Glossary
Every technical word in the course and what it means in plain language, with flashcards for practice.
Every technical term in the course, with a plain-language meaning. On every page these words are marked with a dotted underline: hover or tap one to see its meaning. The “Used in” links point to the sections where each term appears.
Ch 1 · Some Basic Concepts of ChemistryCh 2 · Structure of AtomCh 3 · Classification of Elements and Periodicity in PropertiesCh 4 · Chemical Bonding and Molecular StructureCh 5 · ThermodynamicsCh 6 · EquilibriumCh 7 · Redox ReactionsCh 8 · Organic Chemistry – Some Basic Principles and TechniquesCh 9 · Hydrocarbons
Ch 1 · Some Basic Concepts of Chemistry
| Term | Meaning | Used in |
|---|---|---|
| accuracy | How close a measured value (or the mean of several) is to the true value. | 1.1 5.3 |
| average atomic mass | The abundance-weighted mean of the masses of an element's natural isotopes, $\bar{A}=\sum f_{i}m_{i}$; the value listed in the periodic table. | 1.2 2.1 |
| Avogadro's law | Equal volumes of all gases at the same temperature and pressure contain equal numbers of molecules. | 1.2 |
| dimensional analysis | Converting units by multiplying by unit factors such as $\dfrac{1000\ \mathrm{g}}{1\ \mathrm{kg}}$, each equal to one, so that unwanted units cancel. | 1.1 |
| empirical formula | The formula showing the simplest whole-number ratio of the atoms of each element in a compound, e.g. $\mathrm{CH_{2}O}$ for glucose. | 1.3 4.4 8.4 |
| excess reagent | A reactant present in more than the amount needed to react with the limiting reagent; some of it is left over. | 1.3 |
| formula mass | The sum of the atomic masses of the atoms in one formula unit of an ionic compound, e.g. $58.44\ \mathrm{u}$ for NaCl. | 1.2 1.3 |
| Gay-Lussac's law of gaseous volumes | Gases react, and gaseous products form, in volumes (at the same temperature and pressure) that are in a simple whole-number ratio. | 1.2 |
| heterogeneous mixture | A mixture whose composition is not uniform; different parts can often be seen, as in sand stirred into water. | 1.1 |
| homogeneous mixture | A mixture whose composition is uniform throughout, such as air, brass or a sugar solution. | 1.1 1.3 |
| law of conservation of mass | In a chemical reaction in a closed system the total mass of the products equals the total mass of the reactants; matter is neither created nor destroyed. | 1.2 |
| law of definite proportions | A given compound always contains the same elements combined in the same proportion by mass, whatever its source. | 1.2 |
| law of multiple proportions | When two elements form more than one compound, the masses of one that combine with a fixed mass of the other are in a ratio of small whole numbers. | 1.2 |
| limiting reagent | The reactant that is used up first and so fixes the amount of product; it has the smallest value of moles ÷ coefficient. | 1.3 |
| molality | Moles of solute per kilogram of solvent, in $\mathrm{mol\,kg^{-1}}$; independent of temperature. | 1.3 |
| molar mass | The mass of one mole of a substance, in $\mathrm{g\,mol^{-1}}$; numerically equal to its atomic, molecular or formula mass in u. | 1.1 1.2 1.3 2.1 |
| molarity | Moles of solute per litre of solution, $M=n_{2}/V$ in $\mathrm{mol\,L^{-1}}$; changes slightly with temperature. | 1.3 7.2 |
| mole | The SI unit of amount of substance: exactly $6.022\,140\,76\times10^{23}$ specified entities (atoms, molecules, ions…). | 1.1 1.2 1.3 2.1 |
| mole fraction | The moles of one component divided by the total moles of all components, e.g. $x_{2}=\dfrac{n_{2}}{n_{1}+n_{2}}$; the mole fractions add up to 1. | 1.2 1.3 6.1 6.2 |
| molecular formula | The formula giving the actual number of atoms of each element in one molecule; a whole-number multiple of the empirical formula. | 1.3 8.1 8.2 8.4 |
| percentage composition | The mass percentage of each element in a compound; the percentages add up to 100%. | 1.3 8.4 |
| percentage yield | The actual yield of a product as a percentage of the theoretical yield calculated from the limiting reagent. | 1.3 |
| precision | How close repeated measurements of the same quantity are to one another. | 1.1 1.2 2.3 |
| pure substance | Matter with a fixed composition and the same properties throughout: an element or a compound. | 1.1 4.4 5.3 |
| SI base unit | One of the seven units (m, kg, s, A, K, mol, cd) from which every other SI unit is built. | 1.1 |
| significant figures | The digits of a measured value that are known with certainty, plus the first estimated digit. | 1.1 1.2 1.3 2.1 |
| stoichiometry | The calculation of the amounts of reactants and products in a reaction, using the mole ratios given by the balanced equation. | 1.3 7.2 |
| unified atomic mass unit | One-twelfth of the mass of a carbon-12 atom: $1\ \mathrm{u}=1.66054\times10^{-24}\ \mathrm{g}$. | 1.2 2.1 |
Ch 2 · Structure of Atom
| Term | Meaning | Used in |
|---|---|---|
| (n + l) rule | The subshell with the lower value of $n+l$ fills first; for equal values, the one with the lower $n$ fills first. | |
| absorption spectrum | The spectrum of white light after it has passed through a sample: dark lines at the wavelengths the sample absorbs. | 2.2 |
| angular node | A plane or cone through the nucleus on which the probability of finding the electron is zero; an orbital has $l$ of them. | 2.3 |
| anode rays | Also called canal rays: positive ions of the gas in a discharge tube, moving towards the cathode; their charge-to-mass ratio depends on the gas. | 2.1 |
| atomic nucleus | The tiny (radius about $10^{-15}\ \mathrm{m}$), dense, positively charged centre of an atom, containing its protons and neutrons and almost all its mass. | |
| atomic number | The number of protons in the nucleus, $Z$; it identifies the element and equals the number of electrons in the neutral atom. | 2.1 2.3 2.4 3.1 |
| atomic orbital | A one-electron wave function in an atom, pictured as the region where the electron is most likely to be found; it holds at most two electrons. | 2.3 4.3 4.4 |
| Aufbau principle | In the ground state, electrons occupy orbitals in order of increasing energy, lowest first. | 2.4 3.2 4.4 |
| azimuthal quantum number | $l=0,1,\ldots,n-1$ (s, p, d, f…); fixes the subshell and the shape of an orbital. Also called the orbital angular momentum quantum number. | 2.3 |
| black body | An ideal body that absorbs all radiation falling on it and, when hot, emits radiation whose spread of wavelengths depends only on its temperature. | 2.2 |
| cathode rays | Streams of electrons travelling from the cathode towards the anode in a discharge tube at very low pressure and high voltage. | 2.1 |
| charge-to-mass ratio | The charge of a particle divided by its mass, $e/m$; Thomson measured $e/m_{e}=1.76\times10^{11}\ \mathrm{C\,kg^{-1}}$ for the electron. | 2.1 |
| de Broglie wavelength | The wavelength of the wave associated with a moving particle, $\lambda=h/mv$; noticeable only for particles of very small mass such as electrons. | 2.3 |
