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2P32 – Principles of Inorganic Chemistry Dr. M. Pilkington Lecture 2 - Introduction to Metal Complexes Metal Complexes: What are they? Why are so many transition metals six-coordinate? Werner’s Coordination Theory. Geometry of six-coordinate complexes, geometric isomers? Assignment 1 due in next Monday at 4.30pm, please write your name and student ID on your work and staple all loose sheets together. Transition Metals – located in the d-block of the periodic table alkali metals C, H, N, O, halogens s-block p-block transition metals d-block d-block their d orbitals are filling focus for coordination chem. (f-block) inner transition metals Rows are across are called Periods. Columns down are called Groups Transition Elements begin in the 4th period K Sc Ti V Cr Mn Fe Co Ni Cu Zn Ga Ca 1 2 [Ar]4s [Ar]4s d-block 1 Ground State Electronic Configuration for the first row Transition Metals Sc - Scandium Ti – Titanium V-Vanadium Cr – Chromium Mn – Manganes e Fe – Iron Co – Cobalt Ni – Nickel Cu – Copper Zn – Zinc 10 Ni Pd Pt [Ar] 3d14s2 [Ar] 3d2 4s2 [Ar] 3d3 4s2 [Ar] 3d5 4s1 [Ar] 3d5 4s2 [Ar] 3d6 4s2 [Ar] 3d7 4s2 [Ar] 3d8 4s2 [Ar] 3d10 4s1 [Ar] 3d10 4s2 11 Cu Ag Au 12 Zn Cd Hg Pd – Palladium Pt – Platinum Ag – Silver (Argentum) Au – Gold (Aurum) Cd – Cadmium Hg – Mercury (hydrargyram) “liquid silver” Most transition metals have several oxidation states. Mn exists in 11 oxidation states -3 upto +7 Transition metals and their compounds are usually brightly colored. 2 “Metal Complexes” (coordination compounds) Origin of the name originates from the 1800’s, it was known metal had “valencies” (oxidation numbers) that could be satisfied by combination with elements having opposite valencies. Cr3+ valence of +3 O2- valence of -2 Cl- valence of -1 Examples of metal complexes CrCl3, Cr2O3 BUT CrCl3 reacts with ammonia (NH3) to form a new compound. CrCl3 + 6NH3 → CrCl3.6NH3 Called a complex because nobody at the time understood how they formed. Alfred Werner – late 1800’s the father of coordination chemistry. Studied in Switzerland at the University of Zurich. He lectured in both organic and inorganic chemistry. He developed the theory of coordination chemistry. He prepared and studied coordination compounds and discovered optically active forms of 6-coordinate octahedral complexes. His coordination chemistry extended through a whole range of systematic inorganic chemistry and into organic chemistry and he was awarded the Nobel Prize in Chemistry in 1913. Nobel Lecture http://nobelprize.org/chemistry/laureates/1913/werner-lecture.html 3 Werner studied the following metal complexes: CoCl3 forms four different compounds with NH3. CoCl3.6NH3 AgNO3 (excess) (1:6 mol ratio) Yellow CoCl3.5NH3 Lavender AgNO3 (excess) 3AgCl (ppt) 2AgCl (ppt) One Cl does not react CoCl3.4NH3 Green CoCl3.4NH3 Violet AgNO3 (excess) 1AgCl (ppt) AgNO3 (excess) 1AgCl (ppt) Werner’s Conclusions: 1. 2. 3. 4. 5. In this series of compounds, cobalt has a constant coordination number of 6 (coordination number is the number of groups that can bond directly to the metal). As the NH3 molecules are removed they are replaced by Cl- which acts as if it is covalently bonded to cobalt. Chloride and Ammonia are now called ligands. Ligands are a Lewis base/electron pair donors that can bind to a metal ion. A metal complex – metal ion combined with ligands. 4 6. Coordination complexes are neutral and counter ions are not bonded to the central metal ion but balance the charge. For example: +3 0 -3 [Co(NH3)6]Cl3 counter ions The ligands directly coordinated to the metal are contained within the square bracket. Six NH3 bonded to Co. 3 chloride ions are not bonded to the Co these are counter ions, they balance the charge (Co3+) they are “free” to react with AgNO3 to give 3 moles of AgCl. React with 3 moles of AgNO3 H2O [Co(NH3)6]Cl3 → Yellow [Co(NH3)6]3+ + 3Cl- React with 2 moles of AgNO3 [Co(NH3)5Cl]Cl2 rewrite as [Co(NH3)5Cl]2+ + 2ClLavender (now only two reactive Cl-’s). React with 1 mole of AgNO3 [Co(NH3)4Cl2]Cl rewrite as [Co(NH3)4Cl2]+ + ClTwo isomers green and violet Isomers - have the same formula but different structures, i.e. different spatial arrangements 5 What is the geometry of [Co(NH3)4Cl2]Cl ? Consider Hexagonal Planar Cl Cl Cl H3N H3N Co Co H3N NH3 Cl NH3 Co Cl H3N NH3 H3N NH3 NH3 1,2 isomer ortho 1,3 isomer meta H3N NH3 Cl 1,4 isomer para There are three possibilities so this does not fit with the observations Consider Trigonal Prism Cl Cl Co Cl Cl or Co or Co Cl Cl There are three possibilities so this does not fit with the observations 6 Consider Octahedral http://www.iumsc.indiana.edu/morphology/symmetry/octahedral.html 6 vertices, 8 sides M Metal ion in the center, ligands are on the vertices, all six vertices are identical. Geometry of [Co(NH3)4Cl2]Cl is Octahedral Cl Co Cl 0 180 or 900 Cl Co Cl trans isomer the two Cl ligands are far apart (1800) cis isomer - the two Cl ligands are close to each other (900) Many other examples of complexes of this coordination geometry known. This geometry reduces the steric crowding that is a problem in other geometries and makes them unfavourable. We accept Werner’s conclusions, today further evidence to confirm his conclusions is provided by X-ray crystallography. 7 Why is the coordination number 6 so common? We have to consider the sizes (ionic radii) of the metal ions and ligands. Plot the ionic radii of transition metal ions, most of them are in the range 75-90 pm which can accommodate 6 ligands and hence favours 6-coordination. Consider the size of common ligands such as O as in H2O, N as in NH3 or S. For example: O2- 126 pm 126 pm 126 pm = radius of O (1 pm = 10-12 m) Radius Ratio Rules The structures of many crystals can be rationalised to a first approximation by considering the relative sizes and numbers of ions present. The radius ratio r+/r- can be used to make a first guess at the likely coordination number and geometry around the cation using a set of simple rules: Value of r+/r- Predicted Coordination Number of Cation Predicted Coordination Geometry of Cation < 0.15 2 Linear 0.15-0.22 3 Trigonal Planar 0.22-0.41 4 Tetrahedral 0.41-0.73 6 Octahedral > 0.73 8 Cubic 8 To Summarize: For [M(H20)6]n+ in order for the metal ion to accommodate six ligands it must have at least a 52 pm radius. Six waters fit exactly around a metal ion with 52 pm radius (not crowded). At 92 pm’s the six waters have moved far enough apart that more waters can fit. If the metal ion is smaller than 52pm you will fit fewer H2O’s, if it is larger then you have more space and can go to a larger coordination number. Hence transition metal ions display a number of coordination numbers but 6 is very common. N and O donors are the most abundant in biology (amino acids), S is also present but it is allot larger than O and as a consequence Fe fits 6 O’s but only 4 S’s. This is apparent in thiocluster compounds of Fe. 9