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Chemistry 2810 Lecture Notes 6. Dr. R. T. Boeré Page95 d and f-Block Elements and Coordination Chemistry Sc Y Ti V Cr Mn Zr Nb Mo Tc Hf Ta W Re Rf Db Sg Bh Fe Co Ni Cu Zn Ru Rh Pd Ag Cd Os Ir Pt Au Hg Hs Mt La Ce Pr Nd Pm Sm Eu Gd Tb Dy Ho Er Tm Yb Lu Ac Th Pa U Np Pu Am Cm Bk Cf Es Fm Md No Lr The transition elements comprise those metals in Groups 3 to 12, i.e. the d-block elements, as well as the so-called “inner transition elements” which are the lanthanides and actinides, also know as the f-block elements. These are all metallic elements, so they are often called the transition metals, which we take to be synonymous with the transition elements. The transition metals include many of the most important structural metals for an industrial society, including iron (Fe), chromium (Cr), nickel (Ni) and copper (Cu). The text-book describes how some of these metals are won from their ores. Our focus in lecture is going to be on coordination complexes of these elements. We could study the properties of the transition elements in crystalline solids. However, there are complexities. First of all, many additional crystal structure types would have to be studied. Many of these structures involve a large amount of covalent bonding, so that our simple ionic model is not completely adequate. But in reality, we could do so, and the decision to move to soluble coordination compounds is a deliberate one. The study of the transition metals in solids have been dominated by solid-state physicists, and is closely tied to important technical advances, such as the development of the first Ruby lasers. Chemists have been more involved in studying soluble metal complexes, and today these have become extremely important compounds for industrial chemical processes (e.g. in the petrochemical and pharmaceutical industries), as well as in medicine an as model systems for many plant and animal proteins and enzymes. There is a strong conceptual relationship between the metals in a solid and in solution. In a coordination compound, we suround the metal ion by the same kind of donor atoms as are in the anions found in their crystalline compounds. Thus, to mimic an oxide environment, we can simply use water, or another oxygen-containing species. Of course, the metal ions will act as Lewis acids (both hard and soft acids are found among the transition elements, as well as everything in between), and the donor atoms act as Lewis bases. Thus coordination compounds are nothing more or less than Lewis acid-base complexes. You should review your tables of hard and soft acids and bases at this time. If we want to mimic a crystalline metal halide (from the hard fluoride to the soft iodide), we can employ the halogen anions as the donor atoms. Alternatively, we can employ donor atoms such as nitrogen or carbon that are rare in solids, although nitrides and carbides of these metals do exist. In summary, a coordination compound is a Lewis acid/base complex that solubilizes the metal ion while preserving the local environment found in crystalline solids of these metals. The coordination number of the central ion, the geometry and the nature of the donor atom found in solids can often be mimicked quite accurately in soluble complexes. 6.1 the d-block metals have multiple oxidation states The transition elements are those which have open d-subshells (BACK POINTER: atom electron configurations). There are several important consequences which are due to these d-orbitals. First of all, they are the cause of the large number of common oxidation states for transition metals, as seen in the diagram at the left. The most common oxidation states are given by the square entries. Notice the pattern that the ends of the series, there are fewer common oxidation states, whereas in the middle (i.e. exactly where the d-orbitals become halffilled) there is the greatest variety. Removing d-electrons is relatively easy compared to s,p electrons, and this is the reason that high-oxidation state forms of these elements can be obtained. A second consequence of d-electrons is that coordination compounds are especially stable among the transition elements. It is important to stress that all metallic and metalloid