| degenerate orbitals | Orbitals that have the same energy, such as the three p orbitals of one subshell. | 2.3 2.4 4.4 |
| electronic configuration | The distribution of an atom's or ion's electrons among its subshells, written e.g. $1s^{2}2s^{2}2p^{3}$ for nitrogen. | 2.3 2.4 3.1 3.2 |
| emission spectrum | The spectrum of radiation given out by excited atoms or molecules; for gaseous atoms, a set of bright lines. | 2.2 |
| exchange energy | The stabilisation that arises when electrons with parallel spins in degenerate orbitals can exchange positions; it is greatest for half-filled and filled subshells. | 2.4 5.1 |
| ground state | The lowest-energy state of an atom ($n=1$ for hydrogen); states of higher energy are excited states. | 2.2 2.4 3.2 3.3 |
| Hund's rule | Electrons enter degenerate orbitals singly, with parallel spins, before any pairing occurs. | 2.4 4.4 |
| isobar | Atoms of different elements with the same mass number, e.g. $\mathrm{{}^{14}_{6}C}$ and $\mathrm{{}^{14}_{7}N}$. | 2.1 |
| isotope | Atoms of the same element (same atomic number) with different mass numbers, i.e. different numbers of neutrons, e.g. $\mathrm{{}^{35}Cl}$ and $\mathrm{{}^{37}Cl}$. | 1.2 2.1 3.1 6.1 |
| magnetic quantum number | $m_{l}=-l,\ldots,0,\ldots,+l$; fixes the orientation of an orbital in space. A subshell has $2l+1$ orbitals. | 2.3 |
| mass number | The total number of protons and neutrons in a nucleus, $A$; the number of neutrons is $A-Z$. | 2.1 3.1 |
| noble-gas core | The configuration of the preceding noble gas, written in square brackets as shorthand, e.g. $\mathrm{Na}=[\mathrm{Ne}]\,3s^{1}$. | 2.4 3.2 |
| nucleon | A particle in the nucleus: a proton or a neutron. The number of nucleons is the mass number $A$. | 2.1 |
| orbital diagram | A picture of an electronic configuration in which each orbital is a box and each electron an arrow (↑ or ↓) showing its spin. | 2.4 4.4 |
| Pauli exclusion principle | No two electrons in an atom can have the same four quantum numbers; an orbital holds at most two electrons, with opposite spins. | 2.4 |
| photoelectric effect | The ejection of electrons from a metal surface by light of frequency above a threshold value; explained by Einstein using photons. | 2.2 |
| photon | A quantum (packet) of electromagnetic radiation, with energy $E=h\nu=hc/\lambda$. | 2.2 2.3 3.3 4.1 |
| principal quantum number | $n=1,2,3,\ldots$; fixes the shell, the size of an orbital and (in hydrogen) its energy. | 2.2 2.3 3.2 3.3 |
| radial node | A spherical surface around the nucleus on which the probability of finding the electron is zero; an orbital has $n-l-1$ of them. | 2.3 |
| spin quantum number | $m_{s}=+\tfrac12$ or $-\tfrac12$; describes the spin of an electron (shown as ↑ or ↓). | 2.3 |
| stationary state | In Bohr's model, an allowed orbit of fixed radius and energy in which the electron does not radiate. | 2.2 |
| threshold frequency | The minimum frequency $\nu_{0}$ of light that can eject electrons from a given metal; $h\nu_{0}$ equals the work function. | 2.2 |
| uncertainty principle | The position and momentum of a microscopic particle cannot both be known exactly: $\Delta x\,\Delta p_{x}\ge h/4\pi$. | 2.2 2.3 2.4 |
| unpaired electron | An electron alone in an orbital; species with unpaired electrons are paramagnetic (attracted into a magnetic field). | 2.4 3.2 4.1 4.3 |
| wave function | The solution $\psi$ of the Schrödinger equation for an electron; $\psi^{2}$ gives the probability density of finding the electron at a point. | 2.3 4.4 |
| wavenumber | The number of wavelengths per unit length, $\bar{\nu}=1/\lambda$, usually in $\mathrm{cm^{-1}}$. | 2.2 |
| work function | The minimum energy $W_{0}=h\nu_{0}$ needed to remove an electron from the surface of a metal. | 2.2 |
Ch 3 · Classification of Elements and Periodicity in Properties
| Term | Meaning | Used in |
|---|---|---|
| actinoids | The 14 elements thorium ($Z=90$) to lawrencium ($Z=103$) that follow actinium in period 7; their $5f$ subshell is being filled. | 3.1 3.2 |
| amphoteric | Able to react both as an acid and as a base, e.g. $\mathrm{Al_{2}O_{3}}$ reacts with $\mathrm{HCl}$ and with $\mathrm{NaOH}$. | 3.3 6.3 |
| covalent radius | Half the distance between the nuclei of two identical atoms joined by a single covalent bond, e.g. $r_{\mathrm{Cl}}=\tfrac12(198)=99\ \mathrm{pm}$. | 3.3 4.1 4.3 |
| d-block elements | Elements of groups 3–12 whose last electron enters an inner $(n-1)d$ subshell; outer configuration $(n-1)d^{1-10}ns^{0-2}$. | 2.4 3.1 3.2 |
| diagonal relationship | Similarity between a period-2 element and the period-3 element diagonally below-right of it: Li–Mg, Be–Al, B–Si. | 3.3 |
| Döbereiner's triad | A set of three chemically similar elements in which the atomic mass of the middle one is close to the mean of the other two, e.g. Li, Na, K. | 3.1 |
| effective nuclear charge | The net nuclear charge $Z_{\mathrm{eff}}$ felt by an outer electron after the inner electrons have shielded (screened) part of $Z$. | 2.3 2.4 3.3 |
| eka-element | Mendeleev's name for an undiscovered element one place below a known one, e.g. eka-silicon (germanium). | 3.1 |
| electron gain enthalpy | The enthalpy change when one mole of electrons is added to one mole of isolated gaseous atoms; negative when energy is released (most negative for Cl). | 3.3 4.1 5.3 |
| electronegativity | A relative measure of an atom's ability, in a compound, to attract shared electrons to itself; Pauling scale, F = 4.0. | 3.3 4.1 4.2 4.3 |
| f-block elements | The lanthanoids and actinoids, whose last electrons enter the $(n-2)f$ subshell; shown in two rows below the table. | 3.1 3.2 |
| inner transition elements | Another name for the f-block elements (lanthanoids and actinoids). | 3.2 |
| ionic radius | The radius of an ion in a crystal; cations are smaller and anions larger than their parent atoms. | 3.3 |
| ionisation enthalpy | The enthalpy needed to remove one mole of the most loosely held electrons from one mole of isolated gaseous atoms (or ions): $\mathrm{X(g)\rightarrow X^{+}(g)+e^{-}}$. | 3.2 3.3 4.1 |
| isoelectronic species | Atoms or ions with the same number of electrons, e.g. $\mathrm{O^{2-}}$, $\mathrm{F^{-}}$, $\mathrm{Na^{+}}$ (10 each); their radius falls as $Z$ rises. | 2.4 3.3 4.1 4.2 |
| IUPAC systematic name (of an element) | A temporary name for an element with $Z\gt100$ built from digit roots (nil, un, bi, tri, quad, pent, hex, sept, oct, enn) plus -ium, e.g. 114 = ununquadium (Uuq). | 3.1 |
| lanthanoids | The 14 elements cerium ($Z=58$) to lutetium ($Z=71$) that follow lanthanum in period 6; their $4f$ subshell is being filled. | 3.1 3.2 |
| law of octaves | Newlands' observation (1865) that, with elements in order of atomic mass, every eighth element resembles the first; it held only up to calcium. | 3.1 |
| long form of the periodic table | The present periodic table: elements in order of atomic number in 7 periods and 18 groups, with the lanthanoids and actinoids shown below. | 3.1 |
