This page explores how you write electronic structures for atoms using s, p, and d notation. It assumes that you know about simple atomic orbitals at least as far as the way they are named, and their relative energies. If you want to look at the electronic structures of simple monatomic ions (such as Cl-, Ca2+ and Cr3+), you will find alink at the bottom of the page.
Important! If you haven't already read the page on atomic orbitals you should follow this link before you go any further.
The electronic structures of atoms
Relating orbital filling to the Periodic Table
UK syllabuses for 16 - 18 year olds tend to stop at krypton when it comes to writing electronic structures, but it is possible that you could be asked forstructures for elements up as far as barium. After barium you have to worry about f orbitals as well as s, p and d orbitals - and that's a problem for chemistry at a higher level. It is important that you look through past exam papers as well as your syllabus so that you can judge how hard the questions are likely to get. This page looks in detail at the elements in the shortened version of thePeriodic Table above, and then shows how you could work out the structures of some bigger atoms.
http://www.chemguide.co.uk/atoms/properties/elstructs.html (1 of 9)07/06/2011 04:40:59 p.m.
electronic structures of atoms
Important! You must have a copy of your syllabus and copies of recent exam papers. If you are studying a UK-based syllabus and haven't got them, follow this link to find outhow to get hold of them.
The first period Hydrogen has its only electron in the 1s orbital - 1s1, and at helium the first level is completely full - 1s2. The second period Now we need to start filling the second level, and hence start the second period. Lithium's electron goes into the 2s orbital because that has a lower energy than the 2p orbitals. Lithium has an electronic structure of1s22s1. Beryllium adds a second electron to this same level - 1s22s2. Now the 2p levels start to fill. These levels all have the same energy, and so the electrons go in singly at first. B C N 1s22s22px1 1s22s22px12py1 1s22s22px12py12pz1
Note: The orbitals where something new is happening are shown in bold type. You wouldn't normally write them any differently from the other orbitals.http://www.chemguide.co.uk/atoms/properties/elstructs.html (2 of 9)07/06/2011 04:40:59 p.m.
electronic structures of atoms
The next electrons to go in will have to pair up with those already there. O F Ne 1s22s22px22py12pz1 1s22s22px22py22pz1 1s22s22px22py22pz2
You can see that it is going to get progressively tedious to write the full electronic structures of atoms as the number of electronsincreases. There are two ways around this, and you must be familiar with both. Shortcut 1: All the various p electrons can be lumped together. For example, fluorine could be written as 1s22s22p5, and neon as 1s22s22p6. This is what is normally done if the electrons are in an inner layer. If the electrons are in the bonding level (those on the outside of the atom), they are sometimes written in shorthand,sometimes in full. Don't worry about this. Be prepared to meet either version, but if you are asked for the electronic structure of something in an exam, write it out in full showing all the px, py and pz orbitals in the outer level separately. For example, although we haven't yet met the electronic structure of chlorine, you could write it as 1s22s22p63s23px23py23pz1. Notice that the 2pelectrons are all lumped together whereas the 3p ones are shown in full. The logic is that the 3p electrons will be involved in bonding because they are on the outside of the atom, whereas the 2p electrons are buried deep in the atom and aren't really of any interest. Shortcut 2: You can lump all the inner electrons together using, for example, the symbol [Ne]. In this context, [Ne] means the electronic...