P.A. Cox - Inorganic chemistry (793955), страница 33
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Nevertheless, a thermodynamic analysis of the factorsdetermining Ksp is useful for understanding general solubility trends. According to(see Topic B3), Ksp is related to the standard Gibbs free energy change of solution. Figure 1 shows a thermodynamiccycle that relates the overall ΔG to two separate steps: (i) the formation of gas-phase ions; (ii) their subsequentsolvation.
The enthalpy contributions involve a balance between the lattice energy of the solid and the solvationenthalpies of the ions (see Topics D6 and E1). In a solvent such as water with a very high dielectric constant thesecontributions almost cancel. Nevertheless, some of the solubility trends summarized below can be understood in termsof the changing balance between lattice energies (proportional to 1/(r++r−), where r+ and r− are radii of individualions) and the sum of the individual solvation enthalpies (each roughly proportional to 1/r).
For example, a small ion hasa large (negative) solvation energy, but when partnered by a large counterion cannot achieve an especially large latticeenergy. With ions of very different size, therefore, solvation is relatively favored and solubility tends to be larger than withions of similar size. Entropy terms are, however, also important. The first step in Fig.
1 involves an entropy increase,but solvation produces an ordering of solvent molecules and a negative ΔS contribution. Overall ΔS values fordissolving ions with multiple charges are usually negative, an effect that tends to lower the solubility as mentionedbelow.Major trends in waterThe aqueous solubility of ionic compounds is important in synthetic and analytical chemistry (see Topics B6, B7), and inthe formation of minerals by geochemical processes (see Topic J2). The most significant trends are as follows.(i) Soluble salts are more often found when ions are of very different size rather than similar size.
Thus in comparingsalts with different alkali metal cations, lithium compounds are the least soluble of the series with OH− and F−, butthe most soluble with larger cations such as Cl− or. This principle is often useful in preparative reactions andseparations. If it is desired to precipitate a large complex anion, a large cation such as tetrabutyl ammonium[(C4H9)4N]+ can be helpful.(ii) Salts where both ions have multiple charges are less likely to be soluble than ones with single charges.
Thuscarbonatesand sulfatesof the larger group 2 cations are insoluble. An important factor is thenegative solvation entropies of the ions.(iii) With ions of different charges, especially insoluble compounds result when the lower charged one is smaller (asthis gives a very large lattice energy). Thus with M3+ ions, fluorides and hydroxides are generally very insoluble,whereas heavier halides and nitrates are very soluble.(iv) Lower solubility results from covalent contributions to the lattice energy (see Topic D6). This happens especiallywith ions of less electropositive metals in combination with more polarizable cations. Late transition and posttransition elements often have insoluble sulfides (see Topic J2); insoluble halides (but not generally fluoride) alsooccur, for example with Ag+ and Pb2+.Influence of pH and complexingAny substance in solution that reacts with one of the ions formed in Reaction 1 will shift the equilibrium to the right andhence increase the solubility of the solid.
pH will therefore influence the solubility in a range where one of the ions hassignificant Brønsted acid or base properties (see Topic E2). The solubility of NaCl, for example, should not be affectedE4—SOLUBILITY OF IONIC SUBSTANCES143Fig. 1. Thermodynamic cycle for the solution of an ionic solid MX.by pH, but when the anion is the conjugate base of a weak acid solubility will increase at low pH.
Metal oxides andhydroxides dissolve in acid solution, and conversely such solids may be precipitated from a solution containing a metalion as the pH is increased. The solubility range depends on the Ksp value: for example, Fe(OH)3 is precipitated at muchlower pH than the more soluble Fe(OH)2. At high pH the acidity of the hydrated metal ion may come into play andamphoteric substances such as Al(OH)3 will dissolve in alkaline solution to give [Al(OH)4]−.Sometimes the conjugate acid of the anion is volatile, or decomposes to form a gas.
Thus action of an acid on a sulfidewill liberate H2S, and on a carbonate CO2 from the decomposition of carbonic acid.Any ligand that complexes with the metal ion will also increase solubility. AgCl dissolves in aqueous ammonia by theformation of [Ag(NH3)2]+.
Addition of Cl−, which initially decreases the solubility of AgCl through the common ioneffect (see above), will at high concentrations increase the solubility by forming [AgCl2]−. The solubility of someamphoteric oxides and hydroxides at high pH can be interpreted as a similar complexing effect, with OH− acting as theligand.Other solventsWith its combination of high dielectric constant, good donor ability and hydrogen bonding capability (which contributesto the solvation of anions) water is one of the best solvents for ionic substances. Two other solvents of comparable polarityare HF and H2SO4 (see Topic E1, Table 1).
