Part II: Chemistry
Free study material · concepts, shortcuts & solved questions
Chemistry questions in this exam bracket rarely go beyond "identify the common name/formula," "state the pH," or "identify the process" — the syllabus here is closer to a kitchen-and-body chemistry quiz than an organic-synthesis exam. Master the tables below and you cover the overwhelming majority of what actually gets asked.
Atomic Structure Basics
Every atom is built from three fundamental particles, whose relative charge and mass every general-science paper tests at some point:
| Particle | Charge | Relative mass | Location |
|---|---|---|---|
| Proton | Positive (+1) | ~1 atomic mass unit | Nucleus |
| Neutron | Neutral (0) | ~1 atomic mass unit (slightly more than a proton) | Nucleus |
| Electron | Negative (−1) | ~1/1836 of a proton's mass (essentially negligible) | Orbiting the nucleus, in shells/energy levels |
Atomic number (Z) = the number of protons in an atom's nucleus — this single number defines which element an atom is (every atom with 6 protons is carbon, no matter how many neutrons it has). Mass number (A) = the total number of protons plus neutrons in the nucleus (electrons are too light to meaningfully add to mass). Why the distinction is tested constantly: a question often gives both numbers and asks you to compute the number of neutrons (A − Z) or electrons (equal to protons, Z, in a neutral atom) — a two-second calculation once the definitions are solid.
Isotopes, in one line: atoms of the same element (same atomic number, same number of protons) that differ in mass number (different number of neutrons) — chemically near-identical (since chemical behaviour depends on electrons, which isotopes share) but physically distinguishable, and sometimes radioactive. Everyday examples worth knowing: Carbon-12 and Carbon-14 (the latter radioactive, used in carbon dating of archaeological/fossil material — a fact that recurs constantly in general-science and even history-adjacent questions); Uranium-235 and Uranium-238 (U-235 is the fissile isotope used in nuclear reactors and weapons, while U-238 is far more abundant but not readily fissile); Hydrogen's three isotopes — protium (no neutron, the ordinary form), deuterium (one neutron, "heavy hydrogen," used in heavy water), and tritium (two neutrons, radioactive, used in some fusion research).
Electron shells/orbits, in one line: electrons occupy defined energy levels (shells), labelled K, L, M, N... from nearest to farthest from the nucleus, with a maximum capacity per shell given by 2n² (K=2, L=8, M=18...) — this is the same "Bohr model" picture credited to Niels Bohr in the Physics part above, and the basis for why elements in the same periodic-table group (same number of outermost/valence electrons) behave similarly.
Memory hook — "PNE, positive-neutral-electron-negative": Proton (positive), Neutron (neutral), Electron (negative) — the three particles in alphabetical-and-charge order together, an easy way to never mix up which is which.
Practice Recall: Distinguish atomic number from mass number, define an isotope in one sentence, and name the two isotopes of carbon and the two isotopes of uranium referenced above, with why each matters (dating vs nuclear fuel).
The Periodic Table's Organising Logic
The modern periodic table (based on Mendeleev's original insight, later corrected using atomic number rather than atomic mass by Henry Moseley) arranges elements in periods (horizontal rows, 7 in total) and groups (vertical columns, 18 in total). Elements in the same group share similar chemical properties because they have the same number of electrons in their outermost shell (valence electrons) — this single idea explains almost every periodic-table question this exam asks, so it's worth internalising over memorising individual element facts.
| Group name | Position | 2–3 examples | Key property |
|---|---|---|---|
| Alkali metals | Group 1 | Lithium (Li), Sodium (Na), Potassium (K) | Soft, highly reactive metals; react vigorously with water to release hydrogen gas; one valence electron, easily lost |
| Alkaline earth metals | Group 2 | Beryllium (Be), Magnesium (Mg), Calcium (Ca) | Reactive but less so than Group 1; two valence electrons |
| Halogens | Group 17 | Fluorine (F), Chlorine (Cl), Iodine (I) | Highly reactive non-metals; seven valence electrons, need just one more for a stable outer shell, making them strong oxidisers |
| Noble gases | Group 18 | Helium (He), Neon (Ne), Argon (Ar) | Chemically inert (complete outer shell already, nothing to gain or lose); used in lighting (neon signs) and as inert atmospheres (argon in welding) |
Memory hook: Reactivity in Group 1 increases going down the group (potassium reacts more violently with water than sodium, sodium more than lithium) — but reactivity in Group 17 decreases going down (fluorine is the most reactive element of all; iodine is comparatively mild). One group gets more reactive down the table, the other less — a favourite "which trend is correct" trap.
