Top 200 Vocabulary for JC 2 H2 Chemistry Syllabus | A-Level 9476 Organic Reactions, Equilibrium and Practical

JC2 H2 Chemistry vocabulary for Singapore A-Level 9476: this practical glossary gives 200 separately explained words for chemical bonding, thermodynamics, equilibrium, organic chemistry, electrochemistry and experiments. Start with the term that blocks your understanding, then use it to read questions and construct better explanations.

For the 2027 Singapore-Cambridge H2 Chemistry examination, the syllabus code is 9476. Students need both factual understanding and the ability to apply chemical principles to unfamiliar situations. Paper 4 is a practical examination, so measurement and evidence words belong beside reaction mechanisms in the same guide.

Quick information:

• 200 unique definitions, ten numbered chapters and jump links for quick lookup.
• 2027 H2 Chemistry code 9476; do not confuse it with older 9729 study materials.
• Paper 1: 1 hour, 30 marks, 15%. Paper 2: 2 hours, 75 marks, 30%.
• Paper 3: 2 hours, 75 marks, 35%. Paper 4: 2 hours 30 minutes, 50 marks, 20%.
• Includes worked explanations, practical warnings, a four-week plan, FAQs and the original SEAB syllabus.

This is an independent editorial selection of useful scientific language, not an official SEAB list of exactly 200 compulsory words.

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The 200 H2 Chemistry words and what they mean

1. Particles, amounts and chemical equations

1. atom — smallest particle of an element retaining its chemical identity.
2. element — substance consisting of atoms with the same proton number.
3. compound — substance containing different elements chemically bonded in fixed proportions.
4. ion — charged atom or group formed by loss or gain of electrons.
5. isotope — atom of an element with a different neutron number.
6. relative atomic mass — weighted mean isotopic mass compared with one twelfth of carbon-12.
7. relative molecular mass — sum of relative atomic masses within a molecule.
8. amount of substance — number of specified entities expressed in moles.
9. mole — amount containing exactly 6.02214076×10²³ specified entities.
10. Avogadro constant — number of entities in one mole.
11. molar mass — mass per mole of a specified substance.
12. molar volume — volume occupied by one mole under stated conditions.
13. empirical formula — simplest whole-number atom ratio in a compound.
14. molecular formula — actual numbers of each atom type in one molecule.
15. stoichiometric ratio — mole relationship shown by balanced equation coefficients.
16. limiting reagent — reactant used up first that limits product formation.
17. excess reagent — reactant remaining after the limiting reagent is consumed.
18. percentage yield — actual product divided by theoretical product multiplied by 100%.
19. atom economy — percentage of reactant mass ending in the desired product.
20. concentration — amount of dissolved substance per stated volume of solution.

2. Atomic structure and periodic trends

21. proton — positively charged particle in an atomic nucleus.
22. neutron — uncharged particle in an atomic nucleus.
23. electron — negatively charged particle found outside atomic nuclei.
24. nucleon number — total number of protons and neutrons.
25. proton number — number of protons defining an element.
26. electron configuration — arrangement of electrons among orbitals.
27. orbital — region represented by a wavefunction accommodating up to two paired-spin electrons.
28. subshell — group of orbitals with the same angular-momentum type.
29. ionisation energy — energy required to remove an electron from gaseous particles per mole.
30. effective nuclear charge — net attraction felt by an electron after accounting for shielding.
31. shielding effect — reduced outer-electron attraction caused by inner electrons.
32. atomic radius — measure of the size of an atom based on a specified convention.
33. ionic radius — measure of an ion’s size within an ionic structure.
34. electronegativity — ability of an atom in a bond to attract bonding electrons.
35. electron affinity — enthalpy change when a gaseous atom gains an electron.
36. successive ionisation energies — energies required to remove additional electrons one by one.
37. periodic trend — regular pattern in element properties across the Periodic Table.
38. valence shell — outer shell involved in many chemical reactions.
39. transition element — element forming at least one stable ion with an incomplete d subshell.
40. oxidation state — formal electron-accounting number assigned to an atom.

