energy and fresh water, which must be used with care.
• Salts and Important Chemicals from Common Salt – A salt is made when an acid and a base neutralise each other. Salts can be neutral, acidic or basic. From common salt we make sodium hydroxide, bleaching powder, baking soda and washing soda, and some salts hold water inside their crystals.
• Silica, Glass, Limestone, Gypsum and Fertilisers – Silica (SiO₂) is sand and quartz and is the raw material of glass. Limestone, marble and chalk are all calcium carbonate (CaCO₃); acid makes them fizz. Boiling removes temporary hardness as CaCO₃ scale. Gypsum (CaSO₄·2H₂O) loses water to make plaster of Paris, which sets when water is added. Fertilisers give plants nitrogen, phosphorus and potassium.
• Silicate Minerals: How One Tiny Brick Builds Most of the Earth's Crust – Silicate minerals are minerals built from silicon and oxygen. Their basic brick is the SiO₄ tetrahedron: one silicon atom with four oxygen atoms around it. Tetrahedra can stay alone or share oxygens to form chains, sheets or a 3D framework. The way they link decides the mineral's shape, cleavage and hardness. Silicates make up over 90% of the Earth's crust.
• Smart and Modern Materials – A smart material changes one of its properties (shape, colour, transparency or electrical output) when something in its surroundings changes, such as temperature, light, force or a voltage. The change is reversible. Key examples: shape-memory alloys like nitinol return to a remembered shape when heated; thermochromic pigments change colour with temperature; photochromic materials darken in UV light; piezoelectric crystals make a voltage when squeezed. Modern materials, which are not necessarily smart, include composites such as carbon-fibre and glass-fibre reinforced plastic, graphene and other nanomaterials, technical and smart textiles, liquid crystals and bio-based plastics. Engineers choose materials by their properties, cost and effect on the environment.
• Soaps and Detergents: How a Micelle Cleans Oily Dirt – A soap molecule is the sodium or potassium salt of a long-chain fatty acid. Its ionic head loves water and its long hydrocarbon tail loves oil. In water, soap molecules bury their tails in oily dirt and point their heads outward, forming a micelle that is rinsed away. In hard water, Ca²⁺ and Mg²⁺ ions turn soap into insoluble scum; detergents (sulphonate or ammonium salts) do not form scum, so they work in hard water too.
• Soil Acidity and Plant Growth: A Project – Soil pH tells us if the soil is acidic (below 7), neutral (about 7) or alkaline (above 7). Plants take food from the soil water, and pH changes which nutrients dissolve. Most crops grow best between pH 6 and 7.5, tea likes acid soil near pH 5, and in very acidic or very alkaline soil many plants grow badly. To measure pH we mix soil with distilled water, let it settle and test the clear water with indicator paper or solution or a pH meter. Acidic soil is treated with lime (a base) and alkaline soil with sulfur, compost or acidic fertiliser. In the project we test, compare plants with a fair test, and decide how to treat the soil.
• Solutions: How Things Dissolve and How Much Can Dissolve – A solution is a uniform mixture of a solute dissolved in a solvent. In water (an aqueous solution) the solute breaks into particles too small to see, so it never settles and passes through filter paper. Concentration tells how much solute is present (mass % = solute ÷ solution × 100). Solubility is the most that can dissolve in 100 g of solvent at a given temperature; beyond it the solution is saturated. Evaporation, crystallisation and distillation separate solutions.
• Solutions: Types, Concentration and Henry's Law – A solution is an even mix of two or more substances. The part in bigger amount is the solvent, the smaller part is the solute. We tell 'how strong' a solution is with concentration terms: mass %, volume %, ppm, mole fraction, molarity (per litre of solution) and molality (per kg of solvent). Solids usually dissolve more when hot. Gases dissolve more when their pressure is high (Henry's law, p = KH·x) and less when it is hot.
• Spectroscopy and the Beer-Lambert Law – Spectroscopy studies how matter absorbs or gives out light. Light is made of photons with energy E = hf = hc/λ. Atoms give line spectra because electrons jump between fixed energy levels. A coloured solution absorbs its complementary colour. A spectrophotometer measures transmittance T = I/I₀ and absorbance A = log₁₀(I₀/I). The Beer-Lambert law says A = εlc, so a calibration line of A against c lets us find an unknown concentration.
