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Chapter 22: Co-ordination Complexes Ligands and Co-ordination Complexes The first experiment in the Chemistry 2000 lab is synthesis and analysis of a complex salt. The name of this complex salt is potassium trisoxalatoferrate(III) trihydrate. Its formula is K3[Fe(C2O4)3]·3H2O and its structure is shown below: O C O O K3 O C O C O O C O + Fe O O C C 3 H2O O O Co-ordination complexes are compounds in which several ligands are co-ordinated to a transition metal cation. A ligand is any substance (neutral or anionic) which can act as a Lewis base, donating electrons to the transition metal (which acts as a Lewis acid). e.g. Cu(NH3)42+ is Cu2+ with four :NH3 ligands e.g. Zn(CN)42- is Zn2+ with four :CN- ligands Note that the ligands do not have to be the same! Ligands co-ordinated to a transition metal though one atom are called monodentate ligands. Those co-ordinated to a transition metal through two atoms (like in the picture above) are called bidentate (“two-toothed”) ligands. Polydentate ligands can also be called chelating ligands, or chelates (“claws”). O We saw one such ligand in the Chemistry 1000 Water Hardness lab. EDTA was able to “grip” a cation by co-ordinating to it with six different atoms! (For clarity, individual carbon atoms are not shown.) O N O M O O N O O The number of atoms attached to the transition metal is referred to as the co-ordination number. It doesn’t matter whether these atoms come from the same molecule/ion or from several different ones. e.g. The co-ordination number is 6 for both Fe(C2O4)33- and Fe(OH2)63+ Complex Salts As shown on the previous page, co-ordination complexes can have positive charges, negative charges or be neutral. To make a neutral salt, charged co-ordination complexes (also called complex ions) will need one or more counter ions to balance the charge. This gives a complex salt. e.g. In K3[Fe(C2O4)3]·3H2O, the negative charge of Fe(C2O4)33is balanced out by the three K+ cations. O Some co-ordination complexes and complex salts contain extra water molecules which were trapped during crystallization. These complexes are also hydrates. (Recall Chemistry 1000!) Thus, a co-ordination complex must contain a transition metal cation and several ligands. It may also have counter ion(s) (to balance charge) or extra water molecules. When naming a coordination complex or complex salt, look for these components. Naming Complex Salts The first step in naming a complex salt is to identify the complex ion. To name the complex ion: 1. List the ligands using prefixes to indicate the number of each type of ligand. • For ligands with simple names (e.g. chloro, hydroxo), use di, tri, tetra, etc. • For ligands with complicated names (e.g. oxalato), use bis, tris, and tetrakis. 2. Name the transition metal. If the complex ion is an anion, use the metal’s Latin name and change the suffix to ‘ate’ 3. List the oxidation state using Roman numerals. Once you have named the complex ion, name the complex salt like any other ionic compound: cation then anion then hydration. e.g. Fe(C2O4)33- = 3 C2O42+ Fe3+ K3[Fe(C2O4)3]·3H2O potassium trisoxalatoferrate(III) trihydrate (cation) (complex anion) (hydration) Table of Common Ligands Anions Formula fluoride anion :Fchloride anion :Clnitrite anion :NO2:ONOcarbonate anion :OCO222OO oxalate anion [ cyanide anion thiocyanate anion :OCCO: :CN:SCN:NCS:H:O2:OH- hydride anion oxide anion hydroxide anion Neutral Molecules water :OH2 ammonia :NH3 carbon monoxide :CO nitrogen monoxide :NO Name the following complex salts (a) [Ni(OH2)6]CO3 (b) [Cu(NH3)4]SO4 · H2O ] Name fluoro chloro nitro nitrito carbonato oxalato cyano thiocyano isothiocyano hydrido oxido hydroxo aqua ammine carbonyl nitrosyl Note that there is a difference between water as a ligand and “water of crystallization”. The bright blue crystals commonly referred to as CuSO4·5H2O are really [Cu(OH2)4]SO4·H2O. Give the name corresponding to each of these two formulas. CuSO4·5H2O = [Cu(OH2)4]SO4·H2O = The only way to determine this information is by experiment, but you should recognize that, in many hydrated salts, at least some of the water molecules serve as ligands. Why Are Transition