| Mendeleev's periodic law | The properties of the elements are a periodic function of their atomic masses (1869). | 3.1 |
| metalloid | An element with properties between those of metals and non-metals, lying along the zig-zag line: B, Si, Ge, As, Sb, Te. | 3.2 3.3 |
| modern periodic law | The physical and chemical properties of the elements are periodic functions of their atomic numbers. | 3.1 |
| p-block elements | Elements of groups 13–18 (except He) whose last electron enters an $np$ orbital; outer configuration $ns^{2}np^{1-6}$. | 3.1 3.2 3.3 |
| representative elements | The s- and p-block elements (groups 1, 2 and 13–17; often with 18), whose outer shell alone holds the valence electrons. | 3.2 3.3 4.1 4.2 |
| s-block elements | Elements of groups 1 and 2 (plus H and He) whose last electron enters an $ns$ orbital; outer configuration $ns^{1-2}$. | 3.1 3.2 3.3 |
| transition elements | The d-block elements (groups 3–12): metals with partly filled d subshells that typically show coloured ions and several oxidation states. | 2.4 3.2 7.1 |
| van der Waals radius | Half the distance between the nuclei of identical non-bonded atoms in neighbouring molecules; larger than the covalent radius. | 3.3 9.1 |
Ch 4 · Chemical Bonding and Molecular Structure
| Term | Meaning | Used in |
|---|---|---|
| antibonding molecular orbital | A molecular orbital (marked *) of higher energy than the combining atomic orbitals, with a nodal plane between the nuclei; its electrons weaken the bond. | 4.4 |
| axial and equatorial bonds | In a trigonal bipyramid ($\mathrm{PCl_{5}}$) the two bonds along the axis (axial, longer) and the three in the central plane (equatorial). | 4.2 4.3 |
| bond enthalpy | The energy needed to break one mole of bonds of a given type between atoms in the gaseous state; tables usually give mean values. | 4.1 4.3 4.4 5.3 |
| bond order | The number of bonds between two atoms (Lewis picture), or $\tfrac12(N_{b}-N_{a})$ in molecular orbital theory. Higher bond order means a shorter, stronger bond. | 4.1 4.3 4.4 8.1 |
| bond pair | A pair of electrons shared between two atoms in a covalent bond. | 4.1 4.2 4.3 8.3 |
| bonding molecular orbital | A molecular orbital of lower energy than the combining atomic orbitals, with high electron density between the nuclei; its electrons strengthen the bond. | 4.4 |
| diamagnetic | Weakly repelled by a magnetic field; all electrons are paired, e.g. $\mathrm{N_{2}}$. | 2.4 4.4 |
| dipole moment | Charge × separation, $\mu=Q\times r$, measured in debye ($1\ \mathrm{D}=3.33564\times10^{-30}\ \mathrm{C\,m}$); a vector pointing from the positive to the negative end. | 4.1 4.2 9.2 |
| Fajans' rules | An ionic bond gains covalent character when the cation is small and highly charged (or has a $d^{10}$ shell) and the anion is large. | 4.1 |
| formal charge | A bookkeeping charge on an atom in a Lewis structure: $\text{FC}=V-L-\tfrac12B$. The formal charges add up to the overall charge. | 4.1 |
| hybrid orbital | One of a set of equivalent orbitals formed by hybridisation; hybrid orbitals point as far apart as possible and form σ bonds or hold lone pairs. | 4.3 8.1 9.3 |
| hybridisation | Mixing of atomic orbitals of similar energy on one atom to give an equal number of equivalent hybrid orbitals (sp, sp², sp³, sp³d, sp³d²). | 4.3 8.1 8.3 9.2 |
| hydrogen bond | The attraction between H bonded to F, O or N and a lone pair on another F, O or N atom ($\mathrm{X{-}H\cdots Y}$); intermolecular (water, HF) or intramolecular (2-nitrophenol). | 4.4 7.1 8.2 8.4 |
| ionic bond | The electrostatic attraction between oppositely charged ions formed by complete transfer of electrons, as in $\mathrm{Na^{+}Cl^{-}}$. | 4.1 |
| lattice enthalpy | The energy needed to separate one mole of an ionic solid completely into gaseous ions; $788\ \mathrm{kJ\,mol^{-1}}$ for NaCl. | 4.1 5.3 |
| LCAO | Linear combination of atomic orbitals: molecular orbitals approximated as sums ($\psi_{A}+\psi_{B}$, bonding) and differences ($\psi_{A}-\psi_{B}$, antibonding) of atomic orbitals. | 4.4 |
| Lewis structure | A drawing of a molecule or ion showing all valence electrons: shared pairs as bonds between atoms and lone pairs as dots on atoms. | 4.1 4.2 4.4 8.1 |
| lone pair | A pair of valence electrons held by one atom only and not shared in a bond; it repels other pairs more strongly than a bond pair does. | 4.1 4.2 4.3 4.4 |
| molecular orbital | An orbital spread over the whole molecule, formed by combining atomic orbitals; it holds at most two electrons of opposite spin. | 4.4 |
| octet rule | Main-group atoms tend to bond (by transferring or sharing electrons) until each has eight electrons in its valence shell, the configuration of a noble gas. | 4.1 |
| orbital overlap | Partial merging of orbitals of two atoms in the same region of space; positive (in-phase) overlap builds a bond, and the greater the overlap the stronger the bond. | 4.3 |
| paramagnetic | Attracted into a magnetic field because of unpaired electrons, e.g. $\mathrm{O_{2}}$ (two unpaired electrons in $\pi^{*}2p$). | 2.4 3.2 4.4 8.3 |
| pi bond | A bond formed by sideways overlap of parallel p orbitals, with electron density above and below the internuclear axis; weaker than a σ bond. | 3.3 4.3 8.1 8.2 |
| resonance hybrid | The real structure of a species that needs several Lewis (canonical) structures differing only in electron positions; it is more stable than any one of them. | 4.1 4.3 8.3 9.4 |
| s character | The fraction of s orbital in a hybrid: sp 50%, sp² 33.3%, sp³ 25%. More s character means a shorter bond and a more electronegative atom. | 4.3 9.3 |
| see-saw shape | The shape of a molecule with four bond pairs and one equatorial lone pair in a trigonal bipyramid, e.g. $\mathrm{SF_{4}}$. | 1.2 4.2 4.3 |
| sigma bond | A bond formed by head-on overlap of orbitals along the internuclear axis; every single bond is a σ bond. | 4.3 8.1 8.3 9.1 |
| square planar | The shape with four bond pairs and two trans lone pairs in an octahedral arrangement, e.g. $\mathrm{XeF_{4}}$; all bond angles $90^{\circ}$. | 4.2 4.3 |
| square pyramidal | The shape with five bond pairs and one lone pair in an octahedral arrangement, e.g. $\mathrm{BrF_{5}}$. | 4.2 |
| s–p mixing | Interaction of the σ2s and σ2p_z molecular orbitals in B₂, C₂ and N₂ (small 2s–2p gap), which raises σ2p_z above the π2p pair. | 4.4 |
| T-shaped molecule | The shape with three bond pairs and two equatorial lone pairs in a trigonal bipyramid, e.g. $\mathrm{ClF_{3}}$. | 4.2 |
| trigonal bipyramidal | The arrangement of five electron pairs: three equatorial at $120^{\circ}$ and two axial at $90^{\circ}$ to them, e.g. $\mathrm{PCl_{5}}$. | 4.2 4.3 |
| trigonal pyramidal | The shape of a molecule with three bond pairs and one lone pair on the central atom, e.g. $\mathrm{NH_{3}}$ ($107^{\circ}$). | 4.2 |
| valence bond theory | The theory that a covalent bond forms when half-filled atomic orbitals of two atoms overlap and their electrons, of opposite spin, pair up. | 4.3 |