3. Chemical bonding and structures

41. ionic bond — electrostatic attraction between oppositely charged ions.
42. covalent bond — bond formed through a shared pair of electrons.
43. coordinate bond — covalent bond with both shared electrons supplied initially by one atom.
44. metallic bond — attraction between positive metal ions and delocalised electrons.
45. sigma bond — bond with electron density concentrated along the internuclear axis.
46. pi bond — bond formed by lateral orbital overlap.
47. bond polarity — unequal electron sharing across a covalent bond.
48. dipole moment — measure of separation between positive and negative charge.
49. intermolecular force — attraction between neighbouring molecules.
50. London dispersion force — instantaneous-dipole-induced-dipole attraction.
51. permanent dipole attraction — interactions between polar molecules.
52. hydrogen bonding — specific attraction involving H bonded to N, O or F and an available lone pair.
53. lattice energy — enthalpy change when a mole of ionic solid forms from gaseous ions in this syllabus’s convention.
54. bond energy — energy needed to break a mole of specified bonds in the gas phase.
55. molecular geometry — three-dimensional arrangement of atoms in a molecule.
56. VSEPR theory — model predicting shape from repulsion between electron-pair regions.
57. bond angle — angle between two bonds around a central atom.
58. resonance — representation involving more than one contributing electron arrangement.
59. delocalisation — electrons distributed across several bonded atoms.
60. giant covalent structure — extended network of atoms connected by covalent bonds.

4. Energetics and thermodynamic feasibility

61. exothermic reaction — reaction releasing heat to surroundings with negative enthalpy change.
62. endothermic reaction — reaction absorbing heat from surroundings with positive enthalpy change.
63. enthalpy change — heat exchanged at constant pressure under relevant conditions.
64. standard conditions — specified reference conditions for comparing measured thermodynamic quantities.
65. standard enthalpy of formation — enthalpy change forming one mole of compound from elements in standard states.
66. standard enthalpy of combustion — enthalpy change on complete burning of one mole of substance.
67. enthalpy of neutralisation — enthalpy change per mole of water formed from aqueous acid and base.
68. Hess’s Law — total enthalpy change independent of the reaction pathway.
69. Born–Haber cycle — enthalpy cycle linking formation of an ionic solid with component changes.
70. entropy — thermodynamic quantity associated with energy dispersal and number of accessible arrangements.
71. Gibbs free energy — thermodynamic measure combining enthalpy and entropy to assess feasibility.
72. spontaneity — thermodynamic possibility under the specified conditions, not speed.
73. activation energy — minimum energy barrier associated with a reaction pathway.
74. energy profile — diagram of relative energies along a reaction progress coordinate.
75. calorimetry — measurement of heat transferred through observed temperature change.
76. specific heat capacity — energy needed to raise unit mass by one degree.
77. bond dissociation enthalpy — enthalpy needed to break one mole of specified gas-phase bonds.
78. enthalpy of solution — enthalpy change when one mole of solute dissolves.
79. enthalpy of hydration — enthalpy change when a mole of gaseous ions becomes hydrated.
80. enthalpy of atomisation — enthalpy change when one mole of gaseous atoms forms from an element.

5. Rates, reaction pathways and equilibria

81. reaction rate — change in reactant or product concentration per unit time.
82. rate equation — experimental expression linking rate with reactant concentrations.
83. rate constant — proportionality factor in a rate equation at a fixed temperature.
84. order of reaction — exponent of concentration in the experimentally determined rate law.
85. overall order — sum of concentration exponents in a rate equation.
86. half-life — time required for a reacting concentration to fall to half its value.
87. collision theory — explanation of reaction rates through suitably energetic collisions.
88. rate-determining step — slow elementary step controlling an overall reaction rate in a mechanism.
89. catalyst — substance increasing reaction rate while regenerated overall.
90. heterogeneous catalysis — catalysis where catalyst and reacting medium are different phases.
91. homogeneous catalysis — catalysis where catalyst and reactants share a phase.
92. dynamic equilibrium — state with equal forward and reverse rates and stable macroscopic composition.
93. equilibrium constant — expression relating equilibrium activities or concentrations under a chosen convention.
94. Kc — equilibrium constant written using molar concentrations.
95. Kp — equilibrium constant written using partial pressures of gases.
96. reaction quotient — similar expression using current rather than equilibrium amounts.
97. Le Chatelier’s principle — qualitative prediction of a system’s adjustment after disturbance.
98. partial pressure — pressure a gas would exert if it alone occupied the volume at that temperature.
99. equilibrium position — relative proportions of reactants and products at equilibrium.
100. industrial compromise — balancing rate, equilibrium yield, energy, safety and cost in production.

