Transitional Metal Ions in Aqueous Solutions
We often write transition-metalAn element characterized by a glossy surface, high thermal and electrical conductivity, malleability, and ductility. ions in aqueousDescribing a solution in which the solvent is water. solutionA mixture of one or more substances dissolved in a solvent to give a homogeneous mixture. with symbols such as Cr3+, Cu2+, and Fe3+ as though they were monatomic, but this is far from being the case. These ions are actually hydrated in solution and can be regarded as complex ions. Thus, for example, the grayish-violet color of many chromium(III) salts when dissolved in H2O is due to the species [Cr(H2O)6]3+ rather than to a bare Cr3+ ion. The same color is evident in many crystalline solids such as [Cr(H2O)6]Cl3 which are known to contain the Cr3+ ion surrounded octahedrally by six H2O molecules. In much the same way the blue color of many solutions of copper(II) salts can be attributed to the species [Cu(H2O)4]2+ and the pale violet color of some solutions of iron(III) salts to the [Fe(H2O)6]3+ ion. Because [Fe(H2O)6]3+ is capable of donating a protonThe positively charged particle in an atomic nucleus; its mass is similar to the mass of a hydrogen atom., the conjugate baseThe base formed when an acid releases a hydrogen ion (proton)., [Fe(H2O)5OH]2+ is generally present when Fe3+ is dissolved in water. This imparts a yellow color to the solution. Fig. 1 shows examples of colored ion complexes in aqueous solution.
Not all salts of transition-metal ions yield the hydrated ion when dissolved in H2O. Thus when CuCl2 is dissolved in H2O, a beautiful green color due mainly to the complex [CuCl2(H2O)2] is produced. This is obviously different from the sky-blue color of [Cu(H2O)4]2+ which is obtained when Copper(II) sulfate or copper(II) nitrate are dissolved. This is because the Cl– ion is a stronger Lewis baseA species that donates a pair of electrons to form a covalent bond. with respect to the Cu2+ ion than is H2O. Thus, if there is a competition between H2O and Cl– to bond as a ligandOne of the small molecules or ions attached to a central metal in a coordination complex. to Cu2+, the Cl– ion will usually win out over the H2O.
The superior strength of the Cl– as a Lewis base is easily demonstrated by adding Cl– ions to a sky-blue solution of copper(II) sulfate. A green color immediately appears due to the formation of chloro complexes:
- [Cu(H2O)4]2+ + Cl– [Cu(H2O)3Cl]+ + H2O
- [Cu(H2O)3Cl]+ + Cl– [Cu(H2O)2Cl2] + H2O
If a large excess of Cl– ion is added, the solution changes color again from green to yellow. This is because of even further displacement of H2O ligands by Cl– ligands:
- [Cu(H2O)2Cl2] + Cl– [Cu(H2O)2Cl3] – + H2O
- [Cu(H2O)2Cl3] – + Cl– [CuCl4]2– + H2O
Figure 2 compares these three aqueous copper complexes.
Under favorable circumstances yellow crystals of salts like Cs2[CuCl4], containing the complex ion CuCl42– can be obtained from these solutions.
Because they might very possibly form complexes with it, one must be careful about what ions are added to a solution containing hydrated transition-metal ions. Not only the chloride ions, but the other halide ions are liable to complex, and the same is true of species like NH3 and CN–. These ligands differ quite a lot in their affinity for a particular metal ion, but the rules governing this situation are not simple. One finds, for instance, that although NH3 will complex very readily with Cu2+ it has little or no affinity for Fe3+. In other words, a ligand which is a strong Lewis base with respect to one metal ion is not necessarily a strong baseA base that dissociates completely or ionizes completely in a particular solvent. with respect to another. There are some ions, however, which almost always function as very weak Lewis bases. The perchlorate ion, ClO4– in particular, forms almost no complexes. The nitrate ion, NO3–, and sulfate ion, SO42–, only occasionally form complexes.