| VSEPR theory | Valence shell electron pair repulsion theory: electron pairs around a central atom repel and arrange themselves as far apart as possible, which fixes the shape of the molecule. | 4.2 4.3 |
Ch 5 · Thermodynamics
| Term | Meaning | Used in |
|---|---|---|
| adiabatic process | A process in which no heat passes between the system and the surroundings ($q=0$), so $\Delta U=w$. | 5.1 |
| bomb calorimeter | A sealed steel vessel in which a sample is burned in oxygen at constant volume, inside a water bath; it measures $\Delta U$. | 5.2 5.3 |
| Born–Haber cycle | A Hess's-law cycle linking the formation of an ionic solid from its elements with sublimation, bond dissociation, ionisation, electron gain and lattice steps; used to find lattice enthalpies. | 5.3 |
| calorimetry | Measuring the heat of a process from the temperature change it produces in a body of known heat capacity (the calorimeter). | 5.2 5.3 9.4 |
| coffee-cup calorimeter | An insulated cup open to the atmosphere, used for reactions in solution at constant pressure; it measures $\Delta H$. | 5.2 |
| crossover temperature | The temperature $T=\Delta H/\Delta S$ (when $\Delta H$ and $\Delta S$ have the same sign) at which $\Delta G=0$ and the spontaneity of a process changes. | 5.4 |
| endothermic | Absorbing heat from the surroundings; at constant pressure $\Delta H\gt0$. | 3.3 5.2 5.3 5.4 |
| enthalpy | $H=U+pV$, a state function whose change equals the heat absorbed at constant pressure: $\Delta H=q_{p}$. | 3.2 3.3 4.1 4.3 |
| enthalpy of atomisation | $\Delta_{a}H^{\circ}$: the enthalpy change for breaking one mole of a substance completely into gaseous atoms. | 5.3 |
| enthalpy of combustion | $\Delta_{c}H^{\circ}$: the enthalpy change when one mole of a substance burns completely in oxygen under standard conditions; always negative. | 5.3 9.1 |
| enthalpy of fusion / vaporisation / sublimation | Enthalpies of phase change: the heat absorbed per mole on melting (fusion), boiling (vaporisation) or direct solid-to-vapour change (sublimation), at constant pressure. | 5.2 5.3 |
| enthalpy of solution / hydration / dilution | Heat absorbed when one mole of a substance dissolves ($\Delta_{sol}H=\Delta_{lattice}H+\Delta_{hyd}H$), when gaseous ions are hydrated, or when a solution is diluted further. | 5.2 5.3 |
| entropy | A state function $S$ measuring disorder — how widely energy and particles are spread; for a reversible change $\Delta S=q_{rev}/T$. | 5.1 5.4 |
| exothermic | Giving out heat to the surroundings; at constant pressure $\Delta H\lt0$. | 5.2 5.3 5.4 6.2 |
| extensive property | A property that depends on the amount of matter in the system, such as mass, volume, internal energy or heat capacity. | 5.1 5.2 5.4 |
| first law of thermodynamics | Energy can be neither created nor destroyed; for a closed system $\Delta U=q+w$, with $q$ and $w$ positive when energy enters the system. | 5.1 |
| Gibbs energy | $G=H-TS$; at constant $T$ and $p$, $\Delta G=\Delta H-T\Delta S=-T\Delta S_{total}$, negative for a spontaneous process and zero at equilibrium. | 5.1 5.4 6.2 7.3 |
| heat capacity | The heat needed to raise the temperature of a body by $1\ \mathrm{K}$, $C=q/\Delta T$; at constant volume $C_{v}$, at constant pressure $C_{p}$. | 2.2 5.1 5.2 |
| Hess's law | The enthalpy change of a reaction is the same whether it occurs in one step or in several; the step enthalpies add. It holds because $H$ is a state function. | 5.3 9.4 |
| intensive property | A property independent of the amount of matter, such as temperature, pressure, density or molar heat capacity. | 5.1 5.2 7.3 |
| internal energy | The total energy stored in a system — molecular kinetic and potential energy, bond and nuclear energy; written $U$. Only its change $\Delta U$ can be measured. | 5.1 5.2 |
| isolated system | A system that exchanges neither energy nor matter with its surroundings; its internal energy cannot change. | 5.1 5.4 |
| mean bond enthalpy | The average enthalpy needed to break one mole of a given type of bond (e.g. C–H) in gaseous molecules, averaged over a molecule or over many compounds. | 4.1 4.3 5.3 9.2 |
| molar heat capacity | Heat capacity per mole, $C_{m}$ in $\mathrm{J\,mol^{-1}\,K^{-1}}$; for an ideal gas $C_{p}-C_{v}=R$. | 5.1 5.2 |
| reversible process | An idealised process carried out through states infinitesimally close to equilibrium, so that an infinitesimal change in conditions would reverse it. | 5.1 |
| second law of thermodynamics | In every spontaneous process the total entropy of system and surroundings increases: $\Delta S_{total}\gt0$ (zero at equilibrium). | 5.4 |
| specific heat capacity | Heat capacity per gram, $c$ in $\mathrm{J\,g^{-1}\,K^{-1}}$, so that $q=mc\Delta T$; $4.18\ \mathrm{J\,g^{-1}\,K^{-1}}$ for liquid water. | 5.2 |
| spontaneous process | A process that, once started, continues by itself without outside help; it need not be fast. At constant $T$ and $p$, $\Delta G\lt0$. | 5.4 7.3 |
| standard enthalpy of formation | $\Delta_{f}H^{\circ}$: the enthalpy change when one mole of a compound forms from its elements in their reference states; zero for an element in its reference state. | 5.3 |
| standard molar entropy | $S^{\circ}$: the entropy of one mole of a substance in its standard state, measured from zero at $0\ \mathrm{K}$; not zero for elements at $298\ \mathrm{K}$. | 5.4 |
| standard state | The pure substance in its stable form at $1\ \mathrm{bar}$ (a solute at $1\ \mathrm{mol\,L^{-1}}$) and a stated temperature, usually $298\ \mathrm{K}$; marked by the symbol $^{\circ}$. | 5.3 5.4 6.2 7.3 |
| state function | A property such as $p$, $V$, $T$ or $U$ whose value depends only on the present state of a system, not on the path by which it was reached. | 5.1 5.2 5.3 5.4 |
| surroundings | Everything outside a thermodynamic system that can exchange energy or matter with it; system + surroundings = universe. | 5.1 5.2 5.4 |
| thermochemical equation | A balanced equation with state symbols together with its enthalpy change, e.g. $\mathrm{C(s)+O_{2}(g)\rightarrow CO_{2}(g)}$; $\Delta_{r}H^{\circ}=-393.5\ \mathrm{kJ\,mol^{-1}}$. | 5.3 |
| thermodynamic system | The part of the universe chosen for study in thermodynamics, separated from its surroundings by a real or imaginary boundary. | |
| third law of thermodynamics | The entropy of a perfectly crystalline substance approaches zero as $T\to0\ \mathrm{K}$, which allows absolute entropies $S^{\circ}$ to be tabulated. | 5.4 |
Ch 6 · Equilibrium
| Term | Meaning | Used in |
|---|---|---|
| Brønsted–Lowry acid and base | A Brønsted–Lowry acid is a proton donor and a Brønsted–Lowry base is a proton acceptor. | 6.3 |
| buffer solution | A solution, usually a weak acid with its salt or a weak base with its salt, that resists a change of pH when a little acid or base is added or when it is diluted. | 6.3 6.4 |