6. Acids, bases, buffers and solubility

101. Brønsted–Lowry acid — species donating a proton to another species.
102. Brønsted–Lowry base — species accepting a proton.
103. conjugate acid–base pair — two species differing by one transferred proton.
104. amphiprotic species — species able to donate or accept a proton.
105. strong acid — acid almost completely dissociated in aqueous solution.
106. weak acid — acid only partly dissociated in aqueous solution.
107. acid dissociation constant — Ka, equilibrium constant describing acid dissociation.
108. pKa — negative base-10 logarithm of Ka.
109. pH — negative base-10 logarithm of hydrogen-ion activity, often approximated by concentration.
110. pOH — negative base-10 logarithm of hydroxide-ion concentration in simple treatments.
111. ionic product of water — Kw, equilibrium product involving aqueous H+ and OH−.
112. buffer solution — solution resisting large pH change on modest addition of acid or base.
113. buffer capacity — amount of added acid or base a buffer can absorb before substantial pH change.
114. titration — controlled addition of known concentration to determine an unknown.
115. equivalence point — point at which reacting amounts match their stoichiometric ratio.
116. indicator endpoint — observable indicator change used to estimate equivalence.
117. hydrolysis — bond-breaking reaction involving water or its acidic or alkaline forms.
118. salt — ionic product formed through combinations of cations and anions.
119. solubility product — Ksp, equilibrium constant for dissolution of a sparingly soluble solid.
120. common ion effect — suppression of ionisation or solubility through adding an already-present ion.

7. Oxidation, electrode potentials and electrolysis

121. oxidation — increase in oxidation number or loss of electrons.
122. reduction — decrease in oxidation number or gain of electrons.
123. oxidising agent — species that accepts electrons and becomes reduced.
124. reducing agent — species that donates electrons and becomes oxidised.
125. oxidation number — formal charge assigned by electron-accounting rules.
126. half-equation — electron-balanced representation of an oxidation or reduction.
127. redox reaction — reaction involving both oxidation and reduction.
128. electrochemical cell — arrangement converting chemical change into electrical energy.
129. standard electrode potential — potential of a half-cell measured relative to the standard hydrogen electrode.
130. standard hydrogen electrode — reference electrode defined to have potential zero volts.
131. cell potential — potential difference between two electrodes.
132. anode — electrode where oxidation occurs.
133. cathode — electrode where reduction occurs.
134. electron flow — movement of electrons through a conductive path.
135. electrolysis — use of electrical energy to drive a nonspontaneous reaction.
136. Faraday constant — electric charge carried by one mole of electrons.
137. standard cell conditions — reference concentration, pressure and temperature conditions specified for potentials.
138. disproportionation — reaction in which one oxidation state is both oxidised and reduced.
139. electrodeposition — formation of a solid on an electrode by reduction of ions.
140. fuel cell — electrochemical device supplied continuously with fuel and oxidant.

8. Organic structures and mechanisms

141. homologous series — family of compounds with the same functional group and systematic formula pattern.
142. functional group — structural feature responsible for characteristic reactions.
143. structural isomer — compound sharing a molecular formula but differing in connectivity.
144. stereoisomer — compound with the same bonding connections but different spatial arrangement.
145. geometric isomer — stereoisomer arising from restricted rotation, often E/Z.
146. optical isomer — stereoisomer differing in three-dimensional chirality.
147. chiral centre — atom at which interchanging two groups gives a different stereoisomer.
148. enantiomer — one of two non-superimposable mirror-image molecules.
149. racemic mixture — equal mixture of two enantiomers.
150. electrophile — electron-pair acceptor attacking an electron-rich site.
151. nucleophile — electron-pair donor attacking an electron-poor site.
152. free radical — species containing an unpaired electron.
153. homolytic fission — bond cleavage giving one electron to each fragment.
154. heterolytic fission — bond cleavage giving both electrons to one fragment.
155. inductive effect — electron-density shift transmitted through sigma bonds.
156. reaction mechanism — sequence of elementary bond-breaking and bond-forming steps.
157. substitution — reaction replacing one atom or group with another.
158. addition — reaction combining species across an unsaturated centre.
159. elimination — reaction removing groups to form a multiple bond.
160. condensation — reaction joining units with loss of a small molecule.