• Spontaneity, Entropy and Gibbs Energy – A spontaneous process happens by itself, without outside help (it may be fast or slow). ΔH alone cannot predict it: ice melts on its own although it takes in heat. Entropy S measures how spread out energy and matter are; ΔS = q_rev/T. Second law: for any spontaneous change, ΔS_total = ΔS_system + ΔS_surroundings > 0; at equilibrium it is zero. Third law: a perfect crystal at 0 K has zero entropy. Gibbs energy G = H − TS combines both: at constant T and p, ΔG = ΔH − TΔS, and ΔG < 0 means spontaneous, ΔG = 0 means equilibrium. It links to the equilibrium constant by ΔG° = −RT ln K = −2.303 RT log K.
• Stereochemistry: Chirality, R/S and Z/E – Stereochemistry studies how atoms sit in 3D space. A carbon with four different groups is chiral and has a non-overlapping mirror image (enantiomer). CIP rules rank groups by atomic number; with the lowest priority pointing away, 1 to 2 to 3 clockwise is R and anticlockwise is S. Around a double bond, the higher-priority groups on the same side give Z, on opposite sides E.
• Stoichiometry, Limiting Reagent and Concentration Terms – Stoichiometry means measuring the amounts of substances in a reaction. A balanced equation acts like a recipe: its numbers give the mole ratio of reactants and products. With it you can change any mass into moles, use the ratio, and change back to the mass of any other substance. The reactant that runs out first is the limiting reagent; it decides how much product forms. For solutions, concentration is given as mass percent, mole fraction, molarity or molality.
• Structure and Properties of Ceramics – Ceramics are made of ions held by strong bonds. The ratio of cation size to anion size decides how many neighbours (coordination number) each ion has: 0.225 to 0.414 gives 4, 0.414 to 0.732 gives 6, above 0.732 gives 8. If the ions sit in a repeating pattern the ceramic is a crystal; if the pattern is random it is a glass. Strong bonds make ceramics hard and heat-proof; fixed ions make them brittle.
• Structure and Properties of Materials – What a material does (its property) depends on what is inside it (its structure). Metals conduct because they have free electrons. Solid salt does not conduct because its ions are locked, but melted or dissolved salt does because the ions can move. Sugar dissolves but stays neutral, so it does not conduct. Long plastic chains are strong; a chain with weak links can be cut by bacteria, so it is biodegradable. Chains with OH groups attract water, so they swell into a gel. To predict a property, ask: what are the particles, do they carry charge, can they move, and how strongly do they attract each other or water?
• Structure of the Atom – An atom has a tiny, heavy, positive nucleus made of protons and neutrons. Electrons move around it in fixed shells K, L, M, N. The number of protons (Z) tells the element; protons + neutrons give the mass number (A). Outer electrons decide valency. Isotopes share Z; isobars share A.
• Subatomic Particles and Early Atomic Models – An atom is made of three main particles. Electrons (charge −1, very light) were found with cathode ray tubes. Protons (charge +1, about 1836 times heavier) were found in canal rays. Neutrons (no charge, mass about a proton) were found by Chadwick in 1932. Thomson pictured the atom as a ball of positive charge with electrons stuck in it. Rutherford fired alpha particles at gold foil and found a tiny, heavy, positive nucleus with electrons moving outside. Atomic number Z = number of protons; mass number A = protons + neutrons. Rutherford's model could not explain why atoms are stable or why they give line spectra, which led to Bohr's model.
• Sulfur and Its Compounds – Sulfur (S, atomic number 16, electrons 2, 8, 6) is a yellow non-metal in group 16 with oxygen. Solid sulfur is made of S₈ rings and has rhombic and monoclinic forms. It burns to sulfur dioxide (SO₂), a gas that bleaches, kills microbes and causes acid rain. In the contact process SO₂ is oxidised to SO₃ over a V₂O₅ catalyst and turned into sulfuric acid (H₂SO₄), a strong acid that also dehydrates and oxidises. Hydrogen sulfide (H₂S) smells of bad eggs. Sulfate ions are found with barium chloride: a white precipitate of BaSO₄ that does not dissolve in dilute HCl.
