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 · SolutionsCh 2 · ElectrochemistryCh 3 · Chemical KineticsCh 4 · The d- and f-Block ElementsCh 5 · Coordination CompoundsCh 6 · Haloalkanes and HaloarenesCh 7 · Alcohols, Phenols and EthersCh 8 · Aldehydes, Ketones and Carboxylic AcidsCh 9 · AminesCh 10 · Biomolecules
Ch 1 · Solutions
| Term | Meaning | Used in |
|---|---|---|
| azeotrope | A liquid mixture that has the same composition as its vapour and boils at a constant temperature, so it cannot be separated by fractional distillation. | 1.2 |
| colligative property | A property of a dilute solution that depends only on the number of solute particles, not on their nature: relative lowering of vapour pressure, $\Delta T_{b}$, $\Delta T_{f}$ and osmotic pressure. | 1.3 |
| cryoscopic constant | $K_{f}$, the molal depression constant: the freezing-point fall of a $1\ \mathrm{mol\,kg^{-1}}$ solution; it depends only on the solvent ($1.86\ \mathrm{K\,kg\,mol^{-1}}$ for water). | 1.3 |
| depression of freezing point | The fall in freezing point of a solvent caused by a dissolved solute: $\Delta T_{f}=K_{f}\,m$ (times $i$ for electrolytes). | 1.3 |
| ebullioscopic constant | $K_{b}$, the molal elevation constant: the boiling-point rise of a $1\ \mathrm{mol\,kg^{-1}}$ solution; it depends only on the solvent ($0.52\ \mathrm{K\,kg\,mol^{-1}}$ for water). | 1.3 |
| elevation of boiling point | The rise in boiling point of a solvent caused by a non-volatile solute: $\Delta T_{b}=K_{b}\,m$ (times $i$ for electrolytes). | 1.3 |
| Henry's law | At constant temperature, the solubility of a gas in a liquid is proportional to its partial pressure: $p=K_{H}\,x$. A larger $K_{H}$ means a less soluble gas. | 1.1 1.2 |
| ideal solution | A solution that obeys Raoult's law at every composition, with $\Delta_{mix}H=0$ and $\Delta_{mix}V=0$ (e.g. hexane + heptane). | 1.2 1.3 |
| maximum-boiling azeotrope | An azeotrope from a large negative deviation; it boils above both pure components (nitric acid + water, about 68% acid by mass). | 1.2 |
| minimum-boiling azeotrope | An azeotrope from a large positive deviation; it boils below both pure components (ethanol + water, 95.6% ethanol by mass). | 1.2 |
| molality | Moles of solute per kilogram of solvent, unit $\mathrm{mol\,kg^{-1}}$ (m); it does not change with temperature. | 1.1 1.2 1.3 |
| molarity | Moles of solute per litre of solution, unit $\mathrm{mol\,L^{-1}}$ (M); it changes slightly with temperature. | 1.1 1.3 2.3 |
| mole fraction | Moles of one component divided by the total moles of all components, $x_{i}=n_{i}/\sum n$; the mole fractions of a solution add up to 1. | 1.1 1.2 1.3 |
| negative deviation | Vapour pressure below the Raoult's-law value, because A–B attractions are stronger than A–A and B–B; $\Delta_{mix}H\lt0$ and $\Delta_{mix}V\lt0$. | 1.2 |
| non-ideal solution | A solution whose vapour pressure is higher (positive deviation) or lower (negative deviation) than Raoult's law predicts. | 1.2 |
| osmosis | Net flow of solvent through a semipermeable membrane (one that passes solvent but not solute) from pure solvent, or a dilute solution, into a more concentrated solution. | 1.3 2.3 |
| osmotic pressure | The extra pressure that must be applied to a solution to stop solvent flowing into it through a semipermeable membrane: $\pi=CRT$ (times $i$ for electrolytes). Solutions with equal $\pi$ are isotonic. | 1.3 |
| parts per million | Parts of a component per $10^{6}$ parts of the solution (by mass, volume or moles); used for traces. In dilute water, 1 ppm by mass ≈ 1 mg per litre. | 1.1 1.2 5.1 6.3 |
| positive deviation | Vapour pressure above the Raoult's-law value, because A–B attractions are weaker than A–A and B–B; usually $\Delta_{mix}H\gt0$ and $\Delta_{mix}V\gt0$. | 1.2 |
| Raoult's law | For a solution, the partial vapour pressure of each volatile component is its mole fraction times the vapour pressure of the pure component: $p_{i}=x_{i}\,p_{i}^{\circ}$. | 1.2 |
| relative lowering of vapour pressure | $(p_{1}^{\circ}-p_{1})/p_{1}^{\circ}$; for a dilute solution of a non-volatile non-electrolyte it equals the mole fraction of the solute, $x_{2}$. | 1.3 |
| reverse osmosis | Flow of pure solvent out of a solution through a semipermeable membrane, forced by applying a pressure greater than the osmotic pressure; used to desalinate seawater. | 1.3 |
| saturated solution | A solution holding as much solute as it can at that temperature (and pressure), in dynamic equilibrium with undissolved solute. | 1.1 10.1 |
| solute | A component of a solution present in the smaller amount, dissolved in the solvent (e.g. salt in brine). | 1.1 1.2 1.3 2.1 |
| solvent | The component of a solution present in the largest amount; it usually decides the physical state of the solution. | 1.1 1.2 1.3 2.3 |
| van 't Hoff factor | $i$ = normal molar mass ÷ abnormal molar mass = observed ÷ calculated colligative property; $i\gt1$ for dissociation, $i\lt1$ for association. | 1.3 8.3 |
| vapour pressure | The pressure of a vapour in dynamic equilibrium with its liquid in a closed vessel at a given temperature; it depends only on the liquid and the temperature. | 1.2 1.3 |
Ch 2 · Electrochemistry
| Term | Meaning | Used in |
|---|---|---|
| anode | The electrode at which oxidation occurs. It is the negative terminal of a galvanic cell and is joined to the positive terminal of the supply in electrolysis. | 2.1 2.2 2.4 4.1 |
| cathode | The electrode at which reduction occurs. It is the positive terminal of a galvanic cell and is joined to the negative terminal of the supply in electrolysis. | 2.1 2.2 2.4 4.1 |
| cell constant | $G^{*}=l/A$ for a conductivity cell, found as $R\kappa$ with a standard KCl solution; unit cm⁻¹ or m⁻¹. | 2.3 |
| concentration cell | A galvanic cell with the same couple on both sides at different concentrations; $E^{\circ}_{cell}=0$ and the more dilute electrode is the anode. | 2.2 |
| conductance | The reciprocal of resistance, $G=1/R$, measured in siemens (S = Ω⁻¹). | 2.3 |
| conductivity | The reciprocal of resistivity, $\kappa=1/\rho$: the conductance of a 1 cm cube of the solution; unit S m⁻¹ or S cm⁻¹. It falls on dilution. | 2.3 3.2 4.1 5.1 |
| conductivity cell | Two platinised platinum electrodes fixed a set distance apart, used with an AC Wheatstone bridge to measure the resistance of a solution. | 2.3 |
| corrosion | Slow destruction of a metal by its surroundings through electrochemical oxidation; for iron (rusting) the product is Fe₂O₃·xH₂O. | 2.1 2.4 4.1 |
| electrochemical cell | A device in which a redox reaction is linked to a flow of electrons through an external circuit: galvanic (makes electricity) or electrolytic (uses it). | 2.1 |
| electrochemical equivalent | $Z=M/nF$: the mass of a substance deposited or liberated by one coulomb of charge (g C⁻¹). | 2.4 |
| electrochemical series | Standard reduction potentials listed in order: strong oxidising agents at the top (high $E^{\circ}$), strong reducing agents at the bottom (low $E^{\circ}$). | 2.1 |
