Inorganic Chemistry
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Inorganic Chemistry

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with appropriate electronegativities, 
and arranged in a suitable geometrical arrangement, to optimize the binding of a specific metal ion. 
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Effect of pH 
pH changes will affect complex formation whenever any of the species involved has Brønsted 
acidity or basicity (see Topic E2). Most good ligands (except Cl\u2212, Br\u2212 and I\u2212) are basic, and 
protonation at low pH will compete with complex formation. This is important in analytical 
applications. For example, in titrations with EDTA, Fe3+ (for which K1 is around 10
25) can be 
titrated at a pH down to two, but with Ca2+ (where K1 is about 10
10) a pH of at least seven is 
required because at lower pH values complex formation is incomplete. 
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Section E\u2014Chemistry in solution 
Consider an ionic solid that dissolves in water according to the equation: 
The equilibrium constant for this reaction, 
is known as the solubility product of MnXm. The form of this equilibrium is important in 
understanding effects such as the influence of pH and complexing (see below) and also the common 
ion effect: it can be seen that adding one of the ions Mm+ or Xn\u2212 will shift Reaction 1 to the left and 
so reduce the solubility of the salt. Thus AgCl(s) is much less soluble 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 difference between these quantities is large in concentrated ionic solutions, and Ksp is 
quantitatively reliable as a guide to solubilities (measured in concentration units) only for very dilute 
solutions. Nevertheless, a thermodynamic analysis of the factors determining 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 
thermodynamic cycle that relates the overall \u394G to two separate 
Key Notes 
Thermodynamics The equilibrium constant for dissolving an ionic substance is known as the solubility 
product. It is related to a Gibbs free energy change that depends on a balance of lattice 
energy and solvation energies, together with an entropy contribution. 
Major trends in water Solids tend to be less soluble when ions are of similar size or when both are multiply 
charged. Covalent contributions to the lattice energy reduce solubility. 
Influence of pH and 
Solubility increases in acid conditions when the anion is derived from a weak acid, for 
example hydroxide, sulfide or carbonate. Amphoteric substances may dissolve again 
at high pH. Complexing agents also increase solubility. 
Other solvents Highly polar solvents show parallels with water. Compounds with multiply charged 
ions are often insoluble in less polar ones, but different donor properties and 
polarizability play a role. 
Related topic Lattice energies (D6) 
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steps: (i) the formation of gas-phase ions; (ii) their subsequent solvation. The enthalpy contributions 
involve a balance between the lattice energy of the solid and the solvation enthalpies of the ions (see 
Topics D6 and E1). In a solvent such as water with a very high dielectric constant these contributions 
almost cancel. Nevertheless, some of the solubility trends summarized below can be understood in 
terms of the changing balance between lattice energies (proportional to 1/(r++r\u2212), where r+ and r\u2212 
are radii of individual ions) and the sum of the individual solvation enthalpies (each roughly 
proportional to 1/r). For example, a small ion has a large (negative) solvation energy, but when 
partnered by a large counterion cannot achieve an especially large lattice energy. With ions of very 
different size, therefore, solvation is relatively favored and solubility tends to be larger than with 
ions 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 \u394S 
contribution. Overall \u394S values for dissolving ions with multiple charges are usually negative, an 
effect that tends to lower the solubility as mentioned below. 
Major trends in water 
The aqueous solubility of ionic compounds is important in synthetic and analytical chemistry (see 
Topics B6, B7), and in the formation of minerals by geochemical processes (see Topic J2). The most 
significant trends are as follows. 
Fig. 1. Thermodynamic cycle for the solution of an ionic solid MX.
(i) Soluble salts are more often found when ions are of very different size rather than similar 
size. Thus in comparing salts with different alkali metal cations, lithium compounds are the 
least soluble of the series with OH\u2212 and F\u2212, but the most soluble with larger cations such as 
Cl\u2212 or . This principle is often useful in preparative reactions and separations. If it is 
desired to precipitate a large complex anion, a large cation such as tetrabutyl ammonium 
+ can be helpful. 
(ii) Salts where both ions have multiple charges are less likely to be soluble than ones with single 
charges. Thus carbonates and sulfates of the larger group 2 cations are 
insoluble. An important factor is the negative solvation entropies of the ions. 
(iii) With ions of different charges, especially insoluble compounds result when the lower charged 
one is smaller (as this gives a very large lattice energy). Thus with M3+ ions, fluorides and 
hydroxides are generally very insoluble, whereas heavier halides and nitrates are very 
(iv) Lower solubility results from covalent contributions to the lattice energy (see Topic D6). This 
happens especially with ions of less electropositive metals in combination with more 
polarizable cations. Late transition and post-transition elements often have insoluble sulfides 
(see Topic J2); insoluble halides (but not generally fluoride) also occur, for example with 
Ag+ and Pb2+. 
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Influence of pH and complexing 
Any substance in solution that reacts with one of the ions formed in Reaction 1 will shift the 
equilibrium to the right and hence increase the solubility of the solid. pH will therefore influence the 
solubility in a range where one of the ions has significant Brønsted acid or base properties (see Topic 
E2). The solubility of NaCl, for example, should not be affected by pH, but when the anion is the 
conjugate base of a weak acid solubility will increase at low pH. Metal oxides and hydroxides 
dissolve in acid solution, and conversely such solids may be precipitated from a solution containing a 
metal ion as the pH is increased. The solubility range depends on the Ksp value: for example, Fe(OH)
3 is precipitated at much lower pH than the more soluble Fe(OH)2. At high pH the acidity of the 
hydrated metal ion may come into play and amphoteric 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 sulfide will liberate H2S, and on a carbonate CO2 from the decomposition of carbonic 
Any ligand that complexes with the metal ion will also increase solubility. AgCl dissolves in 
aqueous ammonia by the formation of [Ag(NH3)2]
+. Addition