Metals Special? As we saw in Chemistry 1000, metals in Groups 1 and 2 are limited in what oxidation states they can take on. Transition metals, on the other hand, can take on many different oxidation states. This distribution is not entirely random, as show in the graph below (with common oxidation states in dark red): The elements in the middle can exist in a wider variety of oxidation states than those on either end of the d-block. Why? Because of the valence d electrons! Compared to s and p electrons, d electrons can be added or removed relatively easily. e.g. The electron configuration of neutral vanadium is: The first two electrons removed will be those in the 4s orbital. After that, electrons are removed from the 3d orbitals giving three stable oxidation states: vanadium(III) vanadium(IV) vanadium(V) Electronic Structure and Colour One of the more fun consequences of these partially filled d subshells is that the co-ordination complexes of transition metals are often brightly coloured. The flasks below contain aqueous solutions of several nitrate salts. Note that, since all nitrates are water-soluble, these solutions contain transition metal-water complexes. Fe3+ Co2+ Ni2+ Cu2+ Why is the Zn2+ complex the only colourless one? Zn2+ Consider the electron configurations of the five cations: Fe3+ Co2+ Ni2+ Cu2+ Zn2+ The colourless Zn2+solution is the only one shown containing a cation with a full d subshell! Where does the variety in colour come from? Most co-ordination complexes have octahedral geometry. This means that two of the d orbitals point directly at ligands while the other three do not. point at ligands point between ligands A simple electrostatic model, called the crystal field theory, assumes that there will be a certain degree of electron-electron repulsion between the electron pair a ligand donates and any electrons already in the metal d orbitals. This repulsion is felt most strongly by electrons in d orbitals pointing at the ligands. Thus, the dz2 and dx2-y2 orbitals are pushed to higher energy than the dxy, dxz and dyz orbitals. This separation in energy is referred to as crystal field splitting (∆o where ‘o’ is for ‘octahedral’). How does this make for coloured solutions? Recall that photons are emitted when electrons drop from a higher energy orbital to a lower energy orbital. (see Atomic Line Spectra in Chemistry 1000) Similarly, the electrons get to the higher energy orbital by absorbing photons of light. In co-ordination complexes with crystal field splitting, there are two ways to distribute d electrons. The high spin distribution maximizes the spin pairing of the d electrons while the low spin distribution puts electrons in the lowest energy orbitals first. When the atom (of the ligand) donating the electron pair is oxygen, the d electrons are always distributed high spin. Electrons in the lower energy d orbitals can absorb photons and be excited into the higher energy d orbitals. Since ∆o corresponds to the energy of light in the visible region (and there is more than one way to absorb a photon), some wavelengths of visible light are absorbed. The wavelengths that are not absorbed give the colour of solution. e.g. Ni2+ in Ni(OH2)62+ absorbs both red-yellow and violet light, giving a solution that appears green: Note that different ligands provide different amounts of crystal field splitting. Fe(OH2)63+ and Fe(C2O4)33- are both complexes of Fe3+ but Fe(OH2)63+ is red-orange while Fe(C2O4)33- is green. A spectrophotometer measures the amount of light absorbed by a complex. When analyzing a green complex, it is therefore necessary to look at the absorption of light other than green. Generally, we will try to choose the wavelengths most strongly absorbed by the complex. On the figure above, that would correspond to the peak yellow-red wavelength or the peak violet wavelength. Important Concepts from Chapter 22 • Lewis acids and Lewis bases • ligands (monodentate, bidentate, chelating) • co-ordination number • naming co-ordination complexes and complex salts • electron configurations of cations • d electrons and crystal field splitting • why co-ordination complexes are often coloured • spectrophotometry (especially for labs!)