| electrode potential | The potential difference between an electrode and its electrolyte, quoted as a reduction potential: the tendency of the couple to gain electrons. | 2.1 2.2 4.1 |
| electrolytic cell | A cell in which an external DC supply drives a non-spontaneous redox reaction (electrolysis); its cathode is negative and its anode positive. | 2.1 2.4 |
| electromotive force | The potential difference between the electrodes of a galvanic cell when no current is drawn; standard value $E^{\circ}_{cell}=E^{\circ}_{R}-E^{\circ}_{L}$. | 2.1 2.2 |
| Faraday constant | The charge on one mole of electrons, $F=96\,487\ \mathrm{C\,mol^{-1}}$ (often rounded to 96 500). | 2.2 |
| Faraday's laws of electrolysis | (1) The mass liberated is proportional to the charge passed, $m=ZQ=ZIt$. (2) For the same charge, masses are proportional to equivalent masses $M/n$. Together: $m=MIt/nF$. | 2.4 |
| fuel cell | A galvanic cell that turns the energy of a fuel's combustion directly into electricity, with fuel and oxygen supplied continuously; the H₂–O₂ cell gives 1.23 V and only water. | 2.4 |
| galvanic cell | An electrochemical cell that turns the chemical energy of a spontaneous redox reaction into electrical energy, e.g. the Daniell cell (also called a voltaic cell). | 2.1 2.2 2.4 |
| hydrogen economy | An energy system in which hydrogen made by electrolysing water with renewable electricity is stored, transported and used as a clean fuel, e.g. in fuel cells. | 2.4 |
| intensive property | A property that does not depend on the amount of substance, like temperature or electrode potential; $\Delta_{r}G$ is extensive (it doubles when the equation is doubled), $E$ is not. | 2.1 |
| Kohlrausch's law | At infinite dilution each ion contributes independently: $\Lambda_{m}^{\circ}=\nu_{+}\lambda_{+}^{\circ}+\nu_{-}\lambda_{-}^{\circ}$. Used to find $\Lambda_{m}^{\circ}$, α and $K_{a}$ of weak electrolytes. | 2.3 |
| limiting molar conductivity | $\Lambda_{m}^{\circ}$, the molar conductivity at infinite dilution ($c\rightarrow0$); for strong electrolytes the intercept of $\Lambda_{m}$ against $\sqrt{c}$. | 2.3 |
| molar conductivity | $\Lambda_{m}=\kappa/c$, i.e. $\kappa\times1000/c$ in S cm² mol⁻¹ (κ in S cm⁻¹, c in mol L⁻¹). It increases on dilution. | 2.3 5.1 |
| Nernst equation | Gives an electrode or cell potential at any concentrations: $E_{cell}=E^{\circ}_{cell}-\tfrac{2.303RT}{nF}\log Q$, i.e. $-\tfrac{0.059}{n}\log Q$ volts at 298 K. | 2.1 2.2 2.4 |
| overpotential | The extra voltage, beyond the value from $E^{\circ}$, needed to make a slow electrode reaction (such as O₂ evolution) go at a useful rate; it explains why brine gives Cl₂, not O₂. | 2.4 |
| primary battery | A battery whose reaction occurs only once and which cannot be recharged, e.g. the dry (Leclanché) cell and the mercury cell. | 2.4 |
| reaction quotient | $Q$: products over reactants, each concentration raised to its coefficient (solids omitted, gases as pressures). It equals $K$ only at equilibrium. | 2.2 |
| resistivity | The resistance of a conductor of unit length and unit cross-section, $\rho$ in $R=\rho\,l/A$; unit Ω m or Ω cm. | 2.3 |
| sacrificial protection | Protecting iron by contact with a more reactive metal (Zn, Mg) of lower $E^{\circ}$, which becomes the anode and corrodes instead; galvanising is an example. | 2.1 2.4 |
| salt bridge | A tube of jelly containing an inert electrolyte (KCl, KNO₃) that joins the two half-cells, completes the circuit through ions and keeps both solutions electrically neutral. | 2.1 2.2 |
| secondary battery | A rechargeable battery: passing current in the reverse direction reverses its cell reaction, e.g. the lead storage battery and the nickel–cadmium cell. | 2.4 |
| standard electrode potential | The electrode potential $E^{\circ}$ when all species are in their standard states (1 mol L⁻¹, 1 bar, pure solids), measured against the standard hydrogen electrode. | 2.1 4.1 |
| standard hydrogen electrode | The reference electrode Pt | H₂(g, 1 bar) | H⁺(aq, 1 M), whose potential is defined as exactly 0.00 V at all temperatures (SHE). | 2.1 |
| strong electrolyte | An electrolyte completely ionised in solution (KCl, NaCl, HCl); its $\Lambda_{m}$ rises only gently on dilution, $\Lambda_{m}=\Lambda_{m}^{\circ}-A\sqrt{c}$. | 1.3 2.3 |
| weak electrolyte | An electrolyte only partly ionised in solution (ethanoic acid, NH₄OH); its $\Lambda_{m}$ rises steeply at low concentration, and $\alpha=\Lambda_{m}/\Lambda_{m}^{\circ}$. | 2.3 |
Ch 3 · Chemical Kinetics
| Term | Meaning | Used in |
|---|---|---|
| activated complex | The short-lived, highest-energy arrangement of atoms at the top of the energy profile between reactants and products (transition state). | 3.3 |
| activation energy | The minimum extra energy $E_{a}$ that reacting molecules need to reach the activated complex; the height of the energy hill measured from the reactants. | 3.1 3.3 4.1 10.2 |
| Arrhenius equation | $k=Ae^{-E_{a}/RT}$: the temperature dependence of the rate constant; $\ln k$ against $1/T$ is a straight line of slope $-E_{a}/R$. | 3.3 10.2 |
| average rate | The rate over a finite time interval, $-\Delta[\mathrm{R}]/\Delta t$: the slope of a chord on the concentration–time graph. | 3.1 |
| catalyst | A substance that increases the rate of a reaction by providing a pathway of lower activation energy, and is recovered unchanged; it does not change $\Delta_{r}G$ or $K$. | 2.4 3.1 3.2 3.3 |
| chemical kinetics | The branch of chemistry that studies the rates of reactions, the factors that change them and the mechanisms by which reactions occur. | 3.1 |
| collision frequency | The number of collisions between reacting molecules per second per unit volume, $Z_{\mathrm{AB}}$; proportional to $[\mathrm{A}][\mathrm{B}]$. | 3.3 |
| effective collision | A collision that leads to reaction: the molecules have at least the threshold energy and a suitable orientation. | 3.3 |
| elementary reaction | A reaction that occurs in a single step exactly as written; its rate law can be written from its equation. | 3.1 6.2 |
| first-order reaction | A reaction whose rate is proportional to one concentration: $\ln[\mathrm{R}]=\ln[\mathrm{R}]_{0}-kt$, constant half-life $0.693/k$ (e.g. radioactive decay). | 3.1 3.2 3.3 10.1 |
| half-life | The time for the concentration (or amount) of a reactant to fall to half its value; $t_{1/2}=0.693/k$ for first order. | 3.2 4.3 6.2 |
| instantaneous rate | The rate at one moment, $-d[\mathrm{R}]/dt$: the slope of the tangent to the concentration–time curve at that moment. | 3.1 |
| integrated rate equation | An equation, obtained by integrating a rate law, that gives the concentration of a reactant as a function of time, e.g. $[\mathrm{R}]=[\mathrm{R}]_{0}e^{-kt}$. | 3.2 |
| inversion of cane sugar | Acid-catalysed hydrolysis of sucrose into glucose and fructose; the optical rotation changes sign from + to −. A pseudo-first-order reaction. | 3.2 10.1 |
| Maxwell–Boltzmann distribution | The spread of kinetic energies among the molecules of a gas at a given temperature; its high-energy tail beyond $E_{a}$ holds the molecules able to react. | 3.3 |