The addition of ligands to a solution in order to form a highly colored complex is often used to detect the presence or absence of a given metal in solution. The deep blue color of [Cu(NH3)4]2+ produced when excess NH3 is added to solution of Cu(II) salts is a case in point. This can be seen in the following video, where a aqueous solution of ammonia is added to a copper sulfate solution:
The initial copper sulfate solution is sky blue, due to the [Cu(H2O)4]2+ complex. When ammonia is added, a precipitate of Cu(OH)2(s) is formed. as it settles to the bottom, it can be seen that the remaining solution is a dark blue, due to the [Cu(NH3)4]2+ complex formed by copper with ammonia.
Other well-known color reactions are the blood-red complex formed between Fe(III) ions and the thiocyanate ion, SCN–, as well as the pink-red complex of Ni(II) with dimethylglyoxime.
While most of the reactions we have been describing are very fast and occur just as quickly as the solutions are mixed, this is not always the case. With certain types of complexes, ligand substitution is quite a slow process. For example, if Cl– ions are added to a solution containing [Cr(H2O)6]3+ ions, it is a few days before the grayish-violet color of the original ion is replaced by the green color of the chloro complexes [Cr(H2O5) Cl]2+ and [Cr(H2O)4 Cl]+. Alternatively the solution may he heated, in which case the green color will usually appear within 10 min. The reaction
- [Cr(H2O)6]3+ + Cl– → [Cr(H2O)5Cl]3+ + H2O
is thus a slow reaction with a high activation energyThe energy barrier over which a reaction must progress in order for reactants to form products; the minimum energy that reactants must have if they are to be converted to products.. Ligand substitution reactions of other Cr(III) complexes behave similarly. In consequence Cr(III) complexes are said to be inertUnreactive. Used to describe coordination complexes that exchange ligands slowly or an electrode in an electrochemical cell that serves only as a surface where reaction can occur and is neither consumed nor added to during reaction., as opposed to a complex like Fe(H2O)63+ which swaps ligands very quickly and is said to be labileReactive; often used to describe coordination complexes that exchange ligands rapidly.. Other examples of inert complexes are those of Co(III), Pt(IV), and Pt(II). Almost all the compounds which were used to establish the nature and the geometry of coordination compounds were inert rather than labile. There is very little point in trying to prepare cisDescribes the relationship between two atoms or groups of atoms, each attached to one of two doubly bonded carbon atoms and located on the same side of the double bond. Also refers to groups located adjacent to each other in an octahedral or square planar coordination complex. and trans isomers of a labile complex, for example, because either will quickly react to form an equilibriumA state in which no net change is occurring, that is, in which the concentrations of reactants and products remain constant; chemical equilibrium is characterized by forward and reverse reactions occurring at the same rate. mixtureA combination of two or more substances in which the substances retain their chemical identity. of the cis and trans forms.
A final complication in dealing with aqueous solutions of transition-metal complexes is their acidIn Arrhenius theory, a substance that produces hydrogen ions (hydronium ions) in aqueous solution. In Bronsted-Lowry theory, a hydrogen-ion (proton) donor. In Lewis theory, a species that accepts a pair of electrons to form a covalent bond.-base behavior. Hydrated metal ions like [Cr(H2O)6]3+ are capable of donating protons to water and acting as weak acids. Most hydrated ions with a charge of + 3, like Al3+ and Fe3+ behave similarly and are about as strong as acetic acid. The hydrated Hg(II) ion is also noticeably acidic in this way. Perhaps the most obvious of these cationic acids is the hydrated Fe(III) ion. When most Fe(III) salts are dissolved in water, the color of the solution is yellow or brown, though the Fe(H2O)63+ ion itself is pale violet. The yellow color is due to the conjugate base produced by the loss of a proton. The equilibrium involved is
- [Fe(H2O)6]3+ + H2O [Fe(H2O)5OH]2+ + H3O+
- Pale violet Brown
If solutions of Fe(III) salts are acidified with perchloric acid or nitric acid, the brown base is protonated and the yellow color disappears from the solution entirely.