Download d-Block Elements

d-­‐Block Metals From period 4, we find the d-­‐block metals. Because they are located between the s-­‐ and p-­‐
block elements, most of the d-­‐block metals are also known as the transi/on metals. The two most fundamental aspects of the d-­‐block elements is that they fill their d-­‐orbitals, moving leA to right, but the d-­‐orbitals of the free atoms are actually higher in energy than the ‘next’ s-­‐orbital, e.g. fourth period elements have the configuraEon 4s2 3dn, fiAh period elements have the configuraEon 5s2 4dn and so on. d-­‐Orbitals z Let’s remind ourselves of the shapes of the five d-­‐orbitals: y x dxy dx2–y2 dxz dz2 dyz The chemical aspects of the d-­‐subshell are quite different to those of the s-­‐ and p-­‐
subshells. The complex symmetry of d-­‐orbitals means that they do not overlap easily with orbitals from other atoms and so they are usually (but not always) considered to form non-­‐bonding orbitals in molecules. The simplest contribuEon of the d-­‐subshell in d-­‐block chemistry is to act as a source of electrons that can be gained and lost in redox reacEons. By losing and gaining these electrons, the metal atoms can achieve different oxidaEon states. These oxidaEon states have different charge/size raEos and hence different chemical properEes. Another important but conceptually harder aspect is that the d-­‐orbitals can parEally and temporarily overlap with other molecules. This behaviour leads to transiEon metals forming the ac/ve centre in many important catalysts and natural enzymes. The Special Chemistry of TransiEon Metals As menEoned above, one of the principal features of transiEon metal chemistry is the ability of the elements to easily form different oxidaEon states by ionising and rearranging different combinaEons of s-­‐ and d-­‐electrons. These variable oxidaEon states combine with different atomic and ionic radii to make the d-­‐block elements almost form a new periodic table within the periodic table. Another important characterisEc is that most transiEon metals can form ions, molecules or assemblies that are somewhere in between. This feature is central to their use in homogeneous catalysts. Note that the first group of the d-­‐block elements (3B) and the last group (2B) are not convenEonally considered to be transiEon metals. This is because the group 3B metals always lose their single d-­‐electron easily to form an empty d-­‐shell, while the group 2B metals have full d-­‐orbitals and do not easily lose d-­‐electrons. Both groups easily lose their s-­‐orbital pair so group 3B metals always form +3 caEons and group 2B metals always form +2 caEons. In short, group 3B and 2B metals behave more like simple metals than transiEon metals. Simple CoordinaEon Complexes TransiEon metals are notable primarily because of their ability to form an extremely wide range of compounds of variable stability. The precise nature of compounds available to any parEcular element depends on three main factors: 1) available oxidaEon states 2) orbital geometry and 3) atomic/ionic radius. Compounds of transiEon metals are commonly called coordina/on complexes, or simply ‘complexes’. These compounds are typically composed of a central metal ion surrounded by one to eight molecules called ligands (though four or six are the most common numbers). Some ligands are easy to add or remove, whereas others are very difficult to add/remove. We say that the ligands coordinate to the metal centre. ligand (NH3) [Cd(NH3)4]2+ metal centre (Cd) OAen, the whole complex has an electrical charge and in these cases, counter ions are required to balance the charge. The counter ions do not usually affect the chemistry of the complex but they do affect the physical properEes, such as melEng point and solubility. In soluEon, the counter ions usually separate from the complex. In the example above, the complex has a +2 charge so counter ions with a total charge of –2 are required, for instance two chloride ions (2 x –1). The total formula is wri_en as [Cd(NH3)4]Cl2, showing that the complex is made from one Cd ion and four NH3 molecules, with two chloride ions as counter ions. Structures Ligands usually bind to the metal centres by donaEng pairs of electrons. This means that ligands are usually Lewis bases and the metal centres are usually Lewis