Geometrical isomerism And Oxidation state of Co-ordination Compounds

1.6 Geometrical Types and Isomers Coordination compounds show a wide variety of regular, and an infinite range of irregular, geometries for the arrangement of the ligands about the metal centre. However, for the first row transition metals, a few geometries by far outweigh all of the others. The regular polyhedra upon which complexes are commonly based are the octahedron (six coordination) and the tetrahedron (four coordination). A significant number of four coordinate complexes exhibit a planar geometry and in Chapter 7 we rationalize the occurrence of this structural geometry. One of the consequences of complexes adopting specific geometries is the occurrence of isomers. We review these only briefly, and the interested reader will find more information in the "suggestions for further reading" at the close of this chapter. Several different types of isomers arise in transition-metal coordination compounds, and these are described below. Structural isomers: These are compounds in which the isomers are related by the interchange of ligands inside the coordination sphere for those outside it. A classical example of this phenomenon is observed in the compounds of formula CrCl3(H2O)6- As usually obtained from chemical suppliers, this is a green solid in which only two of the chloride ions are coordinated to the metal. This is formulated [Cr(H2O)4Cl2]Cl^H2O. Solutions of this compound in water slowly turn bluegreen as a coordinated chloride ion is replaced by a water molecule and the complex [Cr(H2O)5Cl]Cl2-H2O may be isolated. More commonly, structural isomers are related by the exchange of anionic ligands and counter ions, rather than neutral ligands. Typical examples include the pair of complexes [Co(en)2Br2]Cl and [Co(en)2BrCl]Br. Linkage isomerism: This is a special type of structural isomerism in which the differences arise from a particular ligand which may coordinate to a metal ion in more than one way. In Table 1-3 we indicated that a ligand such as thiocyanate could bond to a metal through either the nitrogen or the sulfur atom, and the complex ions [Co(NH3)5(7VCS)]2+ and [Co(NH3)5(SCN)]2+ are related as linkage isomers. Coordination isomerism: This is an interesting type of isomerism which can occur with salts in which both the cation and the anion are complex ions. Consider the salt [Co(bpy)3] [Fe(CN)6] containing one cobalt (m) and one iron (m) centre: coordination isomers of this would include [Fe(bpy)3] [Co(CN)6], [Co(bpy)2 (CN)2] [Fe(bpy)(CN)4], [Fe(bpy)2(CN)2][Co(bpy)(CN)4], and [Co(bpy)3][Fe(CN)6]. Geometrical isomerism: This is an important topic which played a crucial role in the development of coordination chemistry. Werner used the number of isomers 10 1 An Introduction to Transition-Metal Chemistry which could be isolated for a range of cobalt(m) complexes to establish the octahedral character of the CoL6 species. A planar complex of the type [Pt(NH3)2Cl2] can exist in two forms depending upon the relative spatial orientation of the two chloride ligands. They can be at 90° to each other to give the cis form (1.1), or at 180° to give the trans isomer (1.2). Cl I H3N — Pt-Cl NH3 1.1 NH3 Cl Pt-Cl I NH3 1.2 In six coordinate complex ions such as [Co(NH3)4Br2]+, a similar situation exists, in which the bromine ligands adopt either a cis (1.3, 1.4) or a trans arrangement (1.5). The reader should note the identity of the cis isomers despite the different drawings (1.3 and 1.4). In a similar manner, complexes of the type [MX3Y3] may adopt two structures, depending upon the relative arrangement of the three identical groups in the octahedron. If the three X groups are arranged about a single triangular face, then the/acia/ (or/ac) isomer (1.7) is obtained, whereas if they are arranged in three of the four sites of the equatorial plane, the meridional (or mer) isomer (1.6) is obtained. NH3 Br H3N, H3N* :Co: -Br H3NH3N* H3NH3N* Br :Co; 'NH3 NH3 1.3 1.4 Br trans 1.5 X- 1.6 Y foe 1.7 Notice the 'loose' use of the term octahedral to describe six-coordinate complexes which are based upon an octahedral geometry, but