JC1 H2 Chemistry vocabulary becomes much easier when students can connect the mole concept, atomic structure, bonding and chemical reactions. This guide explains 200 useful Chemistry words in straightforward language, with examples and a revision route for Singapore junior college students. It follows the key ideas and extensions in the 2027 SEAB H2 Chemistry syllabus 9476.
Do not attempt to learn all 200 terms on the first day of JC1. Your college may teach topics in a different order and some of these concepts belong to later parts of H2 Chemistry. Start with the cluster matching the current lecture, then check the definition against a calculation, molecule, graph, experiment or everyday phenomenon.
- Most useful first: mole, stoichiometry, limiting reagent, electronegativity, bonding, enthalpy, equilibrium and oxidation.
- Explore the full guide: atomic structure, chemical kinetics, acids and bases, electrochemistry, organic reactions and practical science.
- Important: this is an independently authored 200-term study glossary, not an officially ranked SEAB word list or a promise of a particular exam question.
What changed for the 2027 H2 Chemistry A-Level?
The 2027 H2 Chemistry syllabus code is 9476. Older guides using 9729 should not be treated as a substitute for checking the current syllabus. The curriculum emphasises understanding matter, structure and properties, transformation, and applications to environmental sustainability and materials. Scientific practices and practical investigation are central to learning the subject.
Fast navigation to the 200 Chemistry words
- Mole concept and chemical quantities (1–20)
- Atomic structure and electronic arrangement (21–40)
- Bonding, structure and properties (41–60)
- Thermochemistry and thermodynamics (61–80)
- Rates and reaction mechanisms (81–100)
- Chemical equilibrium (101–120)
- Acids, bases and aqueous solutions (121–140)
- Redox and electrochemistry (141–160)
- Organic chemistry (161–180)
- Practical chemistry, materials and society (181–200)
1. Mole concept and chemical quantities: vocabulary 1–20
Start with units. Most calculation mistakes are not arithmetic mistakes: students choose the wrong mole ratio, forget which reagent limits the reaction or confuse grams and moles.
- Atom. The smallest unit of an element retaining its chemical identity.
- Molecule. Two or more atoms joined by covalent bonds.
- Ion. An atom or group with an overall electric charge.
- Element. A substance made of atoms with the same proton number.
- Compound. A substance containing chemically combined atoms of different elements.
- Mixture. Substances combined physically rather than joined into one chemical substance.
- Isotope. Atoms of the same element with different neutron numbers.
- Relative atomic mass. Weighted mean atomic mass compared with one twelfth of carbon-12.
- Relative molecular mass. Sum of relative atomic masses in a molecule.
- Mole. The SI amount unit containing exactly 6.02214076 × 10²³ specified entities.
- Avogadro constant. The number of specified particles per mole.
- Molar mass. Mass of one mole of a specified substance, usually g mol⁻¹.
- Molar volume. Volume occupied by one mole of gas at stated conditions.
- Empirical formula. The simplest whole-number ratio of elements in a compound.
- Molecular formula. The actual number of each element’s atoms in one molecule.
- Stoichiometry. Quantitative relationships between reactants and products in balanced reactions.
- Limiting reagent. The reactant exhausted first according to the reaction proportions.
- Excess reagent. A reactant present in more than the required stoichiometric amount.
- Percentage yield. Actual product yield divided by theoretical yield, multiplied by 100%.
- Atom economy. The fraction of reactant atom mass incorporated into the desired product.
2. Atomic structure and electronic arrangement: vocabulary 21–40
Electronic arrangement is a model of electron behaviour, not a picture of tiny planets. Use its predictions to explain periodic trends and reactivity.
- Proton. A positively charged particle located in the atomic nucleus.
- Neutron. An uncharged nuclear particle contributing to atomic mass.
- Electron. A negatively charged particle occupying atomic orbitals.
- Nucleon number. Total number of protons and neutrons in an atom.
- Proton number. The number of protons identifying an element.
- Electronic configuration. The arrangement of an atom’s electrons among available orbitals.
- Principal quantum number. A number describing an electron shell’s main energy level.
- Orbital. A spatial region associated with an electron probability distribution.
- Subshell. An orbital group of a particular type within a shell.
- s orbital. A spherically symmetric orbital type with one spatial orientation.
