P.A. Cox - Inorganic chemistry (793955), страница 32
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halides, cyanide, oxalate) or neutral molecules (e.g. ammonia, pyridine). The ligandacts as a donor and replaces one or more water molecules from the primary solvation sphere, and thus a complex isdistinct from an ion pair, which forms through purely electrostatic interactions in solvents of low polarity (seeTopic E1). Although complex formation is especially characteristic of transition metal ions it is by no means confined tothem.Several steps of complex formation may be possible, and the successive equilibrium constants for the reactionsand so on are known as the stepwise formation constants K1, K2….
The overall equilibrium constant for thereaction138SECTION E—CHEMISTRY IN SOLUTIONis given bySuccessive stepwise formation constants often decrease regularly K1>K2>…of the maximum value being determined bythe number of ligands that can be accommodated: this is often six except for chelating ligands (see below). The decreasecan be understood on entropic (statistical) grounds, as each successive ligand has one less place available to attach.Exceptional effects may result from the charge and size of ligands, and a reversal of the normal sequence can sometimesbe attributed to specific electronic or structural effects. It is important to remember that each ligand replaces one ormore solvating water molecules. For example, in the Cd2+/Br− system K4>K3 as the octahedral species [Cd(H2O)3Br3]−is converted to tetrahedral [CdBr4]2− with an entropy gain resulting from the increased freedom of three watermolecules.Hard and soft behaviorFor cations formed from metals in early groups in the periodic table the complexing strength with halide ions followsthe sequencewhereas with some later transition metals and many post-transition metals the reverse sequence is found (e.g.
Pt2+, Hg2+, Pb2+; see Topics G4, G6 and H5). The former behavior is known as class a and the latter as class b behavior, andthe difference is an example of hard and soft properties, respectively (see Topic C9). Class b ions form strongcomplexes with ligands such as ammonia, which are softer than water, whereas class a ions do not complex with suchligands appreciably in water.Solvation plays an essential part in understanding the factors behind class a and b behavior.
Trends in bond strengthsshow that almost every ion would follow the class a sequence in the gas phase, and the behavior in water is a partly aconsequence of the weaker solvation of larger anions. With a class b ion such as Hg2+ the bond strengths decrease moreslowly in the sequence Hg-F>Hg-Cl>… than do the solvation energies of the halide ions. With a class a ion such as Al3+, on the other hand, the change in bond strengths is more marked than that in the solvation energies.In solution the difference between the two classes is often manifested in different thermodynamic behavior.
Class bcomplex formation is enthalpy dominated (i.e. driven by a negative ΔH) whereas class a formation is often entropydominated (driven by positive ΔS). The strongest class a ions are small and highly charged (e.g. Be2+, Al3+) and havevery negative entropies of solvation (see Topic E1). Complexing with small highly charged ions such as F− reduces theoverall charge and hence frees up water molecules, which are otherwise ordered by solvation. Hard cations with lowcharge/size ratio, such as alkali ions, form very weak complexes with all ligands except macrocycles (see below).Some polyatomic ions such asandhave very low complexing power to either class a or b metals.They are useful as counterions for studying the thermodynamic properties of metal ions (e.g.
electrode potentials; seeTopic E5) unaffected by complex formation.Chelates and macrocyclesChelating ligands are ones with two or more donor atoms capable of attaching simultaneously to a cation: they aredescribed as bidentate, tridentate,…according to the number of atoms capable of binding. Chelating ligands includeE3—COMPLED FORMATION139bidentate ethylenediamine (1) and ethylenediamine tetraacetate (EDTA, 2), which is hexadentate, having two nitrogendonors and four oxygens (one from each carboxylate). Chelating ligands generally form stronger complexes thanunidentate ones with similar donor properties. They are useful for analysis of metal ions by complexometrictitration and for removing toxic metals in cases of poisoning (see Topic J3).The origin of the chelate effect is entropic.
