Electron Counting. It s easier to show than to explain, I think.

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1 Electron Counting This, for some reason I ve never worked out, isn t as easy as it should be. It is, however, important. It allows us to say things about the stability of organometallic complexes and make predictions about the structures of complexes and deduce what oxidation state a metal centre is in. All useful information. Effectively, electron counting is a method for determining, by visual inspection, how many electrons a given complex has in its molecular orbitals (its covalent bonding system). There are certain magic numbers (which I ll go through below) that tell you your complex is happy. There are two schools as to how to go about determining this, the ionic method and the covalent method. Both must give the same answer, so although you only really need to ever use one method, using the other is a good way to double-check your answer. It s easier to show than to explain, I think. Firstly, using both ways, you need to assess which ligands are dative and which aren t. This partly comes with experience, but if you re not confident you can work it out by looking at the ligand by itself and seeing if the binding atom conforms to the octet rule (remember that?). If it s got eight electrons in its outer shell, contributed from itself and from any other atoms bound to it, it ll act as a dative ligand; if it s got seven or fewer, it ll be negatively charged.

2 Then, using both methods in parallel: ionic: covalent: So, in essence, in the ionic method, you determine which bonds naturally break into ionic pairs (in this, the Pd Aryl and Pd I bonds) and which are purely dative (donation from the phosphorous lone pair). This gives you the charge on the metal centre (positive, 2+). From this, looking at your periodic table, you can determine how many outer shell electrons this metal has (eight), and add this number to the number of electrons the ligands each give (in this case, each gives two electrons). In the covalent method you treat all the bonds as covalent and determine which bonds are dative and which have electrons contributed from both the ligand and the metal. The dative ligands give 2 electrons, whilst the other ligands give one electron. You count how many electrons the metal has assuming it to be uncharged (i.e. exactly as the periodic table has it, in this case 5s 2 4d 8, ten electrons). You do also need to take account of the overall charge of the complex - if we look at Zeise s salt again, you can see how the complex itself is negatively charged as it needs to be balanced by a potassium cation.

3 You only need to count the complex itself, without the counterion. Hence, as another example: ionic: covalent: If you don t quite understand this, there s a very good page online that goes through this again, has a few more worked examples and some opportunities for you to try yourself. It also gives a very useful table which shows many of the most common ligands and how many electrons they donate in the ionic and covalent counting method. See:

4 Task 1 Provide electron-counts (using both the ionic and covalent methods) for the following complexes: [Hint - use the ligand table in the above link to help with any unfamiliar ligands] Eighteen/sixteen electron rule Right, you now know how to count electrons. This is good, though it probably seems pretty useless. It isn t. It s very important in combination with the eighteen and sixteen electron rules. These rules come from the molecular orbital analysis I briefly talked about earlier but didn t actually show. They are what is left when you boil down all the group theory stuff and get to the actually useful information. They re used to say, in essence, that a given complex is stable based on how many electrons it has in its bonding system. The (very straightforward) rules are as follows: The sixteen electron rule: if a complex has a square-planar geometry and has 16 electrons, it is stable. The eighteen electron rule: if a complex has any geometry other than a square-planar geometry and has 18 electrons, it is stable.

5 Easy. Chemistry s at its best when the maths it requires can be done with only fingers and toes. Some common inorganic complex geometries are shown in figure 3. Figure 3 - Common inorganic complex geometries - left to right: square-planar, tetrahedral, octahedral, trigonal bipyramidal, square-based pyramidal Think of the sixteen and eighteen electron rules a bit like the octet rule that you learn at GCSE, except applied to organometallic systems. They re pretty much that simple, but they re very powerful. Just one quick example of their use is in rationalising why some complexes exist in square-planar configurations (e.g. Zeise s salt) when a tetrahedral configuration would allow them to spread their ligands out more (more stable in a purely ionic compound). These rules also allow you to determine what oxidation state the metal at the centre of an organometallic complex is in. Just assume the complex to have eighteen electrons (or sixteen if it s square planar), use the ionic method to strip away the electrons contributed by the ligands, account for the overall charge, then work out how many electrons the metal has to have and how many fewer (or even more) that number is than the periodic table gives. Easy. (Disclaimer - this only works for truly covalent organometallic compounds, the maths doesn t work on some complexes like the titanium complex shown in Figure 1 Task 2 o Why might it be that the above method only works for truly covalent organometallic compounds? o Comment on the stability of all the complexes you ve electron-counted so far in this work sheet.