- p orbital. An orbital type with directional lobes and three orientations.
- d orbital. A type of orbital associated with transition-element electronic structure.
- Aufbau principle. Electrons occupy lower-energy available orbitals before higher-energy ones.
- Pauli exclusion principle. No two electrons in one atom share all four quantum numbers.
- Hund’s rule. Electrons occupy equivalent orbitals singly before pairing when possible.
- Ionisation energy. Energy needed to remove an electron from a gaseous species.
- First ionisation energy. Energy needed to remove one electron from each mole of gaseous atoms.
- Effective nuclear charge. The net nuclear attraction experienced by an electron after shielding.
- Shielding effect. Reduction of nuclear attraction due to intervening electron density.
- Atomic radius. A measure of atomic size defined through a specified interatomic distance.
3. Bonding, structure and properties: vocabulary 41–60
Relate microscopic bonding to observable properties. Why does one material conduct, dissolve, melt or remain rigid? Its structure matters.
- Ionic bond. Electrostatic attraction between oppositely charged ions.
- Covalent bond. A bond formed through a shared pair of electrons.
- Metallic bond. Attraction between positive metal ions and delocalised electrons.
- Electronegativity. The tendency of an atom to attract bonding electrons.
- Bond polarity. Unequal distribution of bonding electron density.
- Dipole moment. A measure of charge separation in a bond or molecule.
- Lone pair. A pair of valence electrons not participating in a bond.
- Bond pair. A pair of electrons shared between bonded atoms.
- Coordinate bond. A covalent bond whose shared pair initially comes from one atom.
- Sigma bond. A covalent bond formed by head-on orbital overlap.
- Pi bond. A covalent bond formed by sideways overlap of orbitals.
- Bond energy. Energy associated with breaking a specified type of gaseous bond.
- Bond length. The average distance between nuclei of bonded atoms.
- VSEPR theory. A model predicting molecular shape using electron-pair repulsion.
- Molecular geometry. The three-dimensional arrangement of atoms in a molecule.
- Tetrahedral. A geometry with four bond directions approximately 109.5° apart.
- Trigonal planar. A three-bond geometry with approximate 120° angles.
- Hydrogen bonding. Strong directional attraction involving H bonded to N, O or F and a suitable lone pair.
- Instantaneous dipole. Temporary charge separation caused by fluctuating electron distribution.
- Giant covalent structure. An extensive network of atoms joined by covalent bonds.
4. Thermochemistry and thermodynamics: vocabulary 61–80
Always specify the system and sign convention. Heat released by the system is an exothermic change, not a positive reaction enthalpy.
- System. The portion of the universe chosen for a chemical energy analysis.
- Surroundings. Everything outside the thermodynamic system.
- Enthalpy change. Heat exchanged at constant pressure under suitable conditions.
- Exothermic reaction. A reaction releasing heat to surroundings, with negative enthalpy change.
- Endothermic reaction. A reaction absorbing heat from surroundings, with positive enthalpy change.
- Standard conditions. The specified reference conditions for comparing thermodynamic quantities.
- Standard enthalpy of formation. Enthalpy change when one mole of compound forms from elements in standard states.
- Standard enthalpy of combustion. Enthalpy change when one mole of substance undergoes complete combustion.
- Enthalpy of neutralisation. Enthalpy change when an acid-base reaction produces one mole of water.
- Enthalpy of atomisation. Enthalpy change forming one mole of gaseous atoms from an element.
- Hess’s law. Reaction enthalpy depends only on initial and final states.
- Born–Haber cycle. An enthalpy cycle for analysing formation of an ionic solid.
- Lattice enthalpy. Enthalpy associated with forming or separating an ionic lattice, by stated convention.
- Hydration enthalpy. Enthalpy change when gaseous ions become hydrated in water.
- Entropy. A thermodynamic measure related to energy dispersal and accessible microstates.
- Gibbs free energy. The quantity ΔG = ΔH − TΔS for constant-temperature, constant-pressure feasibility.
- Spontaneity. Thermodynamic feasibility under the specified conditions, not necessarily rapid reaction.
- Calorimetry. Measurement of heat changes using temperature and heat capacity.
- Specific heat capacity. Heat needed to raise unit mass by one degree of temperature.
- Heat loss. Energy transferred away from a measured system, affecting experimental accuracy.