Chemical Bonding Basics
Why atoms bond at all, in one line: atoms bond with each other to achieve a stable, complete outer electron shell (usually eight electrons — the "octet rule"), either by fully transferring electrons or by sharing them — the two routes give the two fundamental bond types below.
- Ionic (electrovalent) bond: formed when one atom transfers one or more electrons to another, creating oppositely charged ions that then attract each other electrostatically — typically between a metal (which loses electrons easily) and a non-metal (which gains them easily). Example: sodium chloride (common salt) — sodium (Na) loses one electron to become Na⁺, chlorine (Cl) gains that electron to become Cl⁻, and the two ions bond through electrostatic attraction. Typical properties of ionic compounds: high melting/boiling points, conduct electricity when molten or dissolved in water (because ions become free to move), and are often soluble in water.
- Covalent bond: formed when two atoms share one or more pairs of electrons, typically between two non-metals. Example: water (H₂O) — each hydrogen atom shares one electron with oxygen, and oxygen shares one electron with each hydrogen, so both atoms achieve a fuller outer shell without either fully gaining or losing an electron. Typical properties of covalent compounds: generally lower melting/boiling points than ionic compounds, and generally poor conductors of electricity (no free ions).
Memory hook — "ionic transfers, covalent shares": picture ionic bonding as one atom giving away a possession outright (a transfer, like a gift — creates two separately charged owners), and covalent bonding as two atoms sharing custody of the same possession (neither one gives it up entirely).
Practice Recall: Distinguish an ionic bond from a covalent bond with one example each, and state one typical property (conductivity, melting point) that follows from each bond type.
Common Name, Chemical Name, Formula
This is the single highest-yield table in the whole Chemistry section — nearly every general-science paper asks at least one "common name vs chemical name" matching question.
| Common name | Chemical name | Formula |
|---|---|---|
| Baking soda | Sodium bicarbonate | NaHCO₃ |
| Washing soda | Sodium carbonate (decahydrate) | Na₂CO₃·10H₂O |
| Caustic soda | Sodium hydroxide | NaOH |
| Caustic potash | Potassium hydroxide | KOH |
| Table salt / common salt | Sodium chloride | NaCl |
| Chalk / limestone | Calcium carbonate | CaCO₃ |
| Quick lime | Calcium oxide | CaO |
| Slaked lime | Calcium hydroxide | Ca(OH)₂ |
| Gypsum | Calcium sulphate dihydrate | CaSO₄·2H₂O |
| Plaster of Paris | Calcium sulphate hemihydrate | CaSO₄·½H₂O |
| Blue vitriol | Copper sulphate | CuSO₄ |
| Green vitriol | Ferrous sulphate | FeSO₄ |
| Marsh gas | Methane | CH₄ |
| Laughing gas | Nitrous oxide | N₂O |
| Dry ice | Solid carbon dioxide | CO₂ |
| Vinegar | Acetic acid (dilute) | CH₃COOH |
| Alcohol (spirit) | Ethanol | C₂H₅OH |
| Epsom salt | Magnesium sulphate | MgSO₄ |
Practice Recall: Cover the right two columns and try to give the chemical name and formula for baking soda, washing soda, quick lime, and laughing gas — these four alone cover a disproportionate share of repeat questions.