| molecularity | The number of reacting species that collide simultaneously in one elementary step: 1, 2 or 3 (uni-, bi-, termolecular). | 3.1 6.2 |
| order of reaction | The sum of the powers of the concentrations in the rate law; found by experiment and may be zero or fractional. | 3.1 6.2 |
| pre-exponential factor | The constant $A$ in the Arrhenius equation, with the same units as $k$; related to the collision frequency and orientation of molecules. | 3.3 10.2 |
| pseudo-first-order reaction | A reaction with two or more reactants that follows first-order kinetics because the others are in large excess (e.g. acid hydrolysis of an ester in water). | 3.2 6.2 10.1 |
| radioactive decay | The spontaneous breakdown of unstable nuclei; it always follows first-order kinetics, with a half-life characteristic of the isotope. | 3.2 4.3 |
| rate constant | The constant $k$ in a rate law: the rate when all concentrations are $1\ \mathrm{mol\,L^{-1}}$. It depends on temperature; its units depend on the order. | 3.1 3.2 3.3 6.2 |
| rate law | The experimentally found expression $\text{rate}=k[\mathrm{A}]^{x}[\mathrm{B}]^{y}$ giving the rate in terms of concentrations. | 3.1 3.2 6.2 8.2 |
| rate of reaction | The change in concentration of a reactant or product per unit time, divided by its stoichiometric coefficient; usual unit $\mathrm{mol\,L^{-1}\,s^{-1}}$. | 3.1 3.3 6.2 |
| rate-determining step | The slowest step of a mechanism, which limits the rate of the overall reaction. | 3.1 6.2 7.2 |
| steric factor | The fraction $P$ of sufficiently energetic collisions that have a suitable orientation; rate $=P\,Z_{\mathrm{AB}}\,e^{-E_{a}/RT}$. | 3.3 |
| temperature coefficient | The ratio of rate constants at two temperatures 10 K apart, $k_{T+10}/k_{T}$; about 2–3 for many reactions near room temperature. | 3.3 |
| zero-order reaction | A reaction whose rate does not depend on reactant concentration: $[\mathrm{R}]=[\mathrm{R}]_{0}-kt$, $t_{1/2}=[\mathrm{R}]_{0}/2k$ (e.g. $\mathrm{NH_{3}}$ on hot Pt). | 3.1 3.2 |
Ch 4 · The d- and f-Block Elements
| Term | Meaning | Used in |
|---|---|---|
| actinoid | One of the fourteen elements Th–Lr ($Z$ = 90–103) that fill the $5f$ subshell (Ac is usually included); all radioactive, with oxidation states up to $+7$; symbol An. | 4.3 |
| actinoid contraction | The decrease in size of actinoid atoms and $\mathrm{M^{3+}}$ ions across the series; larger from element to element than the lanthanoid contraction because $5f$ electrons shield even more poorly. | 4.3 |
| amphoteric | Able to react both as an acid and as a base; e.g. $\mathrm{Cr_{2}O_{3}}$ dissolves in acids and in fused alkalis. | 4.2 10.2 |
| Bohr magneton | The unit of atomic magnetic moment: $1\ \mathrm{BM}=9.27\times10^{-24}\ \mathrm{A\,m^{2}}$. | 4.1 |
| charge-transfer transition | Absorption of light that moves an electron from a ligand (e.g. oxygen) to the metal; it gives intense colours even to $d^{0}$ species such as $\mathrm{MnO_{4}^{-}}$ and $\mathrm{Cr_{2}O_{7}^{2-}}$. | 4.2 |
| chromite | The chief ore of chromium, $\mathrm{FeCr_{2}O_{4}}$ ($\mathrm{FeO\cdot Cr_{2}O_{3}}$), the starting material for chromates and dichromates. | 4.2 |
| d-block element | An element of groups 3–12 in which the last electron enters an $(n-1)d$ orbital; outer configuration $(n-1)d^{1-10}ns^{1-2}$. | 4.1 |
| diamagnetic | Weakly repelled by a magnetic field; all electrons are paired (e.g. Zn²⁺, Sc³⁺). | 4.1 4.2 4.3 5.3 |
| disproportionation | A redox reaction in which one species is both oxidised and reduced, e.g. $\mathrm{3MnO_{4}^{2-}+4H^{+}\rightarrow2MnO_{4}^{-}+MnO_{2}+2H_{2}O}$. | 4.1 4.2 8.2 |
| d–d transition | The jump of an electron between $d$ orbitals that have been split in energy by ligands, absorbing visible light; it gives most transition-metal ions their colour (complementary to the light absorbed). | 5.3 |
| enthalpy of atomisation | $\Delta_{a}H^{\circ}$: the enthalpy change to turn one mole of an element into separate gaseous atoms; high for transition metals because $d$ electrons join in metallic bonding. | 4.1 |
| f-block element | An element in which the last electron enters an $(n-2)f$ orbital: the lanthanoids ($4f$) and actinoids ($5f$), placed in two rows below the main periodic table. | 4.3 |
| f–f transition | Excitation of an electron within the $4f$ (or $5f$) subshell; gives lanthanoid ions their colours, with very narrow absorption bands because the $f$ orbitals are well shielded. | |
| inner transition element | Another name for an f-block element: the subshell being filled ($4f$ or $5f$) lies inside the $(n-1)d$ and $ns$ levels. | 4.3 |
| interstitial compound | A compound formed when small atoms (H, C, N, B) occupy holes in a metal lattice, e.g. TiC, Mn₄N; usually non-stoichiometric, very hard, high-melting, metallically conducting and inert. | 4.1 |
| lanthanoid | One of the fourteen elements Ce–Lu ($Z$ = 58–71) that fill the $4f$ subshell (La is usually included); characteristic oxidation state $+3$; symbol Ln. | 4.1 4.3 |
| lanthanoid contraction | The steady decrease in atomic and ionic radii from La to Lu ($\mathrm{Ln^{3+}}$: 103 → 86 pm), due to poor shielding by $4f$ electrons; it makes the 4d and 5d metals of a group nearly the same size (Zr ≈ Hf). | 4.3 |
| mischmetall | An alloy of about 95% lanthanoid metals (mainly Ce, La), about 5% iron and traces of S, C, Ca and Al; with magnesium it is used in lighter flints, bullets and shells. | 4.3 |
| n-factor | The number of electrons gained or lost per formula unit in a redox reaction; for $\mathrm{MnO_{4}^{-}}$ it is 5 in acid, 3 in neutral and 1 in strongly alkaline solution. | 4.2 |
| oxoanion | A negative ion in which a central atom is bonded only to oxygen, e.g. $\mathrm{CrO_{4}^{2-}}$, $\mathrm{Cr_{2}O_{7}^{2-}}$, $\mathrm{MnO_{4}^{-}}$; transition metals reach their highest oxidation states in such ions. | 4.1 4.2 |
| paramagnetic | Attracted into a magnetic field because of one or more unpaired electrons. | 4.1 4.2 4.3 5.3 |
| primary standard | A pure, stable, non-hygroscopic substance that can be weighed out accurately to make a solution of exactly known concentration, e.g. $\mathrm{K_{2}Cr_{2}O_{7}}$, sodium oxalate. | 4.2 |
| self-indicator | A titrant whose own colour shows the end point; acidified $\mathrm{KMnO_{4}}$ is decolourised until the first excess drop leaves a permanent pink. | 4.2 |
| spin-only magnetic moment | $\mu=\sqrt{n(n+2)}\ \mathrm{BM}$ for $n$ unpaired electrons; a good estimate for 3d ions, which ignores the orbital contribution. | 4.1 4.3 5.3 |
| transition element | An element whose atom, or an ion in a common oxidation state, has an incompletely filled $d$ subshell ($d^{1}$–$d^{9}$). Zn, Cd and Hg are $d$-block elements but not transition elements. | 4.1 4.3 |
| variable oxidation state | The ability of an element to form stable compounds in two or more oxidation states; for transition metals these usually differ by one unit (e.g. Fe²⁺/Fe³⁺). | 4.1 4.3 |