5. Rates and reaction mechanisms: vocabulary 81–100
Rates concern how quickly concentration changes. A favourable reaction may still be slow, so thermodynamic feasibility and kinetics answer different questions.
- Reaction rate. Change in reactant or product concentration per unit time, with suitable stoichiometry.
- Collision theory. A model linking reaction rates to effective molecular collisions.
- Activation energy. The minimum energy barrier along a reaction pathway.
- Transition state. A high-energy arrangement near the top of a reaction barrier.
- Catalyst. A substance changing reaction rate without being consumed overall.
- Homogeneous catalyst. A catalyst in the same phase as its reactants.
- Heterogeneous catalyst. A catalyst in a different phase from at least one reactant.
- Rate equation. An experimentally determined relationship between rate and reactant concentrations.
- Rate constant. A proportionality factor in a rate equation at fixed conditions.
- Reaction order. The exponent of concentration in a measured rate law.
- Overall order. The sum of exponents in a rate equation.
- First-order reaction. A reaction first-order in a specified reactant concentration.
- Zero-order reaction. A reaction whose rate is independent of a specified reactant concentration.
- Half-life. The time taken for a relevant reactant quantity to halve.
- Initial rate. The rate measured at the start of a reaction.
- Rate-determining step. The step governing the overall rate in a proposed reaction mechanism.
- Elementary step. A single molecular event represented in a mechanism.
- Intermediate. A species produced then consumed within the mechanism.
- Arrhenius equation. A relationship describing temperature dependence of rate constants.
- Maxwell–Boltzmann distribution. A distribution of molecular energies used to explain temperature effects on rates.
6. Chemical equilibrium: vocabulary 101–120
An equilibrium reaction continues in both directions. The mixture looks unchanged only because forward and backward rates balance.
- Reversible reaction. A reaction capable of proceeding in both forward and reverse directions.
- Dynamic equilibrium. A state in which forward and reverse reaction rates are equal.
- Equilibrium constant. A ratio of equilibrium activities or concentrations according to a stated convention.
- Kc. An equilibrium constant expressed using concentration terms for a defined reaction.
- Equilibrium position. The relative composition of reactants and products at equilibrium.
- Le Chatelier’s principle. A guide predicting how equilibrium responds to a change in conditions.
- Reaction quotient. A reaction ratio calculated for current, not necessarily equilibrium, conditions.
- Partial pressure. A gas’s contribution to the total pressure in a mixture.
- Homogeneous equilibrium. An equilibrium involving species in the same phase.
- Heterogeneous equilibrium. An equilibrium involving species in different phases.
- Common ion effect. Reduced dissociation or solubility following addition of a shared ion.
- Solubility product. An equilibrium expression for dissolution of a sparingly soluble ionic solid.
- Ksp. The equilibrium constant for the dissolution process under its defined convention.
- Saturated solution. A solution at the solubility limit under stated conditions.
- Precipitation. Formation of an insoluble solid from a solution.
- Complex ion. A charged species containing a central atom or ion and bonded ligands.
- Ligand. An atom, ion or molecule donating electron density to a metal centre.
- Coordination number. The number of donor atoms directly bonded to a metal centre.
- Buffer solution. A solution resisting pH change upon small additions of acid or base.
- Haber process. Industrial ammonia synthesis involving nitrogen, hydrogen and a catalyst.
7. Acids, bases and aqueous solutions: vocabulary 121–140
Acid strength is not the same as acid concentration. Strength concerns extent of ionisation; concentration concerns quantity per solution volume.
- Acid. A proton donor in the Brønsted–Lowry model.
- Base. A proton acceptor in the Brønsted–Lowry model.
- Conjugate acid. The species formed after a base accepts a proton.
- Conjugate base. The species formed after an acid loses a proton.
- Strong acid. An acid substantially fully dissociated in water under usual conditions.
- Weak acid. An acid only partly dissociated in aqueous solution.
- Strong base. A base producing essentially complete relevant reaction with water under stated conditions.
- Weak base. A base only partly reacting or protonating in water.
- Acid dissociation constant. An equilibrium constant describing proton donation by a weak acid.
- Ka. The equilibrium constant for a specified weak-acid dissociation.
- pKa. The negative base-ten logarithm of Ka.