9. Organic functional groups and transformations

161. alkane — saturated hydrocarbon containing only carbon–carbon single bonds.
162. alkene — hydrocarbon containing a carbon–carbon double bond.
163. arene — aromatic hydrocarbon built around a conjugated ring system.
164. benzene — six-carbon aromatic ring with delocalised pi electrons.
165. halogenoalkane — organic compound with a halogen attached to a saturated carbon.
166. alcohol — compound containing an –OH group attached to saturated carbon.
167. phenol — compound with –OH directly attached to an aromatic ring.
168. aldehyde — compound containing a terminal –CHO carbonyl group.
169. ketone — compound containing a carbonyl bonded to two carbon groups.
170. carboxylic acid — compound containing a –COOH group.
171. acyl chloride — reactive carboxylic acid derivative containing –COCl.
172. ester — compound containing –COOR derived from a carboxylic acid.
173. amide — compound containing carbonyl carbon bonded to nitrogen.
174. amine — nitrogen-containing organic base derived structurally from ammonia.
175. nitrile — compound containing a carbon–nitrogen triple bond.
176. amino acid — molecule containing amino and carboxylic acid groups.
177. peptide bond — amide linkage joining amino-acid residues.
178. addition polymerisation — making polymers by adding unsaturated monomers without small-molecule loss.
179. condensation polymerisation — polymer formation with elimination of small molecules.
180. organic reduction — organic transformation typically involving gain of hydrogen or loss of oxygen.

10. Practical skills and scientific communication

181. experimental design — plan linking a testable question to measurements and controls.
182. independent variable — factor deliberately changed by the experimenter.
183. dependent variable — measured response expected to depend on the changed factor.
184. controlled variable — factor held stable to make comparison interpretable.
185. measurement uncertainty — range expressing limits on the measurement’s precision.
186. precision — closeness of repeated measured values to each other.
187. accuracy — closeness of a measurement to the accepted or true value.
188. significant figures — digits indicating precision in a reported value.
189. calibration — checking and adjusting a measuring instrument against a reference.
190. titration curve — graph of measured pH against volume of titrant added.
191. aliquot — accurately measured portion of a larger solution.
192. standard solution — solution with accurately known concentration.
193. primary standard — suitably stable pure substance for preparing accurate standards.
194. qualitative analysis — identification based on observable chemical tests.
195. confirmatory test — test supporting a particular identification over alternatives.
196. precipitate — insoluble solid formed from a solution by a chemical reaction.
197. reflux — heating with a condenser to minimise loss of volatile material.
198. distillation — separation using evaporation and condensation.
199. explain — give reasons in terms of relevant chemistry principles.
200. evaluate — weigh evidence, limitations and alternatives in a conclusion.

Why these Chemistry words matter beyond the textbook

Many Chemistry errors begin as language errors. “Spontaneous” does not mean “fast”: a thermodynamically feasible process may have a high activation-energy barrier. “Strong” is not “concentrated”: strength describes extent of acid dissociation, while concentration describes how much is present in a volume. “Accuracy” and “precision” are not synonyms: a set of titration readings may closely agree and still be systematically wrong.

These differences matter in healthcare, environmental chemistry, batteries and chemical manufacturing. A buffer helps control pH; an electrochemical cell transfers electrons through a circuit; catalyst selection can reduce energy requirements. Chemistry becomes easier when a student can see the same principle in a school diagram and an everyday decision.