The pH Scale, Acids, Bases and Salts
The pH scale runs from 0 to 14 and measures how acidic or basic (alkaline) a solution is, based on the concentration of hydrogen ions (H⁺) present — pH = −log₁₀[H⁺]. A pH of 7 is neutral (pure water); values below 7 are acidic (the lower, the more acidic); values above 7 are basic/alkaline (the higher, the more basic). Each whole-number step on the scale represents a tenfold change in H⁺ concentration — a solution of pH 4 is ten times more acidic than one of pH 5, not merely "a bit more."
| Substance | Approximate pH |
|---|---|
| Battery acid (gastric/stomach acid, similarly acidic) | ~1–2 |
| Lemon juice | ~2 |
| Vinegar | ~3 |
| Human blood | ~7.4 (slightly basic) |
| Pure water | 7 (neutral) |
| Sea water | ~8 |
| Baking soda solution | ~9 |
| Milk of magnesia | ~10 |
| Ammonia solution | ~11–12 |
| Bleach / caustic soda solution | ~13 |
Why it matters medically and agriculturally: human blood pH must stay tightly within about 7.35–7.45 — even small deviations outside this narrow band are medically dangerous, which is why blood pH (not stomach pH) is the "human body pH" figure most exams actually want. Soil pH similarly determines which crops grow well; most crops prefer near-neutral soil, and lime is added to acidic soil specifically to raise its pH.
Acids, bases and salts, expanded with everyday examples:
- Acid, in one line: a substance that releases H⁺ ions in water, tastes sour, turns blue litmus red, and reacts with a base to form a salt and water (neutralisation). Everyday acids: citric acid (in lemons, oranges), acetic acid (vinegar), lactic acid (produced in sour milk/curd and in tired muscles during intense exercise), tartaric acid (tamarind, grapes), oxalic acid (in tomatoes and spinach in small amounts), hydrochloric acid (naturally present in the stomach, aiding digestion).
- Base/alkali, in one line: a substance that releases OH⁻ ions in water, tastes bitter, feels soapy to the touch, turns red litmus blue, and neutralises acids. Everyday bases: soap (mildly alkaline), baking soda, milk of magnesia (used as an antacid), ammonia solution (household cleaners), lime water.
- Salt, in one line: the product formed when an acid reacts with a base (or a metal), typically an ionic compound made of a metal/ammonium cation and a non-metal/acid-derived anion — common salt (NaCl) is only one example among a whole family; others include potassium nitrate, calcium sulphate, and ammonium chloride.
- Indicators, in one line: substances that change colour depending on whether a solution is acidic or basic — litmus (red in acid, blue in base) is the classic laboratory example; turmeric (turns reddish-brown in a base, a fact used in the classic "why does turmeric stain on soap-washed cloth turn red" household observation) and red cabbage extract are natural indicators sometimes cited in general-science questions.
Metals vs Non-Metals
| Property | Metals | Non-metals |
|---|---|---|
| Physical state | Solid at room temperature (except mercury, liquid) | Solid, liquid, or gas (varies — e.g., bromine is liquid) |
| Conductivity (heat & electricity) | Good conductors | Poor conductors (except graphite, a non-metal that conducts) |
| Malleability/ductility | Malleable (can be hammered into sheets) and ductile (drawn into wires) | Brittle, if solid |
| Lustre | Shiny | Generally dull |
| Reaction with oxygen | Forms basic oxides | Forms acidic oxides |
| Electron behaviour | Tend to lose electrons (form cations) | Tend to gain electrons (form anions) |
| Examples | Iron, copper, aluminium, gold, sodium | Oxygen, nitrogen, sulphur, chlorine, carbon |
Metalloids (like silicon, boron, arsenic) sit at the border and share properties of both — silicon's intermediate conductivity is exactly why it underlies semiconductor electronics.