| catalyst | A substance that speeds up a reaction without being used up; it speeds up the forward and backward reactions equally, so it does not change $K$ or the equilibrium composition. | 1.3 3.2 5.4 6.2 |
| chemical equilibrium | The state of a reversible reaction in a closed system in which the forward and backward rates are equal, so the concentrations of all species stay constant. | 6.1 |
| common-ion effect | The suppression of the ionisation (or solubility) of a weak electrolyte by adding a strong electrolyte that supplies one of the same ions. | 6.3 6.4 |
| conjugate acid–base pair | Two species that differ by exactly one proton, such as $\mathrm{NH_{4}^{+}}$ and $\mathrm{NH_{3}}$; for the pair $K_{a}K_{b}=K_{w}$. | 6.3 6.4 9.3 |
| contact process | The industrial manufacture of sulphuric acid, in which sulphur dioxide is oxidised to sulphur trioxide over a vanadium(V) oxide catalyst. | 6.2 |
| degree of dissociation | The fraction $\alpha$ of the molecules of a substance that have dissociated when equilibrium is reached. | 6.2 |
| degree of ionisation | The fraction $\alpha$ of a weak electrolyte present as ions at equilibrium; for a weak acid $\alpha\approx\sqrt{K_{a}/c}$, so it increases on dilution. | 6.3 |
| dynamic equilibrium | A state in which two opposing processes go on at equal rates, so that nothing observable changes although activity continues at the molecular level. | 6.1 6.4 |
| equilibrium constant | The constant $K_{c}$ (concentrations) or $K_{p}$ (partial pressures) given by the law of chemical equilibrium; it depends only on temperature and on how the equation is written. | 5.4 6.1 6.2 6.3 |
| Haber process | The industrial synthesis of ammonia, $\mathrm{N_{2}+3H_{2}\rightleftharpoons2NH_{3}}$, at about 700–775 K and 200 atm over an iron catalyst, with the ammonia removed and the unreacted gases recycled. | 5.4 6.2 |
| Henderson equation | $\mathrm{pH}=\mathrm{p}K_{a}+\log([\text{salt}]/[\text{acid}])$ for an acidic buffer (and $\mathrm{pOH}=\mathrm{p}K_{b}+\log([\text{salt}]/[\text{base}])$ for a basic one). | 6.4 |
| Henry's law | At constant temperature the mass of a gas dissolved in a given mass of solvent is proportional to the pressure of the gas above the solvent. | 6.1 |
| heterogeneous equilibrium | An equilibrium involving more than one phase, e.g. a solid and a gas; pure solids and liquids are left out of its equilibrium constant. | 6.1 6.4 |
| homogeneous equilibrium | An equilibrium in which all reactants and products are in the same phase, e.g. all gases. | 6.1 |
| hydrolysis constant | The equilibrium constant $K_{h}$ of a hydrolysis reaction: $K_{w}/K_{a}$ for the anion of a weak acid, $K_{w}/K_{b}$ for the cation of a weak base. | 6.4 |
| ICE table | A table of Initial values, Changes (in terms of one unknown $x$) and Equilibrium values used to calculate an equilibrium composition from $K$. | 6.2 6.3 |
| ionic product (of a salt) | $Q_{sp}$, the $K_{sp}$ expression evaluated with the actual ion concentrations; a precipitate forms if $Q_{sp}\gt K_{sp}$. | 6.4 |
| ionic product of water | $K_{w}=[\mathrm{H^{+}}][\mathrm{OH^{-}}]$, equal to $1.0\times10^{-14}$ at 298 K; it increases with temperature. | 6.3 |
| law of chemical equilibrium | At equilibrium at a given temperature, the product of the product concentrations divided by the product of the reactant concentrations, each raised to its stoichiometric coefficient, is a constant. | 6.1 |
| Le Chatelier's principle | If a system at equilibrium is disturbed by a change of concentration, pressure or temperature, the equilibrium shifts in the direction that tends to counteract the change. | 6.2 6.3 6.4 |
| Lewis acid and base | A Lewis acid accepts an electron pair (e.g. $\mathrm{BF_{3}}$, $\mathrm{H^{+}}$, metal ions); a Lewis base donates an electron pair (e.g. $\mathrm{NH_{3}}$, $\mathrm{OH^{-}}$). | 6.3 9.4 |
| molar solubility | The number of moles of a salt that dissolve in one litre of its saturated solution; for a salt AB in water $S=\sqrt{K_{sp}}$, for $\mathrm{AB_{2}}$ $S=\sqrt[3]{K_{sp}/4}$. | 6.4 |
| normal boiling point | The temperature at which the vapour pressure of a liquid equals 1.013 bar (1 atm); 100 °C for water. | 5.4 6.1 8.4 9.1 |
| pH scale | The scale $\mathrm{pH}=-\log[\mathrm{H^{+}}]$; at 298 K a solution is acidic below 7, neutral at 7 and basic above 7. | 6.3 |
| polybasic acid | An acid with more than one ionisable proton, such as $\mathrm{H_{2}S}$ or $\mathrm{H_{3}PO_{4}}$; it ionises in steps with $K_{a_{1}}\gt K_{a_{2}}\gt K_{a_{3}}$. | 6.3 |
| reaction quotient | The expression $Q$ with the same form as the equilibrium constant but using the concentrations (or pressures) present at any moment; $Q\lt K$ means net forward reaction, $Q\gt K$ net backward. | 6.2 6.4 |
| salt hydrolysis | The reaction of the cation or anion of a salt with water, which makes the solution acidic or basic; salts of a strong acid and a strong base do not hydrolyse. | 6.4 |
| saturated solution | A solution in which no more solute dissolves at the given temperature: the dissolved solute is in dynamic equilibrium with undissolved solid. | 6.1 6.3 6.4 |
| solubility product | $K_{sp}$, the product of the ion concentrations of a saturated solution of a sparingly soluble salt, each raised to the power of its coefficient; it depends only on temperature. | 6.4 |
| sparingly soluble salt | A salt of which less than about 0.01 mol dissolves per litre, such as AgCl or $\mathrm{BaSO_{4}}$. | 6.4 8.4 |
| strong electrolyte | A substance that is almost completely ionised in aqueous solution, e.g. HCl, NaOH and most salts. | 6.3 |
| vapour pressure | The pressure of the vapour in equilibrium with its liquid (or solid) in a closed container at a given temperature; it depends only on the substance and the temperature. | 1.2 6.1 8.4 |
| weak electrolyte | A substance only partly ionised in solution, so that its molecules are in equilibrium with its ions, e.g. acetic acid and ammonia. | 6.3 |
Ch 7 · Redox Reactions
| Term | Meaning | Used in |
|---|---|---|
| activity series | The metals listed in order of their tendency to lose electrons in aqueous solution (K, Na, Ca, Mg, Al, Zn, Fe, … H, Cu, Hg, Ag, Au); a metal displaces the ions of any metal below it. | 7.1 7.2 7.3 |
| anode | The electrode at which oxidation occurs; in a galvanic cell it is the negative terminal (zinc in the Daniell cell). | 2.1 7.3 9.1 |
| cathode | The electrode at which reduction occurs; in a galvanic cell it is the positive terminal (copper in the Daniell cell). | 2.1 7.3 9.1 |
| cell EMF | The voltage of a galvanic cell when no current is drawn; $E^{\circ}_{\text{cell}}=E^{\circ}_{\text{cathode}}-E^{\circ}_{\text{anode}}$, positive for a spontaneous cell reaction. | 7.3 |
| combination reaction | Two or more substances combine to form one product, $\mathrm{A+B\rightarrow C}$; it is a redox reaction when at least one reactant is an element, e.g. $\mathrm{C+O_{2}\rightarrow CO_{2}}$. | 7.2 |
| competitive electron transfer reaction | A reaction in which two redox couples compete for electrons, such as a metal displacing another metal's ions from solution; the species that holds electrons more weakly gives them up. | 7.1 7.2 |