• Surface Treatments and Finishes – A finish is a coating or treatment applied to the surface of a material. Finishes do two jobs: protect (against rust, rot, water, UV light, wear, stains, fire) and improve looks or feel (colour, shine, texture). Woods are painted, varnished, oiled, waxed, stained or treated with preservatives. Metals are painted over a primer, powder coated, galvanised with zinc, electroplated or, for aluminium, anodised. Many polymers are self-finishing from the mould, but can be polished or printed. Papers and boards are laminated or varnished. Fabrics get chemical or mechanical finishes such as water-repellent, flame-retardant, anti-crease or brushed, and laminated membrane fabrics keep rain out while letting sweat vapour escape. Choosing a finish means matching it to the material, where the product is used, its cost, how it is applied and its environmental impact.
• Synthesised Materials: Plastics and Medicines – Synthesised (man-made) materials are made in labs and factories by chemical reactions. Plastics are long chains (polymers) built from small units (monomers). Medicines are molecules with a shape that fits a target in the body, like a key in a lock.
• Temperature, Activation Energy, Catalysts and Collision Theory – Reactant molecules must climb an energy hill, the activation energy Ea, to become products. The Arrhenius equation k = A e^(−Ea/RT) shows that a small rise in temperature lets many more molecules cross, so k rises fast. A catalyst gives a lower hill without changing ΔH. Collision theory: molecules react only when they collide with enough energy and the right orientation.
• Tests for Gases: How to Identify Hydrogen, Oxygen, Carbon Dioxide, Chlorine and Ammonia – Most gases have no colour, so we identify them with a simple test that gives a clear sign. Hydrogen: a lit splint gives a squeaky pop. Oxygen: a glowing splint relights. Carbon dioxide: limewater turns milky. Chlorine: damp blue litmus turns red, then white (bleached). Ammonia: damp red litmus turns blue. Water vapour turns white anhydrous copper(II) sulfate blue. Sulfur dioxide turns acidified potassium manganate(VII) from purple to colourless.
• Tests for Ions: Flame Tests, Hydroxide Precipitates, Anion Tests and Flame Emission Spectroscopy – An ionic compound is made of a positive metal ion and a negative ion. Each can be found with its own test. Flame tests show metal ions by colour: lithium crimson, sodium yellow, potassium lilac, calcium orange-red, copper green. Sodium hydroxide gives coloured hydroxide precipitates: Cu²⁺ blue, Fe²⁺ green, Fe³⁺ brown, Al³⁺, Ca²⁺ and Mg²⁺ white (only Al(OH)₃ dissolves in excess). Carbonates fizz with acid and the gas turns limewater milky. Halides with acidified silver nitrate: chloride white, bromide cream, iodide yellow. Sulfates with acidified barium chloride: white precipitate. Instruments such as flame emission spectroscopy are faster, more sensitive and can measure amounts.
• The Carbon Family (Group 14) – Carbon, silicon, germanium, tin and lead sit in one group and each has 4 outer electrons. Going down, atoms get bigger and more metallic: carbon is a non-metal, silicon and germanium are metalloids, tin and lead are metals. Their oxides differ: CO2 is a gas of small molecules, SiO2 is a hard giant network. Carbon makes carbonates; silicon makes silicates, glass, cement and ceramics.
• The Chemistry of Cosmetics – A cosmetic is a product put on skin, hair, nails or teeth to clean, protect or change how they look. Most creams are emulsions: tiny oil drops held in water (or water in oil) by an emulsifier. Sunscreens absorb or reflect UV light: UVB burns the surface, UVA goes deeper and ages skin; SPF tells how much UVB is blocked. Antioxidants mop up free radicals. Using cosmetics responsibly means knowing your skin type, reading the ingredient list, patch testing and thinking about cost and the environment.
• The Composition and Evolution of Earth's Atmosphere – Today dry air is about 78% nitrogen, 21% oxygen and about 1% argon, with tiny amounts of carbon dioxide (about 0.04%) and other gases. Earth's early atmosphere came from volcanoes: mostly carbon dioxide and water vapour, with little or no oxygen. As Earth cooled, water vapour condensed into oceans and some CO2 dissolved. Algae and later plants made oxygen by photosynthesis. Carbon was locked away in sedimentary rocks such as limestone and in fossil fuels, so CO2 fell.
• The Haber Process and NPK Fertilisers – The Haber process makes ammonia from nitrogen (from air) and hydrogen (usually from natural gas): N₂ + 3H₂ ⇌ 2NH₃. The forward reaction is exothermic and reversible. A compromise of about 200 atm, about 450 °C and an iron catalyst gives a reasonable yield at a fast rate. Ammonia is cooled to a liquid and removed; unreacted gases are recycled. Ammonia is used to make nitrogen fertilisers such as ammonium nitrate. NPK fertilisers supply nitrogen, phosphorus and potassium to help crops grow.