acids. (SomeEmes, however, the exact Lewis acid/base chemistry is much more complex than this) Here are some examples of coordinaEon complexes with their associated geometry. Note that the geometry descripEon applies only to the atoms that are directly coordinated to the metal centre. For example, [Cd(NH3)4]2+ is a tetrahedral complex because the nitrogen atoms directly coordinated to the Cd centre form a tetrahedron. The hydrogen atoms are not directly bonded to the Cd ion and so they are not considered when determining the geometry of the complex. Atoms that can coordinate to the metal centre are called donor atoms and the number of atoms coordinated to the metal centre is called the coordina/on number. linear CN = 2 square planar CN = 4 tetrahedral CN = 4 trigonal bipyramidal CN = 5 octahedral CN = 6 OxidaEon States TransiEon metals can form a variety of oxidaEon states. The most stable states for any metal depends on its posiEon in the d-­‐block and which ligands it is bound to. Here is a quick summary: 1.  Forming a complex raises the energy of the s-­‐subshell above the level of the d-­‐
subshell. This means that the valence s-­‐electrons are usually lost before the d-­‐
electrons (we always lose electrons from higher energy orbitals before lower orbitals). 2.  In complexes, the d-­‐orbitals are no longer degenerate. The exact energies depend on the geometry of the complex. 3.  Completely empty, half-­‐full and full d-­‐configuraEons are more stable than parEally occupied and parEally paired configuraEons. d-­‐block elements will ionise and/or transfer s-­‐ and d-­‐electrons to achieve the more stable d-­‐configuraEons. 4.  (Unfortunately) there are many excepEons to the rules in transiEon metal chemistry. Rela/ve energies of s-­‐ and d-­‐orbitals in transi/on metals s d d s Free metal d octahedral complexes tetrahedral complexes Table of OxidaEon States We can see from this table that the elements in the middle of the d-­‐block have the most variable oxidaEon states. Note that, as for all chemistry, different oxidaEon states of the same element can have very different chemical properEes. Ligands and Lability SomeEmes it is easy for transiEon complexes to exchange their ligands in a reversible process. Such ligands are called labile ligands. Labile ligands are usually neutral molecules or stable anions. In the following example, the chromium complex can easily exchange the labile H2O and Cl– ligands. Ligands (properly, the atoms) a_ached directly to the metal centre are part of the inner coordina/on sphere. In [Cr(OH2)6]Cl3 (above leA), six oxygen atoms from the water molecules form the inner coordinaEon sphere. In [Cr(OH2)4Cl2]Cl (above right), four oxygen atoms and two chloride ions form the inner coordinaEon sphere. (There is an outer coordina/on sphere beyond the complex, which is composed of counter ions and solvent molecules but it is not as clearly defined as the inner coordinaEon sphere). MulEdentate ligands MulEdentate ligands (‘mulE’ = many, ‘dentate’ = bite) are ligands that have more than one donor atom that can coordinate to the metal centre. For example, the bidentate ligand ethylenediamine (H2NCH2CH2NH2) has two nitrogen atoms that bind to the same metal centre. Ligands with more donor atoms make more stable complexes since more atoms have to be detached to dissociate the complex. Remember that it is the donor atoms that determine the geometry and coordinaEon number of the metal centre, not the ligands. This means that the ethylenediaminetetraac
etato (EDTA) ligand (2 nitrogen and 4 oxygen donor atoms) makes octahedral complexes with CN = 6, even though there is only one ligand. Structural (ConsEtuEonal) Isomers Complexes with the same chemical formula can have different ligand configuraEons, and these different configuraEons are known as isomers. As for organic chemistry, different isomers have different chemical and physical properEes. Ionisa/on and hydra/on isomers are formed by ligands exchanging between the inner and outer coordinaEon spheres. Note that the formula name contains the same ions/
molecules but arranged in a different order. Linkage isomers are formed by ligands coordinaEng by different donor atoms. In the example below, NO2 can coordinate by either the nitrogen atom or one of the oxygen atoms. ionisaEon isomers hydraEon isomers linkage isomers Stereoisomers Geometrical isomers are formed by different arrangements of ligands around the metal centre. Cis-­‐