Important Ores and Their Metals
An ore is a naturally occurring mineral deposit from which a metal can be economically extracted. Matching an ore's name to the metal it yields is a recurring "identify the correct pair" question:
| Ore | Metal extracted | Chemical form |
|---|---|---|
| Bauxite | Aluminium | Hydrated aluminium oxide |
| Haematite | Iron | Iron(III) oxide, Fe₂O₃ |
| Magnetite | Iron | Iron(II,III) oxide, Fe₃O₄ |
| Cassiterite (tin stone) | Tin | Tin(IV) oxide, SnO₂ |
| Galena | Lead | Lead sulphide, PbS |
| Copper pyrites (chalcopyrite) | Copper | Copper-iron sulphide, CuFeS₂ |
| Pitchblende (uraninite) | Uranium | Uranium oxide |
| Rock salt / Sylvite | Sodium / Potassium | Sodium chloride / Potassium chloride |
| Dolomite | Magnesium (and calcium) | Calcium magnesium carbonate |
Why this table matters beyond memorisation: ore-to-metal matching questions test whether you understand that a metal's chemical form in nature is usually an oxide, sulphide, or carbonate rather than the free, pure metal — because most metals are reactive enough to combine readily with oxygen, sulphur, or carbonate ions in the environment, and only a handful of very unreactive metals (like gold and, to a lesser extent, silver) are commonly found in their pure, "native" state.
Alloys — Composition and Use
An alloy is a homogeneous mixture of two or more elements, at least one of which is a metal, engineered to have properties (strength, corrosion resistance, workability) superior to its individual component metals — a favourite "composition and use" matching table across nearly every general-science paper:
| Alloy | Composition | Key property / use |
|---|---|---|
| Brass | Copper + Zinc | Corrosion-resistant, good acoustic properties; used for musical instruments, fittings, decorative items |
| Bronze | Copper + Tin | Harder and more corrosion-resistant than pure copper; used historically for tools/weapons (the "Bronze Age"), and today for statues, medals, bearings |
| Stainless steel | Iron + Chromium + Nickel (+ small carbon) | Highly resistant to rusting/corrosion (the chromium forms a protective oxide layer); used for cutlery, surgical instruments, kitchenware |
| Duralumin | Aluminium + Copper + Magnesium + Manganese | Very high strength-to-weight ratio; used in aircraft bodies and other aerospace applications |
| Solder | Tin + Lead (or lead-free alternatives today) | Low melting point; used to join electrical/electronic components |
| Steel (carbon/mild steel) | Iron + Carbon (small %) | Much harder and stronger than pure iron; the basic structural-engineering metal |
| Amalgam | Mercury + another metal (e.g., silver) | Historically used in dental fillings; mercury's ability to alloy with almost any metal at room temperature is itself a distinguishing fact about mercury |
| Gunmetal | Copper + Tin + Zinc | Resistant to seawater corrosion; used historically for cannons/gun parts and today for marine fittings |
| Nichrome | Nickel + Chromium | High melting point and high electrical resistance; used for heating elements (electric heaters, toasters) |
| Alnico | Aluminium + Nickel + Cobalt (+ Iron) | Produces very strong permanent magnets |
Memory hook — "Brass has Zinc, Bronze has Tin — both start differently, both add differently": the two most-confused alloys, brass and bronze, both pair with copper — remembering that brass and zinc don't share a letter forces you to actively recall rather than guess, while bronze and tin both contain "n," a small but genuine anchor.
Practice Recall: Name three ores and the metal each yields, then name three alloys with their exact composition and one real-world use each — without checking, say which alloy is chosen specifically for its resistance to rusting.
Important Chemical Reactions and Processes
Rusting of iron is an oxidation reaction: iron reacts with oxygen and moisture in the air to form hydrated iron(III) oxide (Fe₂O₃·xH₂O), the reddish-brown flaky substance we call rust. It requires both oxygen and water to occur — iron doesn't rust in dry air or in water with no dissolved oxygen, which is why galvanising (a zinc coating) and painting both work as rust prevention, by simply keeping air/moisture off the metal surface.
Photosynthesis, the reaction underlying nearly all life on Earth, in its balanced chemical form:
6CO₂ + 6H₂O + (sunlight, chlorophyll) → C₆H₁₂O₆ + 6O₂
Carbon dioxide and water, in the presence of sunlight captured by chlorophyll, become glucose (a sugar the plant uses for energy/growth) and oxygen (released as a by-product) — the reverse, essentially, of respiration.
Saponification is the reaction that makes soap: a fat or oil (an ester) reacts with a strong alkali (typically sodium hydroxide) to produce soap (a sodium salt of a fatty acid) and glycerol as a by-product. This is why traditional soap-making uses caustic soda and animal fat/vegetable oil as its two core ingredients.