elements form coordination compounds, but they are definitely most well established and most stable among the transition elements. Chemistry 2810 Lecture Notes 6.2 Dr. R. T. Boeré Page96 Introduction to coordination chemistry A metal cation acts as a Lewis acid. A Lewis acid is a substance that can accept a pair of electrons from another atom to form a new bond. The electron pairs which are received come from surrounding groups which are called ligands. Ligands are typically either neutral or anionic atoms or molecules. Ligands act as Lewis bases. A Lewis base is a substance that can donate a pair of electrons to another atom to form a new bond. (BACK POINTER: review the Lewis theory which we have already covered.) The combination of a metal cation and all its ligands is called a coordination complex. These difficult-sounding definitions are much clearer from a few examples. First, we construct the Lewis diagrams of the ligands H2O, Cl-, and CO - ¨: :Cl :O: :C O: H H ¨ is dissolved in water, the resulting Fe2+(aq) ion is in reality a coordination complex between When a metal such as Fe2+ iron(II) and water: Fe2+ + 6 H2O → [Fe(OH2)6]2+ (here water is written OH2 to emphasize that the electron pair on oxygen of water is acting as the Lewis base donor site to the metal; the charge on the resulting complex ion is simply the sum of the charges of the metal ion and those on the ligands; square brackets are used to isolated the complex ion, and we refer to the six water molecules as comprising the primary coordination sphere of the Fe(II) ion. Copper(II) ions in the presence of a high concentration of chloride ion forms a chloro complex: Cu2+ + 4 Cl- → [CuCl4]2- (here the overall charge is -2 since that is the sum of the ligands and the metal ion) Nickel when finely divided reacts readily with an atmosphere of gaseous carbon monoxide to form the liquid nickel tetracarbonyl: Ni + 4 CO → [Ni(CO)4] (here the overall charge is 0, since the metal and the CO are all uncharged) In summary, coordination compounds can be overall neutral, cationic or anionic, and the charge is the sum of all the constituents. Charged complex compounds are called complex ions, and these can form salts with appropriate counter ions. For example, the complexes mentioned above might be isolated as the salts [Fe(OH2)6]Br2 and K2[CuCl4] You have already met several transition metal coordination complexes back during the Chemistry 2000 laboratory. For example, in the first lab you prepared, and then analyzed, the complex salt K3[Fe(C2O4)3].3H2O. Later on, you determine the relative number of thiocyanate and water ligands in complex ions of the type [Fe(SCN)n(H2O)6 - n](3-n)-. Several of the compounds you may prepare in lab 6, the Thirteen Mystery Test Tubes, are in fact coordination compounds, while you may also have met copper complexes [Cu(OH2)4]2+ and [Cu(NH3)4]2+. In the 2810 lab you will study the electronic spectra of a number of simple aqua, chloro and ammine complexes of the d and f elements. 6.2.1 Common Ligands There are literally thousands of ligands which have at times been used towards metals to form coordination complexes. Some are just common small molecules or ions, such as water, ammonia and the halogen and pseudo-halogen ions. Such species typically have a single donor atom available, and even if this donor has more than one electron lone pair, it is rare for one atom to donate more than once to a given metal. The most common exception to this rule is the oxo ligand, O2–, for which M=O double bonds are extremely common. In all such cases where a single donor atom is linked to the metal, the ligand is said to be monodentate (literally, "single-toothed"). Chemistry 2810 Lecture Notes Dr. R. T. Boeré Page97 However, using simple organic chemistry, it is fairly easy to link together in one molecule two or more donor atoms linked by an inert "backbone". These more complicated ligands are called chelating ligands, from the Greek word chele, meaning “claw”. Chelating ligands can have two, or more donor sites. With two donors, they are called bidentate ("doublytoothed"), with three, The complex ion [Co(en)3]3+ The neutral complex [CoEDTA] tridentate, etc. Whereas iron(III) has room in the primary coordination sphere for six