which, by virtue of the presence of different ligand types, are of lower symmetry than Oh This is a common usage which should give rise to no difficulties. Note also how introduction of chelating 1.6 Geometrical Types and homers 11 ligands into the coordination shell may reduce the number of isomers which are possible. Thus, although there are two isomers of [Pt(NH3)2Cl2], it is only possible to form the cis isomer of [Pt(en)Cl2] (1.8). This is because the relative positions of the nitrogen donor atoms in the en ligand are dictated by the CH2CH2 linker group - the two donor atoms cannot 'stretch' to occupy trans positions. Similarly, it is only possible to obtain the cis isomer of the cation [Co(NH3)4(en)]3+ (1.9). NH3 Cl I Pt-Cl NH2 1.8 H3N. H2 • Ν- NH3 1.9 A final type of isomerism which we mention here also arises most commonly when chelating ligands are present. If a molecule possesses neither a plane nor a centre of symmetry, it is chiral. (This definition is not strictly correct, but will suffice for most transition-metal complexes.) Chiral species may exist in two forms which are related as mirror images. These have identical chemical and physical properties unless they are interacting with something else which is chiral, in which case they differ. That may be a chiral reagent (to give diastereomeric compounds) or polarized light. A typical example of a chiral complex is found when three chelating ligands are coordinated to an octahedral centre, as in the cation [Ru(bpy)3]2+. Two different forms of this cation, related as mirror images, are possible (1.10 and 1.11). These may be separated by formation of salts with chiral anions, and exhibit different and opposite rotations of polarized light. Note also that the cation [Co(en)2Br2]+ (1.12 and 1.13) is chiral, but [Co(NH3)4Br2]+ is not. CT^D Φ'"Ό 1.10 1.11 1.12 1.13 Another way of drawing these isomers emphasizes the three-fold nature of the basic octahedron rather than its four-fold properties (1.14-1.17). 1.7 Oxidation State Oxidation state is a frequently used (and indeed misused) concept which apportions charges and electrons within complex molecules and ions. We stress that oxidation state is a formal concept, rather than an accurate statement of the charge distributions within compounds. The oxidation state of a metal is defined as the formal charge which would be placed upon that metal in a purely ionic description. For example, the metals in the gas phase ions Mn3+ and Cu+ are assigned oxidation states of +3 and +1 respectively. These are usually denoted by placing the formal oxidation state in Roman numerals in parentheses after the element name; the ions Mn3+ and Cu+ are examples of manganese(m) and copper(i). Older texts often employ an alternative nomenclature in which the suffixes -ous and -ic are encountered. In general, these labels only apply to the most common oxidation states of the metals, -ic referring to the higher oxidation state and -ous to the lower. Using this nomenclature, copper(n) is referred to as cupric and copper(i) as cuprous. The system works well if there are only two common oxidation states for a metal ion, but if there are more, the scheme becomes either ambiguous or unwieldy as a variety of prefixes are added. It is usually easy to define the oxidation state for simple compounds of the transition metals. In the case of neutral compounds, we assign charges as if the compound were ionic. Thus, MnCl2 is regarded as (Mn2+, 2Cl~} and is correctly described as manganese(n) chloride. Similarly, WO3 as (W6+, 3O2~} is tungsten(vi) oxide. Since ligands which bear no formal charges in an ionic formulation may be ignored, [Cr(H2O)3Cl3] is a chromium(m) compound, and Ni(OH)2, NiBr2, NiBr2-SH2O, NiBr2-OH2O and NiBr2-9H2O are all nickel(n) compounds. The assignment of oxidation state makes no implications regarding the nature of the bonding within the molecule - all of the various hydrated forms of CrCl3 are chromium(m) compounds. Oxidation state is merely a formal scheme: there is no implication that tungsten(vi) oxide necessarily contains W6+ ions. Furthermore, problems with the assignment of oxidation