Each ligand molecule can replace more than one solvating watermolecule, thus giving a favorable entropy increase. Structural requirements occasionally subvert the effect: forexample, Ag+ does not show the expected increase of K1 with ethylenediamine compared with ammonia, because it hasa strong bonding preference for two ligand atoms in a linear configuration, which is structurally impossible with thebidentate ligand.The length of the chain formed between ligand atoms is important in chelate formation, the most stable complexesgenerally being formed with four atoms (including the donors) so that with the metal ion a five-membered ring isformed. Smaller ring sizes are less favorable because of the bond angles involved, larger ones because of the increasedconfigurational entropy of the molecule (coming from free rotation about bonds), which is lost in forming the complex.Limiting the possibility of bond rotation increases the complexing power even with optimum ring sizes, so thatphenanthroline (3) forms stronger complexes than bipyridyl (4).Reducing the configurational entropy is important in macrocyclic ligands, where several donor atoms are already‘tied’ by a molecular framework into the optimal positions for complex formation.
Examples are the cyclic crownethers (e.g. 18-crown-6, 5) and bicyclic cryptands (e.g. [2.2.1]-crypt, 6). As expected, complexing strength isenhanced, and the resulting macrocyclic effect allows complexes to be formed with ions such as those of group 1,which otherwise have very low complexing power (see Topic G2). Another feature of macrocyclic ligands is the sizeselectivity corresponding to different cavity sizes. Thus with ligand 6 complex stability follows the order Li+<Na+>K+>Rb+ and the selectivity can be altered by varying the ring size.Chelating and macrocyclic effects are important in biological chemistry (see Topic J3).
Metal binding sites inmetalloproteins contain several ligand atoms, with appropriate electronegativities, and arranged in a suitablegeometrical arrangement, to optimize the binding of a specific metal ion.140SECTION E—CHEMISTRY IN SOLUTIONEffect of pHpH changes will affect complex formation whenever any of the species involved has Brønsted acidity or basicity (seeTopic E2). Most good ligands (except Cl−, Br− and I−) are basic, and protonation at low pH will compete with complexformation. This is important in analytical applications. For example, in titrations with EDTA, Fe3+ (for which K1 isaround 1025) can be titrated at a pH down to two, but with Ca2+ (where K1 is about 1010) a pH of at least seven isrequired because at lower pH values complex formation is incomplete.Section E—Chemistry in solutionE4SOLUBILITY OF IONIC SUBSTANCESKey NotesThermodynamicsMajor trends in waterInfluence of pH and complexingOther solventsRelated topicThe equilibrium constant for dissolving an ionic substance isknown as the solubility product.
It is related to a Gibbs freeenergy change that depends on a balance of lattice energyand solvation energies, together with an entropycontribution.Solids tend to be less soluble when ions are of similar sizeor when both are multiply charged. Covalent contributionsto the lattice energy reduce solubility.Solubility increases in acid conditions when the anion isderived from a weak acid, for example hydroxide, sulfideor carbonate. Amphoteric substances may dissolve again athigh pH. Complexing agents also increase solubility.Highly polar solvents show parallels with water.Compounds with multiply charged ions are often insolublein less polar ones, but different donor properties andpolarizability play a role.Lattice energies (D6)ThermodynamicsConsider an ionic solid that dissolves in water according to the equation:(1)The equilibrium constant for this reaction,is known as the solubility product of MnXm.
The form of this equilibrium is important in understanding effects suchas the influence of pH and complexing (see below) and also the common ion effect: it can be seen that adding one ofthe ions Mm+ or Xn− will shift Reaction 1 to the left and so reduce the solubility of the salt. Thus AgCl(s) is much lesssoluble in a solution containing 1 M Ag+ (e.g. from soluble AgNO3) than otherwise.Equilibrium constants in solution should correctly be written using activities not concentrations. The differencebetween these quantities is large in concentrated ionic solutions, and Ksp is quantitatively reliable as a guide to solubilities142SECTION E—CHEMISTRY IN SOLUTION(measured in concentration units) only for very dilute solutions.
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