isomers have the same ligands on the same side of the molecule, trans-­‐isomers have the same ligands on opposite sides of the molecule. cis-­‐plaEn [Pt(NH3)2Cl2] is an important anEcancer drug whereas trans-­‐plaEn is not nearly so acEve and quite toxic. Op/cal isomers are harder to spot without a 3D model. OpEcal isomers are mirror images of each other but, like hands, they cannot be superimposed on each other. Molecules that have opEcal isomers are called chiral molecules. Their physical properEes are idenEcal and their chemical properEes are idenEcal with non-­‐chiral molecules but they react differently with other chiral molecules. Since most biological molecules are chiral, recognising chiral molecules is essenEal for making molecules for biological applicaEons. Naming CoordinaEon Complexes The following rules are a brief summary of the rules for naming coordinaEon complexes. More details and examples appear in the textbook. 1.  CaEons are named before anions with a space between their names. 2.  Ligands in the inner coordinaEon sphere appear in alphabeEcal order. Greek prefixes, di-­‐, tri-­‐, tetra-­‐ etc are used for mulEple ligands. The names of more complicated ligands appear in brackets. 3.  The names of anionic ligands end in –o, e.g. chloro-­‐ (Cl–), hydroxo-­‐ (OH–) and sulfato-­‐ (SO4–). 4.  The names of neutral ligands are usually unchanged but there are a few important excepEons: ammine (NH3), aqua (H2O), carbonyl (CO) and nitrosyl (NO). 5.  Metals that exhibit variable oxidaEon states have their oxidaEon state appear in roman numerals in brackets aAer their name, e.g. cobalt(III) (Co3+), Fe(II) (Fe2+). 6.  If the complex is an anion overall, the suffix –ate appears aAer the name. If the complex is neutral or caEonic overall, there is no suffix. When wriEng the formulae of coordinaEon compounds with uncoordinated counter ions in the outer sphere, the metal centres and inner coordinaEon sphere ligands appear within square brackets. Examples: K2[Cu(CN)4] — potassium tetracyanocuprate(II) [Ag(NH3)2]Cl — diammine silver chloride Ni(CO)4 — tetracarbonylnickel(0) [Cu(NH3)2(en)]Br2 — diammine(ethylenediamine)
copper(II) bromide Theories of CoordinaEon Complexes It was originally thought that all coordinaEon complexes were formed by the ligands donaEng pairs of electrons to empty orbitals on the central metal. The formaEon of these complexes was also considered to be reversible with relaEvely fast exchange between coordinated and uncoordinated ligands. It is now known, however, that there are more complex interacEons between the metal centre and the ligands. Many complexes are extremely stable and there are some cases where the metal donates electrons back to the ligands. It is even possible for a metal centre to have a negaEve formal oxidaEon state (this never happens for simple metal ions). Some complexes have a disEnctly ionic nature while others are highly covalent and even involve covalent bonds between metal centres. For this reason, chemists use a range of models, of increasing complexity, to explain the behaviour of metal complexes. We will briefly cover those models over the next few slides. A simple Lewis acid/base complex with labile ligands. Iron pentacarbonyl involves donaEon from the carbonyl ligands (CO) to Fe and back from Fe to CO. VBT and CoordinaEon Complexes Valence bond theory, the original and simplest explanaEon, imagines that ligands donate pairs of electrons to empty d-­‐orbitals on the metal centres. This theory involves the use of da/ve bonds (daEve bonds are the same as normal covalent bonds but both electrons originate from the same atom). The valence bond descripEon is very simple and so it is oAen used to introduce the concepts of transiEon metal complexes. However, there are major problems when we try to account for all of the electrons. Co3+ can be used as an ideal case for applying valence bond theory. Co3+ has the 4s0 3d6 4p0 configuraEon so we can use the six d-­‐electrons to fill three non-­‐bonding d-­‐orbitals. This leaves one s-­‐, two d-­‐ and three p-­‐orbitals empty for forming six empty sp3d2 orbitals (remember that six atomic orbitals make six hybrid orbitals). The