Combustion is any reaction of a substance with oxygen that releases heat and light — complete combustion of a hydrocarbon (like cooking gas, LPG) produces carbon dioxide and water; incomplete combustion (insufficient oxygen) produces carbon monoxide (a poisonous gas) and soot (carbon), which is why poorly ventilated gas heaters pose a carbon-monoxide risk.
Neutralisation is the reaction of an acid with a base to produce a salt and water — the chemistry behind why an antacid tablet (a mild base) relieves acidity (excess stomach acid).
Oxidation and reduction (redox), in one line each: oxidation is the loss of electrons by a substance (often, though not always, accompanied by a gain of oxygen); reduction is the gain of electrons (often accompanied by a loss of oxygen or gain of hydrogen) — the two always happen together in a single "redox" reaction, since an electron lost by one substance must be gained by another. Rusting, combustion, and respiration are all, at their core, oxidation reactions.
Electrolysis, one line: the process of using an electric current to drive a non-spontaneous chemical reaction, typically decomposing a compound into its elements — used industrially to extract reactive metals like aluminium and sodium from their molten ores/salts, and in water electrolysis (splitting water into hydrogen and oxygen gas), increasingly relevant to "green hydrogen" discussions in current affairs.
Organic Chemistry Basics
Organic chemistry is, at its simplest, the chemistry of carbon compounds — carbon's unique ability to form long chains and rings by bonding to itself (a property called catenation) is why it alone supports millions of known compounds, far more than any other element, and why organic chemistry gets treated as its own vast sub-field rather than folded into general chemistry.
- Hydrocarbons, in one line: compounds made of only carbon and hydrogen — the simplest class of organic compounds, and the basis of fuels like LPG, petrol, and natural gas. Saturated hydrocarbons (alkanes) contain only single bonds between carbon atoms (methane CH₄, ethane C₂H₆, propane C₃H₈); unsaturated hydrocarbons contain at least one double bond (alkenes, like ethylene C₂H₄) or triple bond (alkynes, like acetylene C₂H₂) — unsaturated hydrocarbons are generally more chemically reactive than saturated ones, since the double/triple bond is a site where other atoms can readily attach.
- Functional groups, one-liners: a functional group is a specific small cluster of atoms within a larger molecule that determines how that molecule chemically behaves, regardless of the rest of the molecule's structure. Alcohol (–OH): found in ethanol, the active ingredient of alcoholic drinks and a common industrial solvent. Carboxylic acid (–COOH): found in acetic acid (vinegar) and citric acid, gives organic acids their sour taste and acidic behaviour. Aldehyde (–CHO) and Ketone (C=O within a chain): reactive carbonyl-containing groups common in many natural and industrial organic compounds. Amine (–NH₂): nitrogen-containing group found in amino acids, the building blocks of proteins.
- Common polymers — plastics, one-liners: a polymer is a large molecule built from many repeating smaller units (monomers) joined together. Polythene (polyethylene): made from repeating ethylene units; the most common plastic, used in bags and containers. PVC (polyvinyl chloride): rigid, durable; used in pipes, cables, flooring. Polystyrene: lightweight, used in disposable cups/packaging (as "thermocol"/expanded polystyrene). PET (polyethylene terephthalate): used in beverage bottles, recognisable by the recycling code "1." Nylon: a synthetic polymer used in fibres/fabrics and ropes, notable as one of the first fully synthetic fibres (developed in the 1930s). Teflon (PTFE): extremely heat- and chemical-resistant; used in non-stick cookware coatings. Why plastics are tested in general science: most common plastics are non-biodegradable, persisting in the environment for centuries — the chemical basis of plastic-pollution discussions that increasingly cross over from pure science into environment/current-affairs sections of these same exams.
Memory hook — "Same C and H, different bonds": alkanes, alkenes, and alkynes all contain only carbon and hydrogen — what changes, in that exact order, is single → double → triple bonding between carbons, and "-ane, -ene, -yne" endings mirror that exact single/double/triple progression alphabetically (a comes before e comes before y).