state can arise with even apparently simple compounds. Consider, for example, Fe3O4. If the compound were ionic, we would have four O2 ions. In order for the entire compound to be neutral, the three iron atoms must possess an overall charge of +8. The ensuing assignment of an oxidation state of +8/3 to each iron is not particularly meaningful. A compound of this type is best regarded as a mixed oxidation state oxide, (FeO + Fe2O3) or Fe11Fe2 111O4, in which there are both iron(n) and iron(m) centres. Cations and anions are treated in an exactly similar manner, remembering to take the overall charge of the species into account. If only neutral ligands are present, the oxidation state of the metal ion is equal to the overall charge on the ion. Thus, [Fe(H2O)6J3+ and [Ni(NH3)6]2+ are iron(m) and nickel(n) complexes respectively. If charged ligands are present, formal charges are assigned on the basis of an ionic description. Thus, the ion [Ni(CN)4]2" is treated as containing a cationic nickel centre 7.7 Oxidation State 13 and four anionic cyanides. Since the four cyanides give a total charge of -4, the nickel must be assigned a charge of +2 in order for the ion to possess an overall charge of -2, and it is therefore a nickel(n) complex. Similarly, [MnO4] ~ is treated as (Mn7+, 4O2"} and is a manganese(vn) compound. Once again* we stress that this in no way implies that the ion [MnO4] ~ actually contains a Mn7+ ion. By the way, aqueous solutions of transition-metal compounds frequently contain ions such as [M(H2O)6]^+: as water is the most common solvent encountered in chemical reactions, these species are often (but incorrectly) referred to as solutions containing M"+ ions . It is quite possible for a metal centre to possess a zero or negative oxidation state. Thus, the species [Cr(CO)6] and [Fe(CO)4]2- are chromium(O) and iron(-2) complexes. We will see in a later chapter that it is not a coincidence that these low formal oxidation states are associated with ligands such as carbon monoxide. Some ligands pose problems in the assignment of a formal oxidation state to a metal centre. Nitric oxide is a case in point. The ligand may be formulated as either anionic NO" or cationic NO+, and there follows the appropriate ambiguity in assignment of the oxidation state of the metal ion to which it is bonded. These problems arise when it is not clear as to what charge is appropriate to assign to the ligands in the ionic limit. We have repeatedly emphasized the formal character of the concept of oxidation state and turn now to a different general concept which helps us address the real electron distributions in compounds. It is very common for inorganic chemists to 'neglect' or 'ignore' the presence of solvent molecules coordinated to a metal centre. In some cases, this is just carelessness, or laziness, as in the description of an aqueous solution of cobalt(n) nitrate as containing Co2+ ions. Except in very concentrated solutions, the actual solution species is [Co(H2O)6J2+. In other cases, it is not always certain exactly what ligands remain coordinated to the metal ion in solution, or how many solvent molecules become coordinated. Solutions of iron(ni) chloride in water contain a mixture of complex ions containing a variety of chloride, water, hydroxide and oxide ligands. When dealing with the kinetic or thermodynamic behaviour of transition-metal systems, square brackets are used to denote concentrations of solution species. In the interests of simplicity, solvent molecules are frequently omitted (as are the square brackets around complex species). The reaction (1.1) is frequently written as equation (1.2). [Co(H2O)6J2+ + 4Cl - = [CoCl4]2- + 6H2O (1.1) Co2+ + 4Cl- = [CoCl4]2- (1.2) Whilst this will be satisfactory when dealing with kinetic data in which reactions involving the solvent will not explicitly appear in the rate equations, it is not appropriate when we consider equilibrium constants. As an exercise, consider the formation of [Ni(en)3]2+ from aqueous solutions of nickel(n) chloride and en (en = H2NCH2CH2NH2); write the equations with the inclusion and the omission of the water molecules.