ligands donate one pair of electrons each into these six orbitals to make the complex. However, VBT cannot easily explain why Co2+ can also form octahedral complexes because it has one electron more than Co3+ and the final complex would therefore have one too many electrons for the available orbitals. Furthermore, En 4+ (Sn4+) can also make octahedral complexes which results in a total of 22 electrons — 4 more than the total spaces available. Crystal Field Theory The next most simple model is called crystal field theory (CFT). CFT imagines the metal centre as a point of posiEve charge and the surrounding ligands as points of negaEve charge. EquaEons based on the symmetry of d-­‐orbitals then predict which arrangements of ligands have the lowest energy and are therefore most stable. Note that CFT is a major simplificaEon of the real situaEon — in CFT, the ligands are only points, not molecules. This simplificaEon means we can ignore the problems of electron counEng. CFT is useful because it explains many of the observaEons of transiEon metal complexes. Direct approach of –ve charges towards the lobes of the dx2-­‐y2 orbital raises its energy. The –ve charges miss the lobes of the dxy orbital and do not change its energy. In CFT, we consider what happens to the energy of electrons in the non-­‐
bonding d-­‐orbitals as they are approached by points of negaEve charge along the x-­‐, y-­‐ and z-­‐axes. When the negaEve charges approach the d-­‐orbitals, the energies of electrons in orbitals lying directly on the axis of approach will be raised in energy as a result of repulsion with the negaEve charges. Electrons in orbitals not directly on the line of approach will not be raised in energy. CFT and Real Ligands To be_er understand CFT, it might be helpful to look at real ligands around a metal centre, in this case, NH3 ligands around a Co3+ ion. dx2–y2 dz2 z y x dxy dxz dyz We can see that only the dx2–y2 and dz2 orbitals point directly at the ligands, the other orbitals all point between them. Any electrons in the dx2–y2 or dz2 orbitals will have a higher energy because they are closer to the electron density of the ligands. d-­‐Orbital Splinng CFT is a simple way to understand why d-­‐orbitals are not degenerate in complexes. Because ligands affect different d-­‐orbitals in different ways, the d-­‐subshell splits into two subshells: the t2g and the eg subshells. When we remember that empty, half-­‐full and full sub-­‐shells are more stable than other configuraEons, we can understand why some oxidaEon states are more dx2–y2 dz2 eg stable than others. dxy dxz dyz dx2–y2 dz2 dxy dxz dyz t2g The Spectrochemical Series The energy difference between the t2g and eg subshells is called the crystal field spliAng energy (∆) and is similar to the energy of visible light. This means that most coordinaEon complexes can absorb visible light, which gives them colour. The colour of the complex therefore tells us about the energy difference. The value of ∆ depends eg on the ligands and the hν ∆ geometry of the complex. t2g ∆octahedral : I– < Br– < S2– < Cl– < NO3– < F– < OH– < H2O < NH3< CN– < CO Ligand Field Theory CFT has a serious problem with some neutral ligands such as CO and NH3, for example. Since they do not have a negaEve charge, the splinng energy of these complexes should be quite small but it is actually very large. To explain this, we need a higher level theory called ligand field theory (LFT), which mixes the concepts of molecular orbital theory with crystal field theory. Because it incorporates interacEons between d-­‐orbitals on the metal centre and p-­‐
orbitals or π-­‐bonds on the ligands, it explains the properEes of more transiEon metal complexes. The large number of bonding and anEbonding orbitals generated by MO theory also easily accounts for the large number of electrons needed in coordinaEon complexes. Ligand field theory explains why carbonyl ligands form stable complexes with metal ions — the carbon monoxide ligand can donate electron density from its π-­‐bond to the d-­‐orbitals of the metal centre. If a ligand has empty p orbitals or empty π-­‐anEbonding orbitals, the metal ion can back donate electron density to the ligand. Back donaEon stabilises low oxidaEon states of transiEon metals since the metal centre can unload excess electron density.