Only one lone pair is shown on each water molecule. By making the maximum number of bonds, it releases most energy and so becomes most energetically stable. Six is the maximum number of water molecules it is possible to fit around an aluminium ion (and most other metal ions). You might wonder why it chooses to use six orbitals rather than four or eight or whatever. It re-organises (hybridises) the 3s, the three 3p, and two of the 3d orbitals to produce six new orbitals all with the same energy. The aluminium uses six of these to accept lone pairs from six water molecules. That means that all the 3-level orbitals are now empty. Explore the atomic properties menu, following the links to atomic orbitals and electronic structures.Ĭome back to this page later using the BACK button, the History file, or the Go menu on your browser. Warning: It is a complete waste of time going any further with this page if you aren't confident about writing electronic structures in this form for elements and ions (including the first transition series). When it forms an Al 3+ ion it loses the 3-level electrons to leave Start by thinking about the structure of a naked aluminium ion before the water molecules bond to it. We are going to look in detail at the bonding in the complex ion formed when water molecules attach themselves to an aluminium ion to give Al(H 2O) 6 3+. It isn't, however, particularly important to the rest of this page that you know anything more than the fact that a substance which forms a co-ordinate bond by donating a lone pair of electrons to something else is known as a Lewis base. Note: If you haven't come across the term Lewis base, and want to find out more, you could follow this link to a page on theories of acids and bases. In other words, all ligands function as Lewis bases.
These are used to form co-ordinate bonds with the metal ion.Īll ligands are lone pair donors. What all these have got in common is active lone pairs of electrons in the outer energy level. Simple ligands include water, ammonia and chloride ions. The molecules or ions surrounding the central metal ion are called ligands. If you follow this link, use the BACK button on your browser to return quickly to this page. Amongst other examples of co-ordinate bonding, that page contains a description of the bonding in the complex ion formed between aluminium ions and water molecules, and that will be repeated below on this page - so you needn't spend a lot of time reading that bit. Note: If you aren't sure about co-ordinate (dative covalent) bonding, you aren't going to make much sense of what follows without first following this link. (In some cases, the bonding is actually more complicated than that.)
These can be considered to be attached to the central ion by co-ordinate (dative covalent) bonds. It discusses various sorts of ligand (including some quite complicated ones), and describes what is meant by co-ordination number.Ĭomplex metal ions containing simple ligandsĪ complex ion has a metal ion at its centre with a number of other molecules or ions surrounding it. This page explains the terms complex ion and ligand, and looks at the bonding between the ligands and the central metal ion. Introducing complex ions - ligands and bonding
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