- pH. The negative logarithm of hydrogen-ion activity, often approximated from concentration.
- pOH. The negative logarithm of hydroxide-ion activity under the chosen convention.
- Ionic product of water. The equilibrium product for water’s autoionisation, denoted Kw.
- Neutralisation. A chemical reaction between acid and base species.
- Titration. An analytical method using a known solution to determine an unknown amount.
- Equivalence point. The point where titrant and analyte have reacted in stoichiometric proportions.
- Endpoint. The observed indicator or instrumental signal marking completion of a titration.
- Indicator. A substance whose visible or measurable property changes over a useful chemical range.
- Hydrolysis. Reaction with water causing cleavage or chemical transformation of a bond.
8. Redox and electrochemistry: vocabulary 141–160
At the anode, oxidation takes place; at the cathode, reduction takes place. This holds even when electrode signs differ between cells.
- Oxidation. Loss of electrons or increase in oxidation state.
- Reduction. Gain of electrons or decrease in oxidation state.
- Oxidising agent. A species causing oxidation while itself being reduced.
- Reducing agent. A species causing reduction while itself being oxidised.
- Oxidation state. An assigned number used to track electron changes in compounds.
- Half-equation. A balanced representation of one oxidation or reduction process.
- Redox reaction. A chemical change coupling oxidation with reduction.
- Electrode. A conductor through which electrons enter or leave an electrochemical system.
- Anode. The electrode at which oxidation occurs.
- Cathode. The electrode at which reduction occurs.
- Electrolyte. A medium containing mobile ions that can carry electric current.
- Electrolysis. Use of electrical energy to drive a non-spontaneous chemical change.
- Galvanic cell. A cell converting spontaneous chemical energy into electrical energy.
- Electrode potential. A measure of an electrode’s reduction tendency relative to a reference.
- Standard electrode potential. An electrode potential measured under specified standard conditions.
- Cell potential. The potential difference between two electrodes in a cell.
- Salt bridge. An ionic connection maintaining charge balance between half-cells.
- Disproportionation. A reaction in which one element is oxidised and reduced simultaneously.
- Faraday constant. The electric charge per mole of electrons.
- Electroplating. Depositing a metal onto a surface through an electrochemical reaction.
9. Organic chemistry: vocabulary 161–180
Organic vocabulary is a language of electron flow, structures and mechanisms. Identify the functional group before guessing the reaction.
- Functional group. A structural unit largely responsible for characteristic organic reactions.
- Homologous series. A family of compounds with similar functional groups and systematic formula changes.
- Structural isomerism. Compounds sharing a molecular formula but differing in atom connectivity.
- Stereoisomerism. Compounds with the same connectivity but different spatial arrangements.
- Enantiomer. One of two non-superimposable mirror-image molecules.
- Chiral centre. An atom commonly attached to four different groups, allowing chirality.
- Optical activity. Rotation of plane-polarised light by a chiral sample.
- Electrophile. An electron-pair acceptor in an organic reaction.
- Nucleophile. An electron-pair donor in an organic reaction.
- Free radical. A species containing at least one unpaired electron.
- Addition reaction. A reaction joining groups across a multiple bond.
- Substitution reaction. A reaction replacing one atom or group with another.
- Elimination reaction. A reaction removing atoms or groups to form a multiple bond.
- Condensation reaction. A reaction forming a larger molecule while releasing a small molecule.
- Hydrocarbon. A compound containing only carbon and hydrogen.
- Alkane. A saturated acyclic hydrocarbon with only single carbon-carbon bonds.
- Alkene. A hydrocarbon containing at least one carbon-carbon double bond.
- Alcohol. An organic compound containing a hydroxyl group on saturated carbon.
- Carboxylic acid. An organic compound containing the –COOH functional group.
- Ester. An organic compound containing a carboxyl-derived –COO– linkage.
10. Practical chemistry, materials and society: vocabulary 181–200
Experimental evidence needs uncertainty and sound technique. Knowing what an instrument measures is as important as producing a numerical answer.
- Precision. Closeness of repeated measurements to one another.
- Accuracy. Closeness of a measurement to a reference or accepted value.
- Uncertainty. An estimate of the doubt or range associated with a measurement.
- Systematic error. A persistent bias shifting results in a consistent direction.
- Random error. Unpredictable variation causing measured values to scatter.