Ch 5 · Coordination Compounds
| Term | Meaning | Used in |
|---|---|---|
| ambidentate ligand | A ligand with two different possible donor atoms that binds through only one at a time: $\mathrm{NO_{2}^{-}}$ (through N or O), $\mathrm{SCN^{-}}$ (through S or N). | 5.1 5.2 |
| chelate | A complex in which a di- or polydentate ligand binds the same metal through two or more donor atoms, forming a ring that contains the metal; such ligands are chelating ligands. Chelates are especially stable. | 5.1 5.2 5.3 |
| coordination compound | A compound containing a central metal atom or ion bonded to a fixed set of ions or molecules (ligands) that stays together in solution, e.g. $\mathrm{K_{4}[Fe(CN)_{6}]}$; unlike a double salt, it does not break up into all its simple ions. | 5.1 5.2 5.3 |
| coordination entity | The central atom or ion together with the ligands bonded to it, written in square brackets: $\mathrm{[Co(NH_{3})_{6}]^{3+}}$, $\mathrm{[Ni(CO)_{4}]}$. It may be a cation, an anion or neutral. | 5.1 |
| coordination isomerism | Structural isomerism in which ligands are exchanged between a complex cation and a complex anion of different metals: $\mathrm{[Co(NH_{3})_{6}][Cr(CN)_{6}]}$ and $\mathrm{[Cr(NH_{3})_{6}][Co(CN)_{6}]}$. | 5.2 |
| coordination number | The number of donor atoms bonded directly to the central atom (not the number of ligands): 6 in $\mathrm{[Co(en)_{3}]^{3+}}$, 4 in $\mathrm{[Ni(CO)_{4}]}$. | 5.1 5.2 5.3 |
| coordination sphere | The central atom and its ligands, written inside square brackets; ions outside the brackets are counter ions and ionise in solution. | 5.1 5.2 |
| crystal field splitting | The energy gap between the two sets of d orbitals in a complex: $\Delta_{o}$ between $e_{g}$ and $t_{2g}$ in an octahedral field, $\Delta_{t}\approx\tfrac49\Delta_{o}$ between $t_{2}$ and $e$ in a tetrahedral field. | 5.3 |
| crystal field stabilisation energy | The lowering of the d-electron energy in a crystal field relative to the average: $(-0.4x+0.6y)\Delta_{o}$ for $t_{2g}^{x}e_{g}^{y}$, plus $P$ for each extra pair when spin states are compared. | 5.3 |
| crystal field theory | A model of complexes that treats ligands as point charges or dipoles and the metal–ligand bond as purely electrostatic; the ligands' field splits the energies of the metal's d orbitals. | 5.3 |
| denticity | The number of donor atoms through which one ligand binds to the central atom: unidentate ($\mathrm{Cl^{-}}$), didentate (en, oxalate), hexadentate ($\mathrm{EDTA^{4-}}$). | 5.1 5.2 5.3 |
| fac / mer isomers | The two geometrical isomers of an octahedral $\mathrm{[MA_{3}B_{3}]}$ complex: facial (fac), with the three B ligands on one face of the octahedron, and meridional (mer), with them on a meridian, two of them trans. | 5.2 |
| geometrical isomerism | Stereoisomerism in which the same ligands occupy different relative positions around the central atom: cis (adjacent, 90°) or trans (opposite, 180°); fac or mer in $\mathrm{[MA_{3}B_{3}]}$. Found in square planar and octahedral, not tetrahedral, complexes. | 5.2 6.3 8.2 |
| high-spin complex | A complex with the maximum number of unpaired d electrons, formed when the splitting is smaller than the pairing energy ($\Delta\lt P$, weak-field ligands); e.g. $\mathrm{[Fe(H_{2}O)_{6}]^{2+}}$, $t_{2g}^{4}e_{g}^{2}$. | 5.3 |
| homoleptic / heteroleptic | A homoleptic complex has only one kind of ligand ($\mathrm{[Co(NH_{3})_{6}]^{3+}}$); a heteroleptic complex has more than one kind ($\mathrm{[CoCl(NH_{3})_{5}]^{2+}}$). | 5.1 5.3 |
| inner orbital complex | An octahedral complex whose metal uses inner $(n-1)d$ orbitals in $d^{2}sp^{3}$ hybridisation, after strong-field ligands have paired its d electrons; low-spin, e.g. $\mathrm{[Co(NH_{3})_{6}]^{3+}}$. | 5.3 |
| ionisation isomerism | Structural isomerism in which a counter ion and a ligand exchange places, so the isomers give different ions in solution: $\mathrm{[Co(NH_{3})_{5}(SO_{4})]Br}$ and $\mathrm{[CoBr(NH_{3})_{5}]SO_{4}}$. | 5.2 |
| ligand | An ion or molecule bonded to the central atom of a coordination entity through a donor atom that gives it a lone pair (a Lewis base), e.g. $\mathrm{NH_{3}}$, $\mathrm{H_{2}O}$, $\mathrm{Cl^{-}}$, $\mathrm{CN^{-}}$. | 4.1 4.3 5.1 5.2 |
| linkage isomerism | Structural isomerism in which an ambidentate ligand binds through different donor atoms: $\mathrm{[Co(NH_{3})_{5}(NO_{2})]^{2+}}$ (nitrito-N) and $\mathrm{[Co(NH_{3})_{5}(ONO)]^{2+}}$ (nitrito-O). | 5.2 |
| low-spin complex | A complex in which d electrons pair in the lower set before entering the upper set, because $\Delta\gt P$ (strong-field ligands); e.g. $\mathrm{[Fe(CN)_{6}]^{4-}}$, $t_{2g}^{6}$. | 5.3 |
| optical isomerism | Stereoisomerism shown by a chiral complex, one that cannot be superimposed on its mirror image; the two mirror-image forms rotate plane-polarised light equally in opposite directions ($d$ and $l$), e.g. $\mathrm{[Co(en)_{3}]^{3+}}$, cis-$\mathrm{[CoCl_{2}(en)_{2}]^{+}}$. | 5.2 9.1 |
| outer orbital complex | An octahedral complex whose metal uses outer $nd$ orbitals in $sp^{3}d^{2}$ hybridisation, its $(n-1)d$ electrons staying unpaired; high-spin, e.g. $\mathrm{[CoF_{6}]^{3-}}$. | 5.3 |
| pairing energy | $P$: the extra energy needed to put a second electron, of opposite spin, into an orbital that already holds one. | 5.3 |
| solvate isomerism | Structural isomerism in which solvent molecules (water: hydrate isomerism) are ligands in one isomer but free lattice molecules in another: $\mathrm{[Cr(H_{2}O)_{6}]Cl_{3}}$ and $\mathrm{[CrCl(H_{2}O)_{5}]Cl_{2}\cdot H_{2}O}$. | 5.2 |
| spectrochemical series | Ligands in order of increasing splitting $\Delta$: $\mathrm{I^{-}\lt Br^{-}\lt SCN^{-}\lt Cl^{-}\lt S^{2-}\lt F^{-}\lt OH^{-}\lt C_{2}O_{4}^{2-}\lt H_{2}O\lt NCS^{-}\lt EDTA^{4-}\lt NH_{3}\lt en\lt CN^{-}\lt CO}$. | 5.3 |
| structural isomerism | Isomerism in which the isomers differ in their bonds (which atoms are joined to which). In coordination compounds: linkage, coordination, ionisation and solvate isomerism. | 5.2 6.1 6.2 7.1 |
| synergic bonding | The mutually reinforcing bonding in metal carbonyls: σ donation of the C lone pair of CO to the metal, and π back-donation from filled metal d orbitals into the empty $\pi^{*}$ orbital of CO. | 5.3 |
| Werner's theory | The first theory of coordination compounds (1893): a metal shows a primary, ionisable valence (its oxidation state) and a secondary, non-ionisable valence (its coordination number); the secondary valences point in fixed directions, giving a definite shape. | 5.1 |
Ch 6 · Haloalkanes and Haloarenes
| Term | Meaning | Used in |
|---|---|---|
| allylic halide | A halide whose halogen is on an $sp^{3}$ carbon next to a $\mathrm{C{=}C}$ double bond, e.g. 3-chloroprop-1-ene $\mathrm{CH_{2}{=}CHCH_{2}Cl}$. | 6.1 6.2 |
| ambident nucleophile | A nucleophile with two different donor atoms, e.g. cyanide (C or N) and nitrite (O or N); KCN gives nitriles, AgCN isocyanides; KNO₂ gives alkyl nitrites, AgNO₂ nitroalkanes. | 6.2 |