• The Ideal Gas Law: PV = nRT – Boyle (P ∝ 1/V), Charles (V ∝ T) and Avogadro (V ∝ n) combine into one rule: PV = nRT, with R = 8.314 J mol⁻¹ K⁻¹. Temperatures must be in kelvin. It also gives molar mass M = dRT/P. Real gases follow it well at low pressure and high temperature; at high pressure or low temperature they deviate, measured by Z = PV/nRT.
• The Law of Conservation of Mass – In a chemical reaction, mass is not made and not destroyed. The total mass of the substances before the reaction equals the total mass after it. This is true because atoms are only rearranged: the same atoms, in the same numbers, are there before and after. If a gas leaves or joins from the air, the mass on a balance seems to change, but when we count every substance, the total stays the same. A related rule, the law of definite proportions, says a pure compound always has its elements in the same mass ratio (water is always 1 g hydrogen to 8 g oxygen).
• The Periodic Table: Groups, Periods and Trends – The periodic table lists elements in order of atomic number (number of protons). Rows are periods; columns are groups. Elements in the same group have the same number of outer-shell electrons, so they react in similar ways. Mendeleev built an early table, left gaps for unknown elements and predicted their properties. Metals are on the left and bottom, non-metals on the right and top. Group 0 noble gases have full outer shells and are unreactive. Group 1 alkali metals get more reactive going down. Group 7 halogens get less reactive going down.
• The pH Scale: How Strong Is an Acid or a Base? – pH is a number from 0 to 14 that tells how acidic or basic a solution is. 7 is neutral, below 7 is acidic and above 7 is basic. Each step down means ten times more H⁺ ions.
• The Solid State – In a solid the particles sit close together and only vibrate in place, so a solid keeps its shape. In crystalline solids the particles repeat in a regular pattern; in amorphous solids (like glass) they do not. The smallest repeating box is the unit cell. Simple cubic holds 1 particle, body-centred cubic 2 and face-centred cubic 4. By the bonds holding them, crystals are ionic, covalent network, molecular or metallic.
• Titration: Finding an Unknown Concentration – Titration finds the unknown concentration of a solution. A solution of known concentration (the titrant) is added slowly from a burette to a measured volume of the unknown (the analyte) until the reaction is just complete. An indicator or a meter shows this point. At the equivalence point the moles react in the ratio of the equation, so for a 1 : 1 reaction C₁V₁ = C₂V₂.
• Towards Bohr's Model: Light, Photons, Spectra and Bohr's Atom – Light is an electromagnetic wave: c = νλ, and wavenumber ν̄ = 1/λ. But some facts need particles: Planck said energy comes in packets (quanta) E = hν. Einstein used photons to explain the photoelectric effect: an electron comes out only if hν is above the work function W₀ = hν₀, and its kinetic energy is hν − hν₀. Atoms give line spectra, which means electron energies are fixed. Bohr put the electron on fixed orbits with angular momentum nh/2π; energy Eₙ = −2.18 × 10⁻¹⁸ Z²/n² J, radius rₙ = 52.9 n²/Z pm. A jump between orbits gives a photon, and 1/λ = R_H(1/n₁² − 1/n₂²) gives the Lyman, Balmer, Paschen, Brackett and Pfund series. Bohr works only for one-electron species.
• Transition Elements (d-Block) – Transition elements are the d-block metals of groups 3 to 12. Their last electron goes into an inner (n−1)d orbital. Because the d-orbitals are only partly filled, they show many oxidation states, coloured ions, magnetism (μ = √(n(n+2)) BM), good catalytic power and easy alloy formation.
• Transition Metals: Basic Properties – Transition metals are the block of elements in the middle of the periodic table (the d-block), such as iron, copper, nickel, chromium, manganese, cobalt, titanium, silver and gold. Compared with Group 1 metals like sodium, they have much higher melting points, higher densities, and are harder and stronger; they react far more slowly with water and oxygen. They form ions with different charges (Fe²⁺ and Fe³⁺, Cu⁺ and Cu²⁺), make coloured compounds, and are useful catalysts. Their atoms fill an inner d sub-shell; chromium and copper 'borrow' an electron from 4s to get a half-full or full 3d. Many of their ions also form complex ions with water or ammonia.