| comproportionation | The reverse of disproportionation: an element in two oxidation states forms a single intermediate state, e.g. $\mathrm{2H_{2}S+SO_{2}\rightarrow 3S+2H_{2}O}$. | 7.2 |
| Daniell cell | The galvanic cell $\mathrm{Zn(s)\,|\,Zn^{2+}(aq)\,\|\,Cu^{2+}(aq)\,|\,Cu(s)}$; with $1\ \mathrm{mol\,L^{-1}}$ solutions its EMF is $1.10\ \mathrm{V}$. | 7.3 |
| decomposition reaction | One compound breaks down into two or more substances; it is a redox reaction when at least one product is an element, e.g. $\mathrm{2KClO_{3}\rightarrow 2KCl+3O_{2}}$. | 7.2 |
| displacement reaction | An atom or ion in a compound is replaced by one of another element, $\mathrm{X+YZ\rightarrow XZ+Y}$: a more reactive metal displaces a less reactive one, or hydrogen, oxygen or a halogen is displaced. | 7.2 |
| disproportionation | A redox reaction in which one element in a single oxidation state is both oxidised and reduced, e.g. $\mathrm{Cl_{2}+2OH^{-}\rightarrow Cl^{-}+ClO^{-}+H_{2}O}$. | 7.2 |
| electrochemical series | Redox couples listed in order of their standard reduction potentials; it predicts which oxidising agent can oxidise which reducing agent. | 7.3 |
| galvanic cell | A device in which a spontaneous redox reaction produces electricity, the oxidation and reduction taking place in separate half-cells joined by a wire and a salt bridge, e.g. the Daniell cell. | 7.3 |
| half-reaction | The oxidation part or the reduction part of a redox reaction written on its own with the electrons shown, e.g. $\mathrm{Zn\rightarrow Zn^{2+}+2e^{-}}$. | 7.1 7.2 7.3 |
| iodometric titration | An oxidising agent liberates iodine from excess iodide, and the iodine is titrated with thiosulphate using starch (blue colour disappears at the end point). | 7.2 |
| ion–electron method | Balancing a redox equation by balancing the oxidation and reduction half-reactions separately (atoms, O with $\mathrm{H_{2}O}$, H with $\mathrm{H^{+}}$, charge with $\mathrm{e^{-}}$) and adding them so the electrons cancel. | 7.2 |
| oxidation number | The charge an atom would carry if all its bonds were ionic, with each shared pair given to the more electronegative atom; assigned by fixed rules, e.g. Mn is $+7$ in $\mathrm{MnO_{4}^{-}}$. | 3.2 3.3 7.1 7.2 |
| oxidation-number method | Balancing a redox equation by making the total increase in oxidation number equal to the total decrease, then adding $\mathrm{H^{+}}$/$\mathrm{OH^{-}}$ and $\mathrm{H_{2}O}$. | 7.2 |
| oxidising agent | A species that accepts electrons from another (raises another element's oxidation number) and is itself reduced, e.g. $\mathrm{Cu^{2+}}$, $\mathrm{Cl_{2}}$, $\mathrm{MnO_{4}^{-}}$. | 3.3 7.1 7.2 7.3 |
| peroxide | A compound with an O–O single bond, in which each oxygen has oxidation number $-1$, e.g. $\mathrm{H_{2}O_{2}}$, $\mathrm{Na_{2}O_{2}}$, $\mathrm{BaO_{2}}$. | 1.2 1.3 3.3 4.1 |
| redox couple | The oxidised and reduced forms of a species that interconvert in a half-reaction, written oxidised/reduced, e.g. $\mathrm{Zn^{2+}/Zn}$, $\mathrm{Fe^{3+}/Fe^{2+}}$. | 7.1 7.3 |
| redox reaction | A reaction in which one species is oxidised and another is reduced at the same time; the oxidation number of at least one element changes, e.g. $\mathrm{Zn+Cu^{2+}\rightarrow Zn^{2+}+Cu}$. | 7.1 7.2 7.3 |
| redox titration | A titration in which titrant and analyte react by electron transfer, e.g. $\mathrm{Fe^{2+}}$ with acidified $\mathrm{KMnO_{4}}$ or iodine with thiosulphate. | 7.2 |
| reducing agent | A species that gives electrons to another (lowers another element's oxidation number) and is itself oxidised, e.g. Zn, $\mathrm{H_{2}}$, CO. | 3.3 7.1 7.2 7.3 |
| sacrificial protection | Protecting iron or steel from corrosion by connecting it to a more reactive metal (Zn, Mg, Al) that is oxidised in its place, as in galvanising. | 7.3 |
| salt bridge | A tube of an inert electrolyte (KCl or $\mathrm{KNO_{3}}$ in agar) joining the two half-cells; it completes the circuit and keeps both solutions electrically neutral. | 7.3 |
| self-indicator | A titrant whose own colour marks the end point, such as purple permanganate, which is decolourised until the first permanent pink excess. | 7.2 |
| spectator ion | An ion present in the solution that takes no part in the reaction and is left out of the ionic equation, e.g. $\mathrm{SO_{4}^{2-}}$ when zinc reacts with copper(II) sulphate. | 7.1 |
| standard electrode potential | $E^{\circ}$: the potential of a half-cell, written as a reduction, with all species in their standard states, measured against the standard hydrogen electrode; more positive means a stronger oxidising agent. | 3.3 7.3 |
| standard hydrogen electrode | Platinised platinum in $1\ \mathrm{mol\,L^{-1}}$ $\mathrm{H^{+}}$ with $\mathrm{H_{2}}$ at $1\ \mathrm{bar}$; the reference half-cell whose potential is taken as $0.00\ \mathrm{V}$. | 7.3 |
| Stock notation | Naming that shows the oxidation number of a metal as a Roman numeral in brackets: iron(III) chloride for $\mathrm{FeCl_{3}}$, copper(I) oxide for $\mathrm{Cu_{2}O}$. | 7.1 |
| superoxide | A compound containing the $\mathrm{O_{2}^{-}}$ ion, in which each oxygen has oxidation number $-\tfrac12$, e.g. $\mathrm{KO_{2}}$. | 3.3 4.4 7.1 |
Ch 8 · Organic Chemistry – Some Basic Principles and Techniques
| Term | Meaning | Used in |
|---|---|---|
| alicyclic compound | A ring compound of carbon atoms that behaves like an open-chain (aliphatic) compound, e.g. cyclohexane or cyclohexene. | 8.1 |
| alkyl group | A group derived from an alkane by removing one hydrogen atom (CₙH₂ₙ₊₁–), named by changing -ane to -yl: methyl, ethyl, propyl, isopropyl (1-methylethyl), tert-butyl… | 8.2 8.3 9.1 9.2 |
| bond-line formula | A skeletal drawing of a molecule: every line end and vertex is a carbon atom carrying enough hydrogens to make four bonds; only other atoms (and H on them) are written. | 8.1 9.1 |
| carbanion | An ion with a negatively charged carbon carrying a lone pair (eight valence electrons); sp³ and pyramidal. Stability: CH₃⁻ > 1° > 2° > 3°. | 8.3 9.3 9.4 |
| carbocation | An ion with a positively charged carbon that has only six valence electrons; sp² and planar with an empty p orbital. Stability: 3° > 2° > 1° > CH₃⁺. | 8.3 9.2 9.4 |
| Carius method | Estimation of halogens (as AgX) or sulphur (as BaSO₄) by heating the compound with fuming nitric acid in a sealed tube and weighing the precipitate. | 8.4 |
| catenation | The ability of atoms of an element to bond to one another in long chains and rings; carbon shows it more than any other element. | 8.1 |
| chain isomerism | Structural isomerism in which the compounds have different carbon skeletons, e.g. pentane, 2-methylbutane and 2,2-dimethylpropane. | 8.2 9.1 9.2 9.3 |
| chiral carbon | A carbon atom bonded to four different atoms or groups; a molecule with one such carbon is not superimposable on its mirror image. | 8.2 |