| benzylic halide | A halide whose halogen is on an $sp^{3}$ carbon attached directly to an aromatic ring, e.g. (chloromethyl)benzene $\mathrm{C_{6}H_{5}CH_{2}Cl}$. | 6.1 6.2 |
| biomagnification | The increase in concentration of a persistent, fat-soluble substance (such as DDT) at each higher level of a food chain. | 6.3 |
| chirality | The property of an object or molecule of not being superimposable on its mirror image; a carbon carrying four different groups usually makes a molecule chiral. | 5.2 6.2 8.1 9.1 |
| chlorofluorocarbon | A compound of carbon, chlorine and fluorine (freon) such as $\mathrm{CCl_{2}F_{2}}$; once used as refrigerants and propellants, they release Cl atoms in the stratosphere that destroy ozone. | 6.1 6.3 |
| DDT | 1,1,1-Trichloro-2,2-bis(4-chlorophenyl)ethane, $\mathrm{C_{14}H_{9}Cl_{5}}$: a chlorinated insecticide that persists in the environment and builds up in food chains. | 6.1 6.3 |
| enantiomers | Stereoisomers that are non-superimposable mirror images; identical physical properties except that they rotate plane-polarised light by equal angles in opposite directions. | 5.2 6.2 8.1 8.2 |
| Finkelstein reaction | Halogen exchange making an alkyl iodide from an alkyl chloride or bromide with sodium iodide in dry acetone; the insoluble NaCl or NaBr precipitates and drives the reaction. | 6.1 6.2 |
| Fittig reaction | Coupling of two haloarene molecules with sodium in dry ether to give a biaryl, e.g. chlorobenzene to biphenyl. | 6.3 |
| geminal dihalide | A dihaloalkane with both halogen atoms on the same carbon, e.g. 1,1-dichloroethane $\mathrm{CH_{3}CHCl_{2}}$ (an alkylidene halide). | |
| Grignard reagent | An alkyl- or arylmagnesium halide RMgX, made from RX and Mg in dry ether; its carbon is carbanion-like, a strong base and nucleophile; water converts it to RH. | 6.3 7.1 7.3 8.1 |
| haloalkane | A compound in which halogen atoms replace hydrogen atoms of an aliphatic hydrocarbon, the halogen being on an $sp^{3}$ carbon (also called an alkyl halide), e.g. $\mathrm{CH_{3}CH_{2}Br}$. Classified as 1°, 2° or 3° by the number of carbons on the carbon bearing X. | 3.3 6.1 6.2 6.3 |
| haloarene | A compound with a halogen bonded directly to a carbon of an aromatic ring (also called an aryl halide), e.g. chlorobenzene $\mathrm{C_{6}H_{5}Cl}$; the C–X carbon is $sp^{2}$. | 3.3 6.1 6.3 7.1 |
| nucleophilic substitution | A reaction in which a nucleophile (electron-pair donor) replaces a leaving group such as a halide ion: $\mathrm{Nu^{-}+R{-}X\rightarrow R{-}Nu+X^{-}}$. | 6.2 6.3 7.3 8.2 |
| ozone depletion | Thinning of the stratospheric ozone layer, chiefly by chlorine and bromine atoms released from CFCs and halons in a catalytic chain reaction. | 6.3 |
| racemic mixture | An equimolar mixture of two enantiomers, written (±) or dl; optically inactive because the rotations cancel. | 5.2 6.2 8.2 |
| racemisation | Conversion of one enantiomer into a racemic mixture, e.g. by an SN1 reaction through a planar carbocation. | 6.2 |
| Sandmeyer reaction | Replacement of the diazonium group of an arenediazonium salt by Cl or Br using copper(I) chloride in HCl or copper(I) bromide in HBr, with loss of $\mathrm{N_{2}}$. | 6.1 9.3 |
| Saytzeff rule | In dehydrohalogenation the major alkene is the more substituted one, with more alkyl groups on the doubly bonded carbons (Zaitsev's rule). | 6.3 7.2 |
| SN1 | Unimolecular nucleophilic substitution: slow ionisation to a planar carbocation, then fast attack by the nucleophile; rate $=k[\mathrm{RX}]$, racemisation; reactivity $\mathrm{3^{\circ}\gt2^{\circ}\gt1^{\circ}\gt CH_{3}X}$. | 6.2 6.3 |
| SN2 | Bimolecular nucleophilic substitution: a one-step reaction with backside attack, a single transition state, rate $=k[\mathrm{RX}][\mathrm{Nu^{-}}]$ and inversion of configuration; reactivity $\mathrm{CH_{3}X\gt1^{\circ}\gt2^{\circ}\gt3^{\circ}}$. | 6.2 6.3 |
| stereocentre | An atom, usually an $sp^{3}$ carbon bonded to four different groups (asymmetric carbon, $\mathrm{C^{*}}$), at which swapping two groups gives a stereoisomer. | 6.2 9.1 10.1 10.2 |
| Swarts reaction | Preparation of an alkyl fluoride by heating an alkyl chloride or bromide with a metallic fluoride such as $\mathrm{AgF}$, $\mathrm{Hg_{2}F_{2}}$, $\mathrm{CoF_{2}}$ or $\mathrm{SbF_{3}}$. | 6.1 |
| vicinal dihalide | A dihaloalkane with the two halogen atoms on adjacent carbons, e.g. 1,2-dichloroethane $\mathrm{ClCH_{2}CH_{2}Cl}$; formed when halogens add to alkenes. | 6.1 |
| vinylic halide | A halide whose halogen is bonded to one of the $sp^{2}$ carbons of a $\mathrm{C{=}C}$ double bond, e.g. chloroethene $\mathrm{CH_{2}{=}CHCl}$; unreactive towards substitution. | 6.1 6.3 |
| Walden inversion | The inversion of configuration at carbon caused by backside attack in an SN2 reaction, like an umbrella turning inside out. | 6.2 |
| Wurtz reaction | Coupling of two haloalkanes with sodium in dry ether: $\mathrm{2RX+2Na\rightarrow R{-}R+2NaX}$; gives alkanes with twice the carbon count. | 6.3 |
| Wurtz–Fittig reaction | Coupling of a haloarene and a haloalkane with sodium in dry ether to give an alkylarene: $\mathrm{ArX+RX+2Na\rightarrow Ar{-}R+2NaX}$. | 6.3 |
| β-elimination | Loss of a hydrogen from a β-carbon and a halogen from the α-carbon of a haloalkane (dehydrohalogenation), usually with alcoholic KOH on heating, giving an alkene. | 6.3 |
Ch 7 · Alcohols, Phenols and Ethers
| Term | Meaning | Used in |
|---|---|---|
| acid-catalysed dehydration | Loss of water from an alcohol heated with conc. $\mathrm{H_{2}SO_{4}}$ or $\mathrm{H_{3}PO_{4}}$ to give an alkene, through a carbocation; ease 3° > 2° > 1°. | 7.2 7.3 8.3 |
| alcohol | A compound with one or more –OH groups, each bonded to an $sp^{3}$ (saturated) carbon atom, $\mathrm{R{-}OH}$, e.g. ethanol $\mathrm{CH_{3}CH_{2}OH}$. | 1.2 4.2 6.1 6.2 |
| alkoxide | The anion $\mathrm{RO^{-}}$ formed when an alcohol loses the proton of its –OH, e.g. sodium ethoxide $\mathrm{C_{2}H_{5}O^{-}Na^{+}}$ from ethanol and sodium; a strong base. | 6.3 7.2 7.3 8.2 |
| alkoxy group | The group $\mathrm{RO{-}}$ (methoxy $\mathrm{CH_{3}O{-}}$, ethoxy $\mathrm{C_{2}H_{5}O{-}}$); ethers are named as alkoxy-substituted hydrocarbons (alkoxyalkanes). | 7.1 |
| allylic alcohol | An alcohol with –OH on an $sp^{3}$ carbon next to a $\mathrm{C{=}C}$ double bond, e.g. prop-2-en-1-ol. | 7.1 |
| anisole | Methoxybenzene, $\mathrm{C_{6}H_{5}OCH_{3}}$: an aromatic ether whose –OCH₃ group activates the ring and directs electrophiles ortho and para; HI cleaves it to phenol and iodomethane. | 7.1 7.3 |
| benzylic alcohol | An alcohol with –OH on an $sp^{3}$ carbon attached directly to an aromatic ring, e.g. phenylmethanol $\mathrm{C_{6}H_{5}CH_{2}OH}$ (not a phenol). | 7.1 7.2 |
| cumene process | Industrial route to phenol: cumene (isopropylbenzene) is oxidised by air to cumene hydroperoxide, which dilute acid splits into phenol and propanone. | 7.1 |
| denatured alcohol | Commercial ethanol made unfit for drinking by adding copper sulphate (colour) and pyridine (foul smell). | 7.2 |