Three mini-examples of precise vocabulary in action

Example 1: strong acid versus concentrated acid

A dilute solution of hydrochloric acid can still be a strong acid because almost every dissolved HCl molecule dissociates into ions. A concentrated ethanoic acid solution can remain a weak acid because it only partly ionises. Concentration affects total dissolved amount; acid strength concerns the equilibrium position of dissociation. When a question compares pH, identify both features before calculating.

Example 2: rate versus equilibrium yield

An iron catalyst in the Haber process accelerates the forward and reverse reactions and helps equilibrium be reached more quickly. It does not change the equilibrium constant at a fixed temperature. Pressure and temperature changes can affect the equilibrium composition, but operating conditions must also consider cost, safety and reaction speed. That is an industrial compromise, not a single magic condition.

Example 3: thermodynamic feasibility versus reaction speed

The expression ΔG°=ΔH°−TΔS° combines enthalpy and entropy at the stated temperature. A negative ΔG° favours thermodynamic feasibility under standard conditions; it does not by itself say the reaction finishes quickly. The activation-energy barrier and reaction pathway address kinetics. Use these two sets of terms in different parts of an explanation.

Six confusions that cost explanation marks

How to prepare for Paper 4 without memorising descriptions blindly

In an acid-base titration, write down the measured initial and final burette readings, calculate titre by subtraction, repeat appropriately and determine whether results agree. In a qualitative analysis experiment, describe what is actually observed before proposing the identity of an ion or gas. An interpretation is not an observation. For any heating experiment, distinguish reflux (condensation returns vapour to the vessel) from distillation (condensate is collected).

SEAB provides Qualitative Analysis Notes in Paper 4; these are not a substitute for learning how to observe, record and reason. The practical assessment includes planning, manipulation and measurement, presentation, analysis and evaluation.

A four-week JC2 Chemistry vocabulary and practice plan

1. Week 1 — matter and structure: learn words 1–60 and connect each definition to an equation, electron diagram or bonding model.
2. Week 2 — changes and equilibrium: learn 61–120; practise enthalpy signs, Gibbs free energy, rates, acids and buffer interpretations.
3. Week 3 — reactions and organic mechanisms: learn 121–180; classify electron transfer, reagents, reaction types and functional groups before predicting products.
4. Week 4 — laboratory thinking: learn 181–200; redo practical and structured questions with a column for observations, explanation and limitations.

The order is a revision scaffold, not an official teaching sequence. Students with weak mole calculations or organic foundations should devote more time to those gaps.

Frequently asked questions

Is H2 Chemistry 9476 the same as older syllabus 9729?

No. This guide follows the 2027 examination code 9476. Always use the original syllabus for your sitting year; do not assume paper structures or topic descriptions from older notes apply unchanged.

How much is the practical paper worth?

Paper 4 is 50 marks, two hours and thirty minutes, or 20% of the assessment in 2027.

Does the 2027 syllabus include Gibbs free energy and entropy?

Yes. Its chemical energetics section includes entropy, thermodynamic spontaneity and the relation ΔG°=ΔH°−TΔS°. Pay attention to the precise calculation boundaries stated in the official syllabus.

What is the best way to learn organic reaction mechanisms?

Begin with the functional group, identify electron-rich and electron-poor centres, state reagents and conditions, then explain bonds broken and formed. Copying a mechanism without recognising the attacking species is fragile learning.

Are laboratory terms useful outside school?

Yes. Precision, uncertainty, validity, concentration and contamination decisions are important in testing drinking water, pharmaceuticals and industrial processes.

Can memorising these 200 terms guarantee a distinction?

No. A strong grade also depends on conceptual understanding, accurate calculations, experiment skills, integrated applications and practice with authentic question formats.

Official sources and related eduKate explanations

• SEAB: 2027 A-Level syllabus listing
• SEAB: H2 Chemistry 9476 syllabus PDF (2027)
• JC2 H2 Chemistry: acidic and alkaline hydrolysis of ethanamide
• JC2 H2 Chemistry: benzene, methylbenzene and electrophilic nitration
• JC2 H2 Chemistry: acidic and alkaline manganate(VII) half-equations
• JC2 Mathematics: 200 terms for calculations and data

Reviewed against the SEAB 2027 subject syllabus on 6 November 2026. This educational glossary is independently written and is not an MOE or SEAB publication.