| chromatography | Separation of a mixture by the different rates at which its components are carried by a mobile phase over a stationary phase (adsorption: column, TLC; partition: paper). | 8.4 |
| condensed structural formula | A structural formula in which the bonds to hydrogen (and often the C–C bonds) are omitted and the atoms on each carbon are grouped, e.g. CH₃CH₂CH₂OH or (CH₃)₃CH. | 8.1 8.2 9.1 9.3 |
| crystallisation | Purification of a solid by dissolving it in a minimum of hot solvent and letting it crystallise on cooling, while the impurities stay in solution. | 1.1 1.2 8.4 |
| differential extraction | Removing an organic compound from water by shaking with an immiscible solvent in which it is more soluble, then separating the layers in a separating funnel. | 8.4 |
| electromeric effect | The temporary, complete transfer of the π electrons of a multiple bond to one atom when an attacking reagent approaches (+E towards, −E away from the atom attacked). | 8.3 |
| electrophile | An electron-poor reagent that accepts an electron pair from an electron-rich centre, e.g. H⁺, NO₂⁺, carbocations, BF₃, AlCl₃. | 8.3 9.2 9.3 9.4 |
| enantiomers | A pair of stereoisomers that are non-superimposable mirror images; they rotate plane-polarised light equally in opposite directions but have the same boiling point, density, etc. | 8.2 |
| fractional distillation | Distillation through a fractionating column, which gives many vaporisation–condensation steps; separates liquids whose boiling points are close, e.g. the fractions of crude oil. | 8.4 9.1 |
| free radical | A neutral species with an unpaired electron, e.g. CH₃• or Cl•; alkyl radicals are nearly planar, very reactive and paramagnetic. | 8.3 |
| functional group | An atom or group of atoms joined in a specific way that gives a family of organic compounds its characteristic chemical properties, e.g. –OH, –CHO, –COOH, C=C. | 8.1 8.2 |
| functional isomerism | Structural isomerism between compounds with different functional groups, e.g. ethanol and methoxymethane, or propanal and propanone. | 8.2 9.3 |
| heterocyclic compound | A ring compound in which at least one ring atom is not carbon (usually N, O or S), e.g. pyridine, furan, tetrahydrofuran. | 8.1 |
| heterolytic fission | Breaking of a covalent bond in which the more electronegative atom takes both electrons, giving a cation and an anion; favoured by polar bonds and polar solvents. | 8.3 |
| homologous series | A family of compounds with the same functional group and general formula in which successive members differ by a CH₂ unit (14 u). | 8.1 9.1 |
| homolytic fission | Breaking of a covalent bond in which each atom keeps one electron of the shared pair, giving two free radicals; favoured by heat, light and non-polar conditions. | 8.3 9.1 |
| hyperconjugation | Delocalisation of the σ electrons of C–H bonds on a carbon next to a carbocation, radical or π bond ('no-bond resonance'); its extent grows with the number of α-hydrogens. | 8.3 9.2 9.4 |
| inductive effect | The permanent displacement of σ electrons along a chain caused by an atom or group of different electronegativity; −I groups withdraw electrons, +I groups (alkyl) release them. It fades within about three bonds. | 8.3 9.2 |
| isomers | Different compounds with the same molecular formula. Structural isomers differ in the order in which atoms are joined; stereoisomers only in the arrangement of atoms in space. | 4.3 4.4 8.1 8.2 |
| IUPAC nomenclature | The systematic naming of compounds agreed by the International Union of Pure and Applied Chemistry: word root (chain length) + suffixes (saturation, principal group) + prefixes (substituents) with locants; each name corresponds to one structure. | 8.1 8.2 9.1 9.2 |
| Kjeldahl method | Estimation of nitrogen by digesting the compound with conc. H₂SO₄ to ammonium sulphate, liberating NH₃ with NaOH and absorbing it in standard acid: %N = 1.4 n(NH₃, mmol)/m(g). | 8.4 |
| Lassaigne's test | Detection of N, S and halogens by fusing the compound with sodium to form NaCN, Na₂S and NaX, then testing the aqueous extract (Prussian blue, black PbS, silver halides). | 8.4 |
| locant | The number giving the position of a substituent, multiple bond or functional group on the parent chain; the chain is numbered to give the lowest set of locants, compared at the first point of difference. | 8.2 9.1 9.2 9.3 |
| metamerism | Isomerism caused by different alkyl groups on the two sides of the same functional group, e.g. ethoxyethane and 1-methoxypropane. | 8.2 |
| nucleophile | An electron-rich reagent that attacks an electron-poor centre by donating an electron pair, e.g. OH⁻, CN⁻, H₂O, NH₃. | 8.3 9.3 9.4 |
| organic compound | A hydrocarbon or a compound derived from one by replacing hydrogen atoms with other atoms or groups; CO, CO₂, carbonates, cyanides and carbides are traditionally counted as inorganic. | 8.1 8.4 |
| parent chain | The chain of carbon atoms that gives a compound its root name: the longest chain (containing the principal group and multiple bonds), and if there is a tie, the one with more substituents. | 8.2 |
| position isomerism | Structural isomerism in which the skeleton is the same but a substituent, functional group or multiple bond is at a different position, e.g. propan-1-ol and propan-2-ol. | 8.2 9.2 9.3 9.4 |
| resonance effect | The permanent polarity produced in a conjugated system by a group that releases (+R: –OH, –NH₂, halogens) or withdraws (−R: –NO₂, –CN, –CHO) π electrons. | 8.3 9.4 |
| retardation factor | Rf = distance moved by a substance ÷ distance moved by the solvent front, both from the base line; between 0 and 1 and characteristic of the compound for given conditions. | 8.4 |
| steam distillation | Distillation of a water-immiscible, steam-volatile compound with steam; the mixture boils below 373 K because p(organic) + p(water) = atmospheric pressure. | 4.4 8.4 |
| stereoisomers | Isomers with the same structural formula but a different arrangement of atoms in space: geometrical (cis–trans) and optical isomers. | 8.2 9.2 |
| sublimation (purification) | A purification method based on the direct change of a solid into vapour on heating (and back on cooling); used to purify camphor, naphthalene and similar solids from non-volatile impurities. | |
| wedge–dash formula | A three-dimensional drawing: plain lines lie in the plane of the paper, a solid wedge points towards the viewer and a hashed wedge away. | 8.1 |
Ch 9 · Hydrocarbons
| Term | Meaning | Used in |
|---|---|---|
| addition polymerisation | Joining of many monomer molecules, e.g. alkenes, each π bond becoming a σ bond to the next unit, with no atoms lost: $\mathrm{nCH_{2}{=}CH_{2}\rightarrow{-}(CH_{2}CH_{2})_{n}{-}}$. | 9.2 9.3 9.4 |
| alkane | An open-chain saturated hydrocarbon, general formula $\mathrm{C_{n}H_{2n+2}}$, with only single bonds and sp³ carbon atoms (methane, ethane, propane…). | 8.1 8.2 9.1 9.2 |
| alkene | An unsaturated hydrocarbon with a C=C double bond (one σ + one π bond); general formula $\mathrm{C_{n}H_{2n}}$ for one double bond and no ring. | 8.1 8.2 8.3 9.1 |