| epoxide | A cyclic ether with a strained three-membered C–C–O ring (e.g. oxirane, $\mathrm{C_{2}H_{4}O}$); far more reactive than ordinary ethers. | 7.3 |
| ether | A compound in which an oxygen atom joins two alkyl or aryl groups, $\mathrm{R{-}O{-}R'}$; named as an alkoxyalkane, e.g. ethoxyethane. | 6.1 6.2 6.3 7.1 |
| fermentation | Breakdown of sugars by enzymes of yeast, in the absence of air: $\mathrm{C_{6}H_{12}O_{6}\rightarrow2C_{2}H_{5}OH+2CO_{2}}$ (zymase); the commercial source of ethanol. | 7.1 7.2 |
| hydroboration–oxidation | Addition of diborane to an alkene followed by oxidation with $\mathrm{H_{2}O_{2}/OH^{-}}$: adds water anti-Markovnikov, so propene gives propan-1-ol. | 7.1 |
| Kolbe's reaction | Sodium phenoxide heated with $\mathrm{CO_{2}}$ under pressure, then acidified, gives 2-hydroxybenzoic acid (salicylic acid). | 7.2 |
| Lucas test | Test with anhydrous $\mathrm{ZnCl_{2}}$ in conc. $\mathrm{HCl}$: 3° alcohols turn it cloudy at once, 2° within about five minutes, 1° not at room temperature. | 7.2 |
| MTBE | Methyl tert-butyl ether, 2-methoxy-2-methylpropane $\mathrm{(CH_{3})_{3}COCH_{3}}$: a former petrol additive, withdrawn in many places because it contaminates groundwater. | 7.1 7.3 |
| phenol | A compound with an –OH group bonded directly to a carbon of an aromatic ring, $\mathrm{Ar{-}OH}$; also the name of the simplest one, $\mathrm{C_{6}H_{5}OH}$. | 1.2 6.1 6.3 7.1 |
| phenoxide ion | The anion $\mathrm{C_{6}H_{5}O^{-}}$ formed when phenol loses a proton; its negative charge is delocalised over the oxygen and the ortho and para ring carbons. | 6.3 7.1 7.2 7.3 |
| picric acid | 2,4,6-Trinitrophenol, a yellow solid and a strong acid ($\mathrm{p}K_{a}$ about 0.5) because three nitro groups stabilise its phenoxide ion. | 6.3 7.1 7.2 |
| primary alcohol | An alcohol whose carbinol carbon (the one carrying –OH) is bonded to one other carbon, $\mathrm{RCH_{2}OH}$ (methanol is included). | 6.1 7.1 7.2 7.3 |
| pyridinium chlorochromate | PCC, $\mathrm{C_{5}H_{5}NH^{+}CrO_{3}Cl^{-}}$: a mild oxidising agent that converts primary alcohols into aldehydes without further oxidation to acids. | 7.2 8.1 |
| Reimer–Tiemann reaction | Phenol warmed with chloroform and aqueous NaOH, then acidified, gives 2-hydroxybenzaldehyde (salicylaldehyde). | 7.2 |
| secondary alcohol | An alcohol whose carbinol carbon is bonded to two other carbons, $\mathrm{R_{2}CHOH}$, e.g. propan-2-ol. | 6.1 7.1 7.2 7.3 |
| symmetrical ether | An ether whose two groups on oxygen are identical, e.g. ethoxyethane $\mathrm{C_{2}H_{5}OC_{2}H_{5}}$. | 7.1 7.3 |
| tertiary alcohol | An alcohol whose carbinol carbon is bonded to three other carbons, $\mathrm{R_{3}COH}$, e.g. 2-methylpropan-2-ol. | 6.1 7.1 7.2 7.3 |
| unsymmetrical ether | An ether whose two groups on oxygen are different, e.g. methoxyethane or methoxybenzene. | 7.1 7.3 |
| Williamson synthesis | Preparation of an ether by $\mathrm{S_{N}2}$ reaction of a sodium alkoxide or phenoxide with a methyl or primary alkyl halide: $\mathrm{RX+R'ONa\rightarrow ROR'+NaX}$. | 7.3 |
Ch 8 · Aldehydes, Ketones and Carboxylic Acids
| Term | Meaning | Used in |
|---|---|---|
| acyl chloride | A derivative $\mathrm{RCOCl}$ in which the OH of a carboxylic acid is replaced by Cl, made with $\mathrm{SOCl_{2}}$, $\mathrm{PCl_{5}}$ or $\mathrm{PCl_{3}}$; e.g. ethanoyl chloride, $\mathrm{CH_{3}COCl}$. | 7.2 8.1 8.3 9.2 |
| aldehyde | A carbonyl compound in which the carbonyl carbon carries at least one hydrogen atom: $\mathrm{R{-}CHO}$ (or $\mathrm{HCHO}$). IUPAC names end in -al (ethanal) or -carbaldehyde. | 7.1 7.2 8.1 8.2 |
| aldol condensation | Reaction of an aldehyde or ketone having an α-hydrogen in dilute alkali: the enolate of one molecule adds to the carbonyl group of another, giving a β-hydroxy aldehyde or ketone (an aldol), which loses water on heating to give an α,β-unsaturated carbonyl compound. | 8.1 8.2 |
| Cannizzaro reaction | Disproportionation of an aldehyde with no α-hydrogen in concentrated alkali: one molecule is reduced to an alcohol and another oxidised to a carboxylate, e.g. $\mathrm{2HCHO\rightarrow CH_{3}OH+HCOO^{-}}$. | 8.2 |
| carbonyl group | The group $\mathrm{{>}C{=}O}$: an $sp^{2}$ carbon doubly bonded to oxygen, planar and polar ($\mathrm{\overset{\delta+}{C}{=}\overset{\delta-}{O}}$). It is the functional group of aldehydes and ketones and part of the carboxyl group. | 8.1 8.2 8.3 9.1 |
| carboxyl group | The functional group $\mathrm{{-}COOH}$: a carbonyl group and a hydroxyl group on the same carbon. Resonance with the OH lone pair makes its carbon less electrophilic than that of an aldehyde or ketone. | 8.3 10.2 |
| carboxylate ion | The anion $\mathrm{RCOO^{-}}$ left when a carboxylic acid loses its proton. Its two resonance structures are equivalent, so the charge is shared by both oxygens and both C–O bonds are equal. | 8.3 |
| carboxylic acid | An organic acid containing the carboxyl group, $\mathrm{R{-}COOH}$. IUPAC names end in -oic acid (ethanoic acid) or -carboxylic acid; typical $\mathrm{p}K_{a}$ values are about 2–5. | 7.1 7.2 8.1 8.2 |
| Clemmensen reduction | Reduction of a carbonyl group to $\mathrm{CH_{2}}$ with zinc amalgam and concentrated hydrochloric acid, converting an aldehyde or ketone into a hydrocarbon. | 8.2 |
| cross aldol condensation | An aldol condensation between two different carbonyl compounds. If both have α-hydrogens, four products form; it is useful when one partner (such as benzaldehyde, in the Claisen–Schmidt reaction) has no α-hydrogen. | 8.2 |
| cyanohydrin | A compound with $\mathrm{OH}$ and $\mathrm{CN}$ on the same carbon, $\mathrm{{>}C(OH)CN}$, formed by adding $\mathrm{HCN}$ (with a trace of base) to an aldehyde or ketone. | 8.2 8.3 10.1 |
| decarboxylation | Loss of $\mathrm{CO_{2}}$ from a carboxylic acid or its salt; heating a sodium carboxylate with sodalime gives an alkane with one carbon fewer: $\mathrm{RCOONa+NaOH\rightarrow RH+Na_{2}CO_{3}}$. | 8.3 |
| DIBAL-H | Diisobutylaluminium hydride, $\mathrm{AlH(i{-}Bu)_{2}}$: a reducing agent that converts nitriles and esters into aldehydes (after hydrolysis) and leaves $\mathrm{C{=}C}$ bonds untouched. | 8.1 |
| esterification | Formation of an ester from a carboxylic acid and an alcohol in the presence of an acid catalyst: $\mathrm{RCOOH+R'OH\rightleftharpoons RCOOR'+H_{2}O}$. It is reversible; the C–OH bond of the acid breaks. | 7.2 8.3 |
| Etard reaction | Oxidation of a methyl group on a benzene ring to $\mathrm{CHO}$ by chromyl chloride ($\mathrm{CrO_{2}Cl_{2}}$) in $\mathrm{CS_{2}}$, through a chromium complex that is then hydrolysed: toluene → benzaldehyde. | 8.1 |
| fatty acid | A long, unbranched aliphatic carboxylic acid (about $\mathrm{C_{12}}$–$\mathrm{C_{18}}$), such as octadecanoic (stearic) acid, found as esters in fats and oils; its sodium salts are soaps. | 8.3 |
| Fehling's test | Heating with Fehling's solution (copper(II) sulphate mixed with alkaline sodium potassium tartrate): aliphatic aldehydes give a red-brown precipitate of $\mathrm{Cu_{2}O}$; aromatic aldehydes and ketones do not. | 8.2 8.3 10.1 |