water molecules in the complex ion [Fe(OH2)6]3+, it only has room for three chelating bidentate oxalate ions in the complex ion [Fe(C2O4)3]3-. Can you see why? The pictures show some computerr-drawn "ball and stick" diagrams of coordination complexes with chelating ligands. The structure of [Co(en)3]3+ is that of a complex cation, which can be isolated from solution as a salt with any suitably charged anions. On the other hand, [CoEDTA] where the cobalt is in the +3 oxidation state can be a neutral molecule (an hence is not particularly soluble in water.) Alternatively, you should get used to interpreting simple line diagrams of typical coordination complexes, such as the following examples: 3- O OH2 2+ NH 3 H2O OH 2 OH2 H 3N NH3 NH3 C 3+ OH2 O OH2 C O C O O Fe Cu Cu H2O 2+ O Fe H2O OH 2 O OH2 C O O O C C O O Perspective line drawings of some typical coordination complexes We identify the total number of Lewis base attachments to a given metal center as the coordination number. Thus in the example discussed above, both [Fe(OH2)6]3+ and [Fe(C2O4)3]3- have a coordination number of six. In fact the iron(III) ion does not particularly care whether the six oxygen donor atoms come from six monodentate water molecules or three bidentate oxalate ions. One of the most amazing ligands is EDTA, which supplies up to six donor atoms to a single metal ion, the ligand wrapping around the metal and encasing it. EDTA4- is a hexadentate ligand. In general, chelating ligands have additional stability over non-chelating or monodentate ligands. One of the reasons that the oxalate ion displaces H2O during the preparation of [Fe(C2O4)3]3- in the Chem2000 lab experiment is that oxalate is a chelating ligand. Formula Name Abbreviation H2 O aqua use formula F–, Cl–, Br–, I– fluoro, chloro, bromo, iodo use formula ammine use formula Structure :O: H3 N - ¨ Chg monodentate 0 monodentate –1 monodentate 0 monodentate (through C) monodentate (through C) monodentate (often M=O) 0 H H - : F̈ : Donor number :C̈l : , ¨ ¨ N , etc. H H H CO carbonyl use formula CN– cyano use formula O2– oxo use formula :C O: - :C N: -2 : Ö : ¨ –1 –2 Chemistry 2810 Lecture Notes CO32– carbonato Dr. R. T. Boeré Page98 O use formula C C 2 O4 2– oxalato -O O¨ ox H2NCH2CH2NH2 ethyelenediamine C12H8N2 phenanthroline C10H12O8 ethylenediaminetetraacetato -O ¨ en ¨O O H2 H2 C C ¨H C 2 N ¨ -O C 2 bidentate 0 bidentate 0 hexadentate 4– NH2 N -O C –2 - ¨ ¨ ¨ EDTA bidentate C H2N phen –2 - O C bidentate N 2 H2 C ¨ N - ¨ CO2 CO2 The table above lists a group of common ligands that you need to learn to be able to answer the assignment and test questions for this course. You need to remember the name and abbreviation of each complex, the number of donor atoms, as well as the overall charge on the ligand (typically either zero or some negative number; there do exist a very small number of positively charged ligands, but these are extremely rare.) All questions will be restricted to complexes using only this small set of ligands. 6.2.2 Naming Coordination Compounds Just as there are rules for naming simple inorganic and organic compounds, coordination compounds are named according to an established system. For example, the following compounds are named according to the rules outlined below. Compound [Ni(H2O)6]SO4 [Cr(en)2(CN)2]Cl K[Pt(NH3)Cl3] Systematic Name Hexaaquanickel(II) sulfate Dicyanobis(ethylenediamine)chromium(III) chloride Potassium amminetrichloroplatinate(II) As you read through the rules, notice how they apply to the examples above: 1. In naming a coordination compound that is a salt, name the cation first and then the anion. (This is how all salts are commonly named). 2. When giving the name of the complex ion or molecule, name the ligands first, in alphabetical order, followed by the name of the metal. a. If a ligand is an anion whose name ends in -ite or -ate, the final e is changed to o (as in sulfate → sulfato or nitrite → nitrito). b. If the ligand is an anion whose name ends in -ide, the ending is changed to o (as in chloride → chloro or cyanide → cyano). c. If the ligand is a neutral molecule, its common name is usually used. The important exceptions to this rule are water, which is called aqua, ammonia, which is called ammine, and CO, called carbonyl. d. When there is more than one of a particular monodentate ligand with a simple name, the number of ligands is designated by the appropriate pre-fix: di, tri, tetra, penta, or hexa. If the ligand name is complicated (whether monodentate or bidentate), the prefix changes to bis, tris, tetrakis, pentakis, or hexakis, followed by the ligand name in parentheses. 