Practice Recall: Distinguish a saturated from an unsaturated hydrocarbon, name one functional group and one everyday compound that contains it, and name three common plastics with one use each — without checking, say which of those plastics is associated with the "1" recycling code on beverage bottles.
Chemistry in Everyday Life
- Why onions make you cry: cutting an onion ruptures its cells and releases an enzyme that reacts with sulphur compounds to form a volatile gas (syn-propanethial-S-oxide); on reaching the eye, it dissolves in tears to form a mild sulphuric-acid-like irritant, triggering more tears as a defence.
- Why ice floats on water: water is one of the very few substances that is less dense as a solid than as a liquid, because ice's crystal structure holds water molecules in a hexagonal lattice with more open space than liquid water — this anomaly is also why ponds freeze from the top down, letting aquatic life survive winter beneath the ice.
- Why a pressure cooker cooks faster: trapping steam raises internal pressure, which raises the boiling point of water above 100°C — food cooks faster at a higher temperature.
- Why silver/copper vessels were traditionally used for water storage: both metals have mild antimicrobial properties (oligodynamic effect — even trace ion concentrations can kill many bacteria).
- Why a fire is put out by water but not by an electrical or oil fire: water conducts electricity (making it dangerous on live wires) and is denser than oil (so it sinks under burning oil, spreading rather than smothering it, and can even cause dangerous splattering) — CO₂ or dry-chemical extinguishers are used for electrical and oil fires instead, since they smother the fire by cutting off oxygen without either risk.
- Why the sky is blue and sunsets are red: both are the same phenomenon — Rayleigh scattering — by which shorter wavelengths (blue) scatter more strongly than longer wavelengths (red) off molecules in the atmosphere; at sunset, sunlight travels through much more atmosphere, scattering out almost all the blue and leaving mostly red/orange to reach your eye directly.
- Why soap actually cleans (and why water alone often doesn't): a soap molecule has a "double personality" — one end (a long hydrocarbon tail) is hydrophobic (water-repelling, but attracted to oils/grease), and the other end (an ionic head) is hydrophilic (water-attracting). In water, soap molecules cluster into tiny spherical structures called micelles, with their oil-loving tails pointing inward (trapping grease/dirt) and their water-loving heads pointing outward — this lets soap physically surround and lift away grease that plain water, with its high surface tension and no affinity for oil, cannot dislodge on its own.
- Food preservation methods, one line each: refrigeration/freezing slows microbial growth and enzyme activity by lowering temperature; drying/dehydration removes the water microbes need to survive (traditional sun-drying of grains, fruit); salting/curing draws water out of food (and out of microbial cells) via osmosis, creating an environment too dry/saline for most microbes; pickling uses a high-acid (vinegar) or high-salt brine, both hostile to microbial growth; canning seals food after heating, destroying microbes and then excluding new air/microbes entirely; pasteurisation (see Pasteur, below) uses controlled heating to kill most disease-causing microbes in milk and other liquids without fully cooking the product.
- Common gases — everyday uses and hazards: Oxygen (O₂) supports combustion and respiration; Carbon dioxide (CO₂) is used in fire extinguishers (it's denser than air and doesn't support combustion) and carbonated drinks, but is also the principal greenhouse gas driving climate change; Nitrogen (N₂) makes up about 78% of the atmosphere and is chemically unreactive, used to create an inert packaging atmosphere (chip packets) to prevent oxidation/spoilage; Carbon monoxide (CO) is a colourless, odourless, highly toxic gas produced by incomplete combustion, dangerous because it binds haemoglobin far more strongly than oxygen does, effectively suffocating a person from the inside without any obvious warning smell; Chlorine (Cl₂) is used to disinfect drinking water and swimming pools (it kills bacteria) but is itself toxic in concentrated form; LPG (liquefied petroleum gas), mainly propane and butane, is the household cooking-gas mixture, deliberately given a foul "rotten egg" smell (via an additive, ethyl mercaptan) precisely because pure LPG is odourless and a leak would otherwise go undetected.
Practice Recall: Explain, chemically, why soap can remove grease that water alone cannot, then name three food-preservation methods with the mechanism each relies on, and finally name one gas that is denser than air and used specifically for that reason in fire extinguishers.