| Friedel–Crafts acylation | Attachment of an acyl group $\mathrm{RCO{-}}$ to an aromatic ring by an acyl chloride and anhydrous $\mathrm{AlCl_{3}}$, giving an aryl ketone: benzene + $\mathrm{CH_{3}COCl}$ → acetophenone. | 6.3 8.1 8.2 |
| Gattermann–Koch reaction | Introduction of a $\mathrm{CHO}$ group into benzene (or an alkylbenzene) with carbon monoxide and hydrogen chloride in the presence of anhydrous $\mathrm{AlCl_{3}}$ or $\mathrm{CuCl}$. | 8.1 |
| Hell–Volhard–Zelinsky reaction | α-Halogenation of a carboxylic acid that has an α-hydrogen, by $\mathrm{Cl_{2}}$ or $\mathrm{Br_{2}}$ with a little red phosphorus, followed by hydrolysis: $\mathrm{RCH_{2}COOH\rightarrow RCH(X)COOH}$. | 8.3 |
| iodoform test | Warming with iodine and sodium hydroxide: compounds containing $\mathrm{CH_{3}CO{-}}$ (methyl ketones, ethanal) or $\mathrm{CH_{3}CH(OH){-}}$ give a yellow precipitate of iodoform, $\mathrm{CHI_{3}}$. | 8.2 |
| ketone | A carbonyl compound in which the carbonyl carbon is bonded to two carbon atoms: $\mathrm{R{-}CO{-}R'}$. IUPAC names end in -one (propanone, butan-2-one). | 7.1 7.2 8.1 8.2 |
| Kolbe electrolysis | Electrolysis of an aqueous sodium or potassium carboxylate: at the anode two alkyl groups join, giving $\mathrm{R{-}R}$ and $\mathrm{CO_{2}}$ (e.g. ethane from sodium ethanoate). | 8.3 |
| nucleophilic addition | The typical reaction of aldehydes and ketones: a nucleophile attacks the δ+ carbonyl carbon, the carbon becomes tetrahedral ($sp^{3}$) in an alkoxide intermediate, and protonation gives the product, e.g. a cyanohydrin or an alcohol. | 8.2 8.3 |
| ozonolysis | Cleavage of a $\mathrm{C{=}C}$ bond by ozone followed by $\mathrm{Zn/H_{2}O}$: each doubly bonded carbon becomes a carbonyl carbon, giving aldehydes and/or ketones. | 8.1 |
| Rosenmund reduction | Hydrogenation of an acyl chloride to an aldehyde over palladium poisoned with barium sulphate: $\mathrm{RCOCl+H_{2}\rightarrow RCHO+HCl}$. | 8.1 8.3 |
| saponification | Hydrolysis of an ester by aqueous alkali, giving the salt of the carboxylic acid and the alcohol: $\mathrm{RCOOR'+NaOH\rightarrow RCOONa+R'OH}$. Soap is made this way from fats. | 8.3 |
| Stephen reaction | Reduction of a nitrile by $\mathrm{SnCl_{2}/HCl}$ to an imine, followed by hydrolysis to an aldehyde: $\mathrm{RCN\rightarrow RCH{=}NH\rightarrow RCHO}$. | 8.1 |
| Tollens' test | Warming with ammoniacal silver nitrate, $\mathrm{[Ag(NH_{3})_{2}]^{+}}$: aldehydes (aliphatic and aromatic) are oxidised and deposit a silver mirror; ketones do not react. | 8.2 10.1 |
| Wolff–Kishner reduction | Reduction of a carbonyl group to $\mathrm{CH_{2}}$ by forming the hydrazone and heating it with $\mathrm{KOH}$ in ethylene glycol; nitrogen is evolved. Used for compounds stable to base. | 8.2 |
Ch 9 · Amines
| Term | Meaning | Used in |
|---|---|---|
| acetanilide | $N$-phenylethanamide, $\mathrm{C_{6}H_{5}NHCOCH_{3}}$, made by acetylating aniline; used to protect the amino group so that the ring is substituted only once, mainly at the para position. | 9.2 |
| acylation of amines | Replacement of an H on the nitrogen of a 1° or 2° amine by an acyl group $\mathrm{RCO{-}}$, using an acid chloride, anhydride or ester (with pyridine to remove HCl); the product is an amide. With benzoyl chloride it is called benzoylation. | 9.2 |
| amine | A derivative of ammonia in which one, two or three hydrogen atoms are replaced by alkyl or aryl groups; nitrogen keeps a lone pair, so amines are bases and nucleophiles. | 5.1 6.1 6.2 6.3 |
| ammonolysis | Cleavage of the C–X bond of a haloalkane by ammonia (ethanolic $\mathrm{NH_{3}}$, sealed tube, about 373 K), giving a mixture of 1°, 2°, 3° amines and the quaternary salt. | 9.1 |
| anilinium ion | $\mathrm{C_{6}H_{5}NH_{3}^{+}}$, protonated aniline; it has no lone pair, is deactivating and meta-directing, which is why direct nitration of aniline gives much 3-nitroaniline. | 9.1 9.2 |
| arylamine | An amine whose nitrogen is bonded directly to a carbon of an aromatic ring, e.g. aniline $\mathrm{C_{6}H_{5}NH_{2}}$; benzylamine, with N on a $\mathrm{CH_{2}}$, is not one. | 6.1 9.1 9.2 9.3 |
| azo dye | A coloured compound containing the azo group $\mathrm{{-}N{=}N{-}}$ joining two aromatic rings, e.g. orange 4-hydroxyazobenzene; the extended conjugation absorbs visible light. | 9.2 9.3 |
| Balz–Schiemann reaction | Preparation of a fluoroarene: the diazonium salt is precipitated as its fluoroborate with $\mathrm{HBF_{4}}$, and heating the dry salt gives $\mathrm{ArF+BF_{3}+N_{2}}$. | 9.3 |
| basicity of amines | The readiness of an amine's nitrogen lone pair to accept a proton, measured by $K_{b}$ or $\mathrm{p}K_{b}$ (smaller $\mathrm{p}K_{b}$ = stronger base); set by the inductive effect, solvation, steric hindrance and, in arylamines, resonance. | 9.2 |
| carbylamine reaction | A primary amine heated with chloroform and ethanolic KOH gives a foul-smelling isocyanide: $\mathrm{RNH_{2}+CHCl_{3}+3KOH\rightarrow RNC+3KCl+3H_{2}O}$; a test for 1° amines. | 9.2 |
| coupling reaction | Electrophilic substitution of a phenol (mildly alkaline) or an arylamine (mildly acidic), usually at the para position, by an arenediazonium ion; both nitrogens are kept and an azo compound $\mathrm{Ar{-}N{=}N{-}Ar'}$ forms. | 9.3 |
| diazonium salt | A salt $\mathrm{R{-}N_{2}^{+}X^{-}}$ containing the $\mathrm{{-}N_{2}^{+}}$ group; arenediazonium salts such as benzenediazonium chloride are stable for a while in cold solution because the positive charge is delocalised over the ring. | 6.1 7.1 9.2 9.3 |
| diazotisation | Conversion of a primary aromatic amine into its diazonium salt with nitrous acid ($\mathrm{NaNO_{2}}$ + HCl) at 273–278 K: $\mathrm{ArNH_{2}+NaNO_{2}+2HCl\rightarrow ArN_{2}^{+}Cl^{-}+NaCl+2H_{2}O}$. | 6.1 7.1 9.3 |
| Gabriel phthalimide synthesis | Potassium phthalimide + haloalkane → $N$-alkylphthalimide, then alkaline hydrolysis → a pure primary amine; it cannot make arylamines. | 9.1 |
| Gattermann reaction | Replacement of the diazonium group by Cl or Br using copper powder with HCl or HBr; similar to the Sandmeyer reaction but usually in lower yield. | 9.3 |
| Hinsberg's reagent | Benzenesulphonyl chloride, $\mathrm{C_{6}H_{5}SO_{2}Cl}$: with a 1° amine it gives a sulphonamide soluble in alkali, with a 2° amine one insoluble in alkali, and with a 3° amine no reaction. | 9.2 |
| Hoffmann bromamide degradation | $\mathrm{RCONH_{2}+Br_{2}+4NaOH\rightarrow RNH_{2}+Na_{2}CO_{3}+2NaBr+2H_{2}O}$: a primary amide gives a primary amine with one carbon fewer, R migrating from C to N. | 9.1 |
| isocyanide | A compound $\mathrm{R{-}N{\equiv}C}$ (carbylamine), isomeric with the nitrile $\mathrm{R{-}C{\equiv}N}$; toxic, with an extremely unpleasant smell. | 6.2 9.2 |
| phosphinic acid | $\mathrm{H_{3}PO_{2}}$ (hypophosphorous acid), a reducing agent that replaces the diazonium group by hydrogen, removing an amino group from a ring after it has served as a directing group. | 9.3 |