• Types of Chemical Reactions – Most reactions fit a few patterns. Combination: A + B → AB. Decomposition: AB → A + B. Displacement: A + BC → AC + B, where the more reactive A pushes out B. Double displacement: AB + CD → AD + CB, where partners swap; if an insoluble solid forms it is a precipitation reaction. Reactions that give out heat are exothermic; those that take in heat are endothermic.
• Types of Organic Reactions and Their Mechanisms – Organic reactions are sorted by what happens to the molecule: substitution (one atom swaps for another), addition (atoms join across a double bond), elimination (a small molecule leaves and a double bond forms), condensation (two molecules join and lose water) and hydrolysis (water splits a molecule). The mechanism shows how electrons move: bonds break homolytically (radicals) or heterolytically (ions), and nucleophiles (electron-rich) attack electrophiles (electron-poor). Esters form from acids and alcohols in a reversible, acid-catalysed condensation and are split by hydrolysis.
• Valence Bond Theory and Hybridisation – Valence bond theory says a covalent bond forms when half-filled orbitals of two atoms overlap and the two electrons pair up with opposite spins. Head-on overlap makes a strong σ (sigma) bond; side-on overlap of p orbitals makes a weaker π (pi) bond. To explain real shapes, an atom first mixes its orbitals into equal hybrid orbitals: sp (linear, 180°), sp² (trigonal planar, 120°), sp³ (tetrahedral, 109.5°), sp³d (trigonal bipyramidal) and sp³d² (octahedral).
• Valency – Valency tells how many bonds one atom can make. For groups 1, 2 and 13 it equals the group digit (1, 2, 3). For groups 14 to 17 the valency with oxygen equals the group digit (4 to 7) and the valency with hydrogen is 8 minus it. Oxide formulae come from matching hooks.
• Vitamins, Nucleic Acids and Hormones: Types, Roles, DNA and RNA – Vitamins are small organic compounds we need in tiny amounts from food; A, D, E, K dissolve in fat and B, C dissolve in water, and a lack of each causes a known disease. Nucleic acids (DNA and RNA) are long chains of nucleotides. Each nucleotide = a pentose sugar + a phosphate + a nitrogen base. DNA has deoxyribose, bases A, G, C, T and two strands in a double helix, where A pairs with T and G with C. RNA has ribose, U instead of T and usually one strand. Hormones are chemical messengers made by glands, and can be steroids, proteins/peptides or amino acid derivatives.
• VSEPR Theory: Shapes of Molecules – VSEPR stands for Valence Shell Electron Pair Repulsion. The electron pairs around a central atom repel each other and move as far apart as possible. The number of pairs sets the basic arrangement: 2 linear, 3 trigonal planar, 4 tetrahedral, 5 trigonal bipyramidal, 6 octahedral. Lone pairs take more room than bond pairs (lp–lp > lp–bp > bp–bp), so they squeeze bond angles and change the shape we see, for example NH₃ (pyramidal, 107°) and H₂O (bent, 104.5°).
• Water Purification: From Dirty Water to Safe Drinking Water – Potable water is water that is safe to drink. It is not pure: it still has some dissolved salts, but only small amounts and no harmful germs. To make it, we choose a good source, let solids settle, filter it, and kill germs with chlorine, ozone or UV light. Where fresh water is short, sea water is desalinated by distillation or reverse osmosis. Waste water from homes and farms is treated before it goes back to rivers. Hard water has dissolved calcium and magnesium salts; it can be softened by boiling or ion exchange.
• Water Quality: Safe Drinking Water, Pollution and Soil – Safe drinking water has no harmful germs, is clear, has no bad taste or smell, has pH about 6.5 to 8.5 and has dissolved chemicals (nitrate, arsenic, lead, salts) below set limits. Water gets polluted by farms (fertiliser, pesticide), factories (metals, chemicals) and towns (sewage). Soil helps: its clay and humus hold positive ions and filter particles and germs, but nitrate slips through and soil can fill up.
• What Is Chemistry? A Short History and How Chemists Work – Chemistry is the science of substances: what they are made of, how they change and how we can make new ones. It grew from ancient crafts (metals, glass, dyes) and alchemy to careful weighing (Lavoisier) and atoms (Dalton). Chemists use observation, experiment and measurement, write in symbols and formulas, and check facts in reliable sources.