Solubility trends for metal fluorides in HF show close parallels with those forhydroxides and oxides in water. Thus they follow the sequence MF>MF2>MF3, and for a given charge the solubilitytends to increase with the cation size.As the dielectric constant decreases, the solvation energies become less able to compensate for the lattice energy.This is especially true for solids with multiply charged ions, and such compounds are much less soluble in liquidammonia than in water. Detailed comparisons are complicated, however, by the occurrence of ion pairing, and by theincreased importance of other interactions.
Ammonia is a better donor than water for soft class b cations such as Ag+(see Topic E3) and compounds such as AgCl are much more soluble. The solubility trend AgI>AgBr >AgCl in ammoniais also the reverse of that found in water, reflecting another difference: ammonia has a larger polarizability than waterand so van der Waals’ forces are more important. They contribute significantly to the solvation of heavier anions such asI−. Some iodides such as LiI are soluble in solvents of low polarity, a fact that is sometimes wrongly used to suggest thatthe solids have appreciable covalent character.
In fact, LiI often dissolves as an ion pair, the donor solvent coordinatingLi+ and the van der Waals’ forces solvating I−.Section E—Chemistry in solutionE5ELECTRODE POTENTIALSKey NotesStandard potentialsDirection of reactionNonstandard conditionsDiagrammaticrepresentationsKinetic limitationsRelated topicsAn electrode potential is a measure of the thermodynamics of a redoxreaction. It may be expressed as the difference between two half-cellpotentials, which by convention are measured against a hydrogenelectrode. Tabulated values refer to standard conditions (ions at unitactivity).Comparison of two electrode potentials allows prediction of thefavorable direction of a redox reaction, and of its equilibrium constant.Only the differences between electrode potentials are significant;individual potentials have no meaning.The Nernst equation shows how electrode potentials vary withactivity (approximately equal to concentration).
Potentials may beinfluenced by pH and complexing.Latimer and Frost diagrams are different ways of representing theelectrode potentials for different oxidation states of an element. Frostdiagrams are useful for visual comparisons between elements, and forshowing which species are likely to disproportionate.Electrode potentials give no information about the rate of a redoxreaction. Reactions where covalent bonds are involved may be veryslow.Stability and reactivity (B3)3d series: aqueous ions (H3)Oxidation and reduction(B4)Standard potentialsIn an electrochemical cell a redox reaction occurs in two halves (see Topic B4). Electrons are liberated by theoxidation half reaction at one electrode and pass through an electrical circuit to another electrode where they are usedfor the reduction.
The cell potential E is the potential difference between the two electrodes required to balance thethermodynamic tendency for reaction, so that the cell is in equilibrium and no electrical current flows. E is related tothe molar Gibbs free energy change in the overall reaction (see Topic B3) according to(1)where F is the Faraday constant (9.6485×104 C mol−1) and n the number of moles of electrons passed per mole ofreaction.SECTION E—CHEMISTRY IN SOLUTION145Table 1.
Some standard electrode potentials in aqueous solution at pH=0 and 25°CIt is useful to think of the cell potential as the difference between the potentials associated with the two half-cellreactions, although these are not separately measurable. Standard electrode potentials are the half-cell potentialsmeasured against a hydrogen electrode, where the half-cell reaction isall reagents being under standard conditions (unit activity and pressure). Some values are shown in Table 1 for species inaqueous solution. By convention, tabulated potentials refer to reduction reactions, with electrons on the left as in theabove equation. Only differences in electrode potential are significant, the absolute values having no meaning exceptin comparison with the H+/H2 potential (zero by definition).Direction of reactionComparing two couples Ox-Red, a more positive potential means that the corresponding species Ox is a strongeroxidizing agent.
Thus from Table 1 Br2 is a stronger oxidizing agent than Fe3+ and will so oxidize Fe2+, the productsbeing Br− and Fe3+. Conversely, a lower (more negative) potential means that the corresponding Red is a strongerreducing agent. Thus zinc metal is a stronger reducing agent than dihydrogen, and will reduce H+:(2)The equilibrium constant K for such a reaction can be calculated fromwhere R is the gas constant, T the temperature in kelvin,the difference in the two electrode potentials and n thenumber of electrons in each half reaction: this must be the same for both half reactions in a balanced equation (seeexample in Topic B4).
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