| alkyne | An unsaturated hydrocarbon with a C≡C triple bond (one σ + two π bonds, sp carbons, linear); general formula $\mathrm{C_{n}H_{2n-2}}$ for one triple bond and no ring. | 4.3 8.1 9.1 9.2 |
| alkynide (acetylide) | The salt formed when a terminal alkyne loses its ≡C–H proton to a strong base, e.g. sodium ethynide $\mathrm{HC{\equiv}C^{-}Na^{+}}$; decomposed by water. | 4.4 9.3 |
| arene (aromatic hydrocarbon) | A hydrocarbon containing one or more benzene-type rings, e.g. benzene, methylbenzene, naphthalene. | 9.1 9.4 |
| arenium ion (σ-complex) | The resonance-stabilised carbocation formed when an electrophile bonds to one carbon of a benzene ring; that carbon becomes sp³ and the ring is temporarily not aromatic. | 9.4 |
| aromaticity | The special stability of a cyclic, planar, fully conjugated system with $(4n+2)$ delocalised π electrons, as in benzene. | 9.4 |
| Baeyer's reagent | Cold, dilute, alkaline aqueous $\mathrm{KMnO_{4}}$; converts an alkene into a 1,2-diol and loses its purple colour — a test for unsaturation. | 9.2 |
| calcium carbide | $\mathrm{CaC_{2}}$, the ionic compound $\mathrm{Ca^{2+}C_{2}^{2-}}$, made from CaO and coke in an electric furnace; water converts it into ethyne. | 4.4 9.3 |
| carcinogen | A substance that can cause cancer; benzene (leukaemia) and polynuclear aromatic hydrocarbons such as benzo[a]pyrene are examples. | 9.4 |
| chain reaction (free-radical) | A reaction that proceeds by initiation, a cycle of propagation steps in which each step consumes one radical and produces another, and termination; e.g. the halogenation of alkanes. | 4.4 8.3 9.1 9.2 |
| conformation | One of the spatial arrangements of a molecule that interconvert by rotation about single bonds, described by the dihedral angle. | 9.1 |
| cyclic polymerisation | Joining of three ethyne molecules into one benzene ring when ethyne is passed through a red-hot iron tube at about $873\ \mathrm{K}$: $\mathrm{3C_{2}H_{2}\rightarrow C_{6}H_{6}}$. | 9.3 9.4 |
| decarboxylation | Removal of $\mathrm{CO_{2}}$ from a carboxylic acid salt; heating $\mathrm{RCOONa}$ with soda lime gives the alkane $\mathrm{RH}$ with one carbon fewer. | 9.1 9.4 |
| dehydrohalogenation | β-Elimination of HX from an alkyl halide by heating with alcoholic KOH, giving an alkene; ease 3° > 2° > 1° and RI > RBr > RCl. | 8.3 9.2 9.3 |
| dihedral angle | Viewed along a bond, the angle between a bond on the front atom and a bond on the back atom. | 9.1 |
| electrophilic substitution | Replacement of a ring H of an arene by an electrophile $\mathrm{E^{+}}$ via an arenium ion, which then loses $\mathrm{H^{+}}$; nitration, halogenation, sulphonation and Friedel–Crafts reactions. | 9.4 |
| enol | A compound with an OH group on a C=C carbon, e.g. $\mathrm{CH_{2}{=}CHOH}$; usually unstable, rearranging at once to an aldehyde or ketone (formed in the hydration of alkynes). | 8.3 9.3 |
| Friedel–Crafts reaction | Alkylation ($\mathrm{RCl}$) or acylation ($\mathrm{RCOCl}$ or anhydride) of an arene catalysed by anhydrous $\mathrm{AlCl_{3}}$, via $\mathrm{R^{+}}$ or $\mathrm{RCO^{+}}$. | 9.4 |
| geminal dihalide | A compound with two halogen atoms on the same carbon, e.g. $\mathrm{CH_{3}CHBr_{2}}$; formed when HX adds twice to an alkyne by Markovnikov's rule. | 9.3 |
| geometrical (cis–trans) isomerism | Stereoisomerism caused by restricted rotation about a double bond: cis — similar groups on the same side; trans — on opposite sides. Needs two different groups on each doubly bonded carbon. | 4.3 8.1 8.2 9.1 |
| hydrocarbon | A compound containing only carbon and hydrogen: alkanes, alkenes, alkynes, cycloalkanes and arenes. | 1.3 8.1 8.3 8.4 |
| Hückel's rule | A planar, cyclic, fully conjugated molecule or ion is aromatic if it has $(4n+2)$ π electrons ($n=0,1,2,\ldots$): 2, 6, 10, 14… | 9.4 |
| Kekulé structure | A six-membered carbon ring with alternating single and double bonds; benzene is a resonance hybrid of two such structures. | 4.3 5.3 8.1 9.4 |
| Kolbe's electrolysis | Electrolysis of an aqueous sodium or potassium carboxylate $\mathrm{RCOO^{-}}$: alkyl radicals formed at the anode pair up to give $\mathrm{R{-}R}$, with $\mathrm{CO_{2}}$; hydrogen forms at the cathode. | 9.1 |
| Lindlar's catalyst | Palladium on charcoal or calcium carbonate partly deactivated with sulphur compounds or quinoline; reduces an alkyne only as far as the cis alkene. | 9.2 9.3 |
| Markovnikov's rule | In adding HX to an unsymmetrical alkene, the negative part X goes to the doubly bonded carbon with fewer hydrogen atoms, because the reaction goes through the more stable carbocation. | 9.2 9.3 |
| Newman and sawhorse projections | Drawings of conformations: Newman — viewed along the C–C bond, front carbon a point, back carbon a circle; sawhorse — the C–C bond drawn as a slanting line with the bonds on each carbon. | 9.1 |
| nitronium ion | $\mathrm{NO_{2}^{+}}$, the electrophile in nitration, formed from concentrated $\mathrm{HNO_{3}}$ and $\mathrm{H_{2}SO_{4}}$ (the nitrating mixture). | 9.4 |
| ortho/para and meta directors | Groups on a benzene ring that send a second electrophile to the 2/4 positions (electron-releasing groups and halogens) or to the 3 position (withdrawing groups –NO₂, –CHO, –COOH, –CN, –SO₃H). | 9.4 |
| ozonolysis | Reaction of an alkene with ozone (forming an ozonide) followed by Zn and water; the C=C is cut and each doubly bonded carbon becomes C=O, giving aldehydes and/or ketones. | 9.2 9.4 |
| peroxide effect (anti-Markovnikov addition) | In the presence of a peroxide, HBr (only) adds to an unsymmetrical alkene against Markovnikov's rule, by a free-radical chain in which Br• adds first to the CH₂ end. | 9.2 |
| pyrolysis (cracking) | Breaking a large alkane into smaller alkanes, alkenes and hydrogen by strong heating in the absence of air, often with a catalyst. | 9.1 9.2 |
| sodamide | $\mathrm{NaNH_{2}}$, a very strong base (the amide ion $\mathrm{NH_{2}^{-}}$); removes HX from vinylic halides and H⁺ from terminal alkynes. | 9.3 |
| staggered and eclipsed conformations | In ethane, staggered (dihedral $60^{\circ}$): back C–H bonds midway between front ones, lowest energy; eclipsed ($0^{\circ}$): back bonds directly behind front ones, about $12.5\ \mathrm{kJ\,mol^{-1}}$ higher. | 9.1 |
| terminal alkyne | An alkyne with the triple bond at the end of the chain, $\mathrm{R{-}C{\equiv}C{-}H}$; its ≡C–H hydrogen is weakly acidic. Internal alkynes $\mathrm{R{-}C{\equiv}C{-}R'}$ have none. | 4.3 9.3 |
| torsional strain | The extra energy of a conformation caused by repulsion between bond electron clouds on adjacent atoms; greatest when bonds are eclipsed. | 9.1 |
| vicinal dihalide | A compound with two halogen atoms on adjacent carbon atoms, e.g. 1,2-dibromoethane; zinc removes both halogens to give an alkene. | 9.2 9.3 |
| Wurtz reaction | Joining two alkyl groups by heating an alkyl halide with sodium in dry ether: $\mathrm{2RX+2Na\rightarrow R{-}R+2NaX}$. | 9.1 |