3. If the complex ion is an anion, the suffix -ate is added to the metal name. 4. Following the name of the metal, the oxidation number of the metal is given in Roman numerals. Complexes can be considerably more complicated than those described in this chapter; then, even more rules of nomenclature must be applied. The brief rules just outlined, however, are sufficient for the vast majority of complexes. Here follow some examples of how the rules can be applied: 1. [Cu(NH3)4]SO4 The sulfate ion has a 2– charge, so the complex ion has a 2+ charge (that is, [Cu(NH3)4]2+ ). Because NH3 is a neutral molecule, the copper ion is Cu 2+ . The compound’s name is therefore tetraamminecopper(II) sulfate. 2. K2[CoCl4] Two K+ ions occur in this compound, so the complex ion has a 2– charge ([CoCl4]2– ). Because four Cl– ions occur in the complex ion, the cobalt center is Co2+ . Thus, the name of the compound is potassium tetrachlorocobaltate(II). Chemistry 2810 Lecture Notes Dr. R. T. Boeré Page99 Co(phen)2Cl2 This is a neutral compound. Because two Cl– ions and two neutral phen (phenanthroline) ligands are bonded to a cobalt ion, the metal ion must be Co2+ . This means the compound name is dichlorobis(phenanthroline) cobalt(II). [Co(en)2(H2O)Cl]Cl2 Here the complex ion has a 2+ charge because it is associated with two uncoordinated Cl– ions. The cobalt ion must be Co3+ because it is bonded to two neutral en (ethylenediamine) ligands, one neutral water, and one Cl– . The name is aquachlorobis(ethylenediamine)cobalt(III) chloride. K3[Fe(C2O4)3].3H2O First of all, the .3H2O are waters of crystallization. These are additional water molecules which are incorporated into empty spaces in the crystals of the product, and are called hydrates. Thus the salt is a trihydrate. (NB: many so-called hydrates of simple salts of metal ions are in fact metal aqua complexes. Thus CuSO4.5H2O is in fact actually the complex salt [Cu(OH2)4]SO4.H2O. Just how many water molecules in a given hydrated salt are coordinated to a metal ion must be established by experiment. It is not something you can know automatically. However, you should from now on be aware that hydrated salts may in fact contain coordinated water molecules!). Complex salts obey the same convention as simple salts in that the name of the cation is always given first, followed by the name of the anion. Thus the first part of the name of the salt is potassium. Since K always gives a 1+ ion, its oxidation state is not specified. The complex anion is named by giving the name of the ligand first, followed by that of the metal and the oxidation state of the metal in Roman numerals. If the ligand is an anion whose name ends in ite or ate, the ending is changed to o. Thus oxalate becomes oxalato. When there is more than one ligand of the same type (the normal situation) the number is given by di, tri, tetra, penta, hexa, etc. Thus our example is trioxalato. If the complex ion is an anion, the suffix ate is added to the metal name. Iron is a little bit odd in that we revert to the Latin name for anionic forms. Thus ironate is never used; instead we call it ferrate. The oxidation state of the iron is +3, so this is given by (III). We are now ready to name the complete salt which you will prepare next week in the laboratory: potassium trioxalatoferrate(III) trihydrate. That’s quite a mouthful! You can see why on the whole chemists prefer a picture of a molecule to its name. 3. 4. 5. 