| pKb | $\mathrm{p}K_{b}=-\log K_{b}$, where $K_{b}=[\mathrm{RNH_{3}^{+}}][\mathrm{OH^{-}}]/[\mathrm{RNH_{2}}]$; for an amine and its conjugate acid $\mathrm{p}K_{a}+\mathrm{p}K_{b}=14.00$ at 298 K. | |
| primary amine | An amine with one carbon group on nitrogen, $\mathrm{R{-}NH_{2}}$, e.g. ethanamine; the label counts the groups on N, not the type of carbon. | 6.2 8.2 9.1 9.2 |
| quaternary ammonium salt | An ionic compound $\mathrm{R_{4}N^{+}X^{-}}$ in which nitrogen carries four carbon groups and a positive charge and has no lone pair, e.g. $\mathrm{(CH_{3})_{4}N^{+}Cl^{-}}$. | 6.2 9.1 9.2 |
| secondary amine | An amine with two carbon groups on nitrogen, $\mathrm{R_{2}NH}$ or $\mathrm{RR'NH}$, e.g. $N$-methylethanamine; it has one N–H. | 9.1 9.2 |
| sulphanilic acid | 4-Aminobenzenesulphonic acid, made by heating anilinium hydrogensulphate at 453–473 K; it exists as the zwitterion $\mathrm{{}^{+}H_{3}NC_{6}H_{4}SO_{3}^{-}}$. | 9.2 9.3 |
| tertiary amine | An amine with three carbon groups on nitrogen, $\mathrm{R_{3}N}$, e.g. trimethylamine; it has no N–H, so it cannot hydrogen-bond to itself. | 6.2 9.1 9.2 |
Ch 10 · Biomolecules
| Term | Meaning | Used in |
|---|---|---|
| aldose / ketose | A monosaccharide containing an aldehyde group (aldose, e.g. glucose, an aldohexose) or a keto group (ketose, e.g. fructose, a ketohexose). | 10.1 |
| amino acid | A compound containing an amino group and a carboxyl group; proteins are built from about twenty α-amino acids, $\mathrm{H_{2}N{-}CH(R){-}COOH}$, which differ in the side chain R. | 8.3 9.1 9.2 10.2 |
| anomers | Ring forms of a sugar that differ only in configuration at the anomeric carbon (C1 of an aldose, C2 of a ketose), the carbon of the former C=O group: α- and β-D-glucopyranose. | 10.1 |
| carbohydrate | An optically active polyhydroxy aldehyde or polyhydroxy ketone, or a compound that gives these on hydrolysis: sugars, starch, cellulose, glycogen. | 10.1 10.3 |
| complementary base pairing | The specific hydrogen bonding of bases between the two strands of DNA: adenine with thymine (two H-bonds) and guanine with cytosine (three), so each strand fixes the sequence of the other. | 10.3 |
| denaturation | Loss of the biological activity of a protein when heat or a change of pH disrupts the hydrogen bonds and other weak interactions, destroying the secondary and tertiary structure while the primary structure remains; e.g. boiling egg white. | 7.1 10.2 |
| disaccharide | An oligosaccharide that gives two monosaccharide units on hydrolysis, $\mathrm{C_{12}H_{22}O_{11}}$ for two hexoses: sucrose (glucose + fructose), maltose (2 glucose), lactose (galactose + glucose). | 10.1 |
| DNA | Deoxyribonucleic acid: two complementary polynucleotide strands (sugar 2-deoxyribose; bases A, G, C, T) coiled into a double helix; it stores hereditary information and replicates. | 5.2 6.2 10.2 10.3 |
| enzyme | A biological catalyst, almost always a globular protein, that speeds up a specific reaction of its substrate under mild conditions by lowering the activation energy; named with the ending -ase (maltase, invertase). | 3.3 5.3 6.2 7.2 |
| essential amino acids | Amino acids that the human body cannot make and that must be supplied in the diet (valine, leucine, isoleucine, arginine, lysine, threonine, methionine, phenylalanine, tryptophan, histidine); those the body can make are non-essential. | 10.2 |
| fat-soluble / water-soluble vitamins | Vitamins A, D, E and K dissolve in fats and are stored in the liver and fat tissue; the B group and C dissolve in water, are excreted in urine and (except B₁₂) must be supplied regularly. | 10.3 |
| glycosidic linkage | The C–O–C link that joins two monosaccharide units, formed by loss of water between the anomeric OH of one unit and an OH of the other; e.g. α(1→4) in maltose. | 10.1 |
| hormone | A chemical messenger made by an endocrine gland, carried by the blood to target organs where it acts in tiny amounts; steroids (testosterone), polypeptides (insulin) or amino acid derivatives (thyroxine). | 6.1 10.3 |
| invert sugar | The equimolar, laevorotatory mixture of D-glucose and D-fructose formed by hydrolysis of dextrorotatory sucrose; the sign of rotation is inverted. | 10.1 |
| isoelectric point | The pH at which an amino acid (or protein) has zero net charge, existing mainly as the zwitterion, so that it does not move in an electric field; about 6 for glycine. | 10.2 |
| monosaccharide | A carbohydrate that cannot be hydrolysed into simpler polyhydroxy aldehydes or ketones, e.g. glucose, fructose, ribose, galactose. | 10.1 |
| nucleic acid | A polymer of nucleotides (a polynucleotide) that stores or transmits genetic information: DNA and RNA. | 10.1 10.3 |
| nucleoside | A nitrogenous base joined to C1′ of a pentose sugar (ribose or 2-deoxyribose), e.g. adenosine; base + sugar − H₂O. | 10.3 |
| nucleotide | A nucleoside carrying a phosphate group on the sugar, usually at C5′ (base + sugar + phosphate); the monomer unit of nucleic acids. | 10.3 |
| oligosaccharide | A carbohydrate that gives two to ten monosaccharide units on hydrolysis (disaccharides, trisaccharides, …). | 10.1 |
| peptide bond | The amide link $\mathrm{{-}CO{-}NH{-}}$ formed between the carboxyl group of one amino acid and the amino group of another with loss of water; planar and rigid. | 10.2 |
| phosphodiester linkage | The phosphate group that joins C3′ of one nucleotide's sugar to C5′ of the next, forming the sugar–phosphate backbone of a nucleic acid. | 10.3 |
| polypeptide | A chain of more than ten amino acid residues joined by peptide bonds; one with more than about a hundred residues (molar mass above about 10 000 u) is called a protein. | 10.2 10.3 |
| polysaccharide | A carbohydrate that gives a large number of monosaccharide units on hydrolysis; not sweet (a non-sugar): starch, cellulose, glycogen. | 10.1 |
| primary structure | The sequence in which amino acids are joined in each polypeptide chain of a protein, held by peptide bonds. | 10.2 |
| reducing sugar | A carbohydrate that reduces Fehling's solution and Tollens' reagent because it has a free anomeric carbon whose ring can open to an aldehyde or keto group: all monosaccharides, maltose, lactose. Sucrose is a non-reducing sugar. | 10.1 |
| RNA | Ribonucleic acid: usually a single polynucleotide strand (sugar ribose; bases A, G, C, U). Messenger, ribosomal and transfer RNA work together in protein synthesis. | 10.3 |
| secondary structure | The local shape of a polypeptide chain held by hydrogen bonds between C=O and N–H groups of peptide bonds: the α-helix (a right-handed coil) or the β-pleated sheet (extended chains side by side). | 10.2 |
| vitamin | An organic compound needed in small amounts in the diet for normal growth and health, which the body cannot make in sufficient amount; its long-term lack causes a deficiency disease (e.g. vitamin C: scurvy). | 5.1 5.3 10.3 |
| zwitterion | The dipolar ion $\mathrm{H_{3}\overset{+}{N}{-}CH(R){-}COO^{-}}$ formed when the carboxyl group of an amino acid passes a proton to its own amino group; neutral overall and amphoteric. | 9.2 10.2 |