- Calibration. Checking or adjusting an instrument against known standards.
- Standard solution. A solution with accurately known concentration.
- Volumetric flask. Glassware designed to contain a specified solution volume accurately.
- Pipette. Equipment for delivering a measured liquid volume.
- Burette. Graduated apparatus for controlled delivery of titrant.
- Filtration. Separating a solid from fluid using a porous barrier.
- Recrystallisation. Purifying a solid by dissolution and controlled crystal formation.
- Chromatography. A separation method exploiting different interactions with moving and stationary phases.
- Retention factor. A chromatography distance ratio used under stated experimental conditions.
- Reflux. Heating a mixture while condensing vapour back into the vessel.
- Distillation. Separating substances using differences in volatility and condensation.
- Polymer. A macromolecule made from repeating structural units.
- Monomer. A smaller molecular unit used to build a polymer.
- Biodegradability. The capacity of a substance to be broken down biologically under conditions.
- Green chemistry. Chemical design aimed at reducing hazards, waste and resource use.
Worked example: why the limiting reagent matters
In the reaction 2H₂ + O₂ → 2H₂O, suppose a mixture contains 3 moles of hydrogen and 1 mole of oxygen. The equation needs 2 moles of hydrogen per mole of oxygen. Oxygen is the limiting reagent, because 1 mole of oxygen can react with only 2 moles of hydrogen, leaving 1 mole of hydrogen in excess. The maximum water formed is 2 moles, assuming complete conversion. Four vocabulary terms now work together: stoichiometry, limiting reagent, excess reagent and theoretical yield.
Worked example: strength versus concentration
A dilute strong acid may contain fewer total moles per litre than a concentrated weak acid. ‘Strong’ refers to the extent of ionisation in water, whereas ‘concentrated’ refers to the quantity present per unit solution volume. This difference affects explanations of pH, titration curves and equilibrium. Do not substitute one word for the other.
Worked example: spontaneity does not mean speed
A negative Gibbs free energy change, under the relevant conditions, indicates thermodynamic favourability. It does not tell you how fast the reaction will occur. A high activation-energy barrier can make a feasible reaction extremely slow without a catalyst. In a question, the words feasible, spontaneous, activation energy and rate must be kept separate.
Four-week chemistry vocabulary routine
- Week 1: Learn 40 terms from your current lectures. Write the correct formula or particle-level sketch beside each term.
- Week 2: Add another 40 terms and explain them using a practical situation such as dissolving salt, measuring a titration or comparing reaction rates.
- Week 3: Compare commonly confused pairs: molecule versus compound; bond energy versus enthalpy change; strong versus concentrated; equilibrium position versus equilibrium constant.
- Week 4: Use marked tutorial questions to find errors. Write a corrected sentence, a labelled calculation and one new application for each misunderstood term.
Frequently asked questions
Are all these words taught in JC1?
No. This is a reference spanning the H2 syllabus and useful extensions. Topic sequence varies by junior college, so your school scheme of work is the immediate guide to what to study first.
Is 9476 the same as the old H2 Chemistry 9729 syllabus?
No. 9476 is the revised H2 Chemistry syllabus used for the 2027 examination. Check the 2027 document for precise content and assessment requirements rather than relying only on older notes.
How can I get better at Chemistry explanations?
Practise moving between three descriptions: what is observed, what the particles do, and what an equation or symbol represents. An answer becomes stronger when all three descriptions support one another.
Should I memorise every organic reaction before learning structures?
No. Identify functional groups, bonding and electron movement first. Reaction patterns become easier to understand when you know why a reagent reacts with a particular site.
How can parents check progress?
Ask the student to teach one term using an everyday object or process and then answer one unfamiliar example. The quality of explanation reveals more than how long the student spent copying definitions.
Official syllabus and next reading
Check the official SEAB 2027 H2 Chemistry syllabus (9476) and 2027 A-Level syllabus listing. For guided calculations, visit Limiting reagents, percentage yield and purity. For the JC1 transition, read the O-Level to H2 Chemistry bridging plan. Return to eduKate Singapore for connected study guides.
Teaching guide
Explain a concept using the correct term, an equation and an observable result. Ask the student to identify which word in a question signals the method. If the explanation is memorised but cannot predict an outcome, rebuild from particle-level reasoning.