6.2.3 Structure and Isomerism The structure of coordination compounds is a combination of the coordination number (CN) and the possible geometries of the attached ligands. If CN = 2, the structure is always linear, for example in [AgCl2]- and [Au(CN)2]-. Cl Cl Ag If CN = 4, the geometry may be one of two structures: square planar or tetrahedral. For example, the complexes [Cu(OH2)4]2+ and [Cu(NH3)4]2+ are both tetrahedral, as indicated in the picture above. On the other hand, the complex [Ni(CN)4]2- has the square-planar shape. Note that these shapes in general are not predictable by VSEPR theory. They must be established by experiment. Isomerism is not possible in tetrahedral complexes unless all four attached groups are different. This is an extremely rare situation for coordination compounds (but happens fairly frequently for sp3-hybridized carbon atoms). One of the consequences of square planar geometry is that cis and trans isomers are possible. Such geometrical isomers only show up if there are two different types of ligands present on the molecules. A very famous example of this kind of isomerism occurs for the formula [PtCl2(NH3)2]. The cis form of this neutral coordination complex is a potent anticancer drug know in medicine as cisplatin. The trans form has no antitumor activity at all, and is useless as a drug. If CN = 6, the structure is almost always octahedral. This is also the most common geometry for transition metal coordination complexes. Octahedral complexes have a wide range of possible isomers, depending on the type of ligands attached. If the general formula is MX4Y2, cis and trans geometrical isomers are possible. On the other hand, if the general formula is MX3Y3, mer and fac geometrical isomers are possible. Multidentate ligands can change the general picture. For example M(X∩ X)3 complexes, in which X∩ X represents a chelating bidentate ligand, have only one feasible geometry; however such complexes do form optical isomers. This means that they are chemically indistinguishable, but rotate plane polarized light in opposite directions. You will learn more about optical isomers in introductory organic chemistry classes. You should be aware, however, that optical isomerism is rooted purely in the geometry of the molecule, not in the types of elements involved. Chelating complexes of the general formula M(X∩ X)2Y2 can exist as cis and trans geometrical isomers. The cis form of M(X∩ X)2Y2 exists as two optical isomers as well. NH3 Cl Pt NH3 Cl NH3 trans-[PtCl2(NH3)2] Cl Pt NH3 Cl H 3N NH3 Pt H 3N NH3 Cl 2+ NH3 2+ Cl H3 N NH3 Pt H3 N Cl Cl + Cl H 3N NH3 Pt H 3N cis-[PtCl2(NH3)4]2+ NH3 Pt Cl Cl Cl Cl NH3 mer-[PtCl3(NH3)3]+ fac-[PtCl3(NH3)3]+ cis-[PtCl2(NH3)2] trans-[PtCl2(NH3)4]2+ + Cl H 3N Chemistry 2810 Lecture Notes Dr. R. T. Boeré Page100 All of these isomers are due to only to geometry. You are encouraged to make some models and prove to yourself that these differences are possible, and that they are the only possible variations. In addition, you should obtain the point groups for each structure. The computergenerated ball-and-stick diagrams at right illustrate cis and trans isomerism for sixcoordinate octahedral shapes. Note that the cis isomer in this case involves a chelating ligand, so it will exist in two optical isomers, i.e. it is an cis isomer of [Co(en)2Cl2]+ trans isomer of [Co(en)2Cl2]+ example of M(X∩ X)2Y2. Remember from our section on point group symmetry that optical isomers exist for point groups which do not have any form of Sn symmetry axis, including a centre of symmetry or a mirror plane. Thus cis-[Co(en)2Cl2]+ belongs to the point group C2, which has E, C2 symmetry elements only. 6.2.4 Electronic Structure and Colour of Transition Metal Coordination Compounds One thing distinguishes transition metal coordination compounds from all other coordination compounds, and that is colour. Transition metal complexes are very often coloured, whereas the metals and metalloids of the s and p blocks form colourless complexes. Consider the nitrate salts of Fe3+, Co2+, Ni2+, Cu2+ and Zn2+. All of these are metal aqua complexes of the type [M(OH2)n]m+. Why are four of them coloured, while the last is colourless? a) Crystal field theory for octahedral and tetrahedral complexes The origin of these phenomena can be understood by considering what happens to the d orbitals when a coordination compound forms. Remember what the shapes of the five d orbitals are. Important here is also the orientation of these orbitals: The effect on the d orbitals depends on geometry. If the complex is octahedral, with CN = 6, the most common situation, the ligands are found in the sites indicated by the dark circles in the figure at right. A simple electrostatic model, called the crystal field theory, assumes that there will be a significant electronelectron repulsion between the electron pair the ligands donate and any electrons already in the metal d orbitals. The greatest repulsion is felt by those d orbitals directed most closely at the ligands. This leads to an alteration of the energies of the formerly degenerate d orbitals, and in fact causes dx2-y2 and dz2 to become slightly higher in energy than the set dxy, dxz, and dyz. The resulting pattern is shown below. We call the separation between the two sets the crystal field splitting, ∆o, where the “o” stands for octahedral. In an octahedral field, electrons in the shaded d orbitals experience greater repulsion than in the unshaded d orbitals because of greater geometric proximity to the ligand lone pairs. Chemistry 2810 Lecture Notes Dr. R. T. Boeré Page101 The effect of an octahedral crystal field on the five metal d orbitals For tetrahedral complexes, exactly the inverse situation is obtained. Here the CN=4, and the ligands are located where the dark circles are in the geometrical diagram at right. The consequence is that the dxy, dxz, and dyz become higher in energy than dx2-y2 and dz2. Here the crystal field splitting is called ∆T, standing for the tetrahedral case. It is easy to show by geometry that the magnitude of ∆T < ∆o, so long as ligands of these same type are used in both cases. In a tetrahedral field, electrons in the unshaded d orbitals experience greater repulsion than in the shaded d orbitals because of greater geometric proximity to the ligand lone pairs. Finally, for square planar complexes, a third possibility occurs. The pattern here is more complicated, with dxz, and dyz lowest, followed by dz2, dxy, and then dx2-y2. The spacing between the last two is called ∆SP, the square-planar crystal field splitting energy. We will not consider the square planar case, except to say that typically ∆SP is very large. b) The spectrochemical series Changing the type of atoms in the ligand primarily affects the size of the crystal field splitting energies. This ligand effect is called the spectrochemical series. A simplified form of this series is: {small orbital splitting, ∆} Halides < C2O42- < H2O < NH3 = en < phen < CN- {large orbital splitting, ∆} Ligands at the left of the series are often called weak field ligands, while those at the right are strong field ligands. c) Possible electron configurations for d-element complexes The real importance of all this becomes clear when we start putting the electrons in. We need the electron configurations of the cations, e.g. Fe2+ and Fe3+. These are [Ar]3d6 and [Ar]3d5, respectively. Consider the complexes [Fe(OH2)6]2+ and [Fe(C2O4)3]4-, iron(II), and [Fe(OH2)6]3+ and [Fe(C2O4)3]3-, iron(III), all of which can be prepared under suitable conditions. For these cases, and in fact for any configuration from d4 to d7, there are two possible ways of filling the metal orbitals. Let us see how this works out by putting the electrons into these diagrams for octahedral geometry: Chemistry 2810 Lecture Notes d1 d4 Dr. R. T. Boeré d2 Page102 d3 d5 d6 "high spin" forms d8 d7 d4 d9 d10 d5 d6 "low spin" forms d7 The high spin forms predominate for small orbital splitting, and with both water and oxalate ligands, most complexes will be high spin. So the iron(II) and iron(III) complexes will have 5 and 4 unpaired electrons. On the other hand, if cyanide ligands were present there would be large orbital splitting and the low spin forms would predominate, so that [Fe(CN)6]4- and [Fe(CN)6]3- will have 1 and 0 unpaired electrons. Notice also that this last example is diamagnetic (for a definition, see Kotz & Treichel, p. 358), while a free Fe3+ ion is definitely paramagnetic. Exercise - Work out the possibilities that exist for tetrahedral crystal fields as a function of the d counts. For which can high and low spin conditions exist (i.e. those cases where there is a net difference between the two distributions)? In fact, we need only consider the high spin case, since tetrahedral crystal fields are too small to cause spin pairing. high spin situation d1 d4 d8 d2 low spin situation d3 d5 d6 d9 d10 d1 d7 d4 d8 d2 d3 d5 d6 d9 d10 d7