Chemguide: Support for CIE A level Chemistry
Learning outcome 1.3
Electrons, energy levels and atomic orbitals
Learning outcomes 1.3.1 to 1.3.8
These statements deal with the arrangement of electrons in atoms and ions. I am blocking them together because the Chemguide pages I want you to read aren't arranged in precisely the same way as the syllabus. There are a couple of terms used in the syllabus which I haven't used in these pages, and I will discuss those after you have read the pages. You don't need them before that.
Take your time over this. It is absolutely essential that you fully understand it all before you go any further with your chemistry. There are a lot of other topics which rely on this basic material.
You should read the following pages, following the links from one to the next.
Start with the page about atomic orbitals.
Near the top of that page, you will find a link to another page about the difference between the words "orbit" and "orbital". Follow that link! It is really important that you understand the difference between these words, and then permanently forget any ideas that you ever had about electrons orbiting a nucleus!
Follow the link at the bottom of the atomic orbitals page to the page about writing electronic structures.
At the bottom of that page, there is a link to another page which explains how to write electronic structures for ions .
Don't go on to this last page until you are sure you understand the previous one about the electronic structures of atoms.
The only way of being sure is to practise doing it. Use a Periodic Table to find the proton numbers, and then practise working out the electronic structure for random atoms up to krypton in the Periodic Table, checking your answers against examples given on the first page.
Once you are certain that you can do it for atoms from hydrogen to krypton, try a few examples of bigger atoms in the s and p blocks. Start with iodine and barium so that you can check your answers with the text on the electronic structures page. Then try a few more - you will have to use your initiative to find a way of checking them. (Actually, they aren't hard to find from a Google search!)
Once you are sure you can do this, go on and read, and then practise, the follow-up page about the electronic structures of ions. You will have to restrict yourself to ions whose charges you know.
Important! On the page about the electronic structures of ions, you will find a link to a discussion page about the order of filling of the 3d and 4s orbitals. To keep things simple for yourself, I would suggest that you don't follow this link.
The link challenges the accepted view of the order of filling of these orbitals for the d-block elements. At the time of writing (October 2012), this isn't generally taught at this level. CIE will expect you to use the traditional view that 4s orbitals fill before 3d orbitals.
If you don't read that page, you won't get confused by this.
You will find references to this discussion page at various other places on Chemguide. I would suggest that you ignore it on each occasion. For CIE exam purposes, you need to stick to the generally accepted view.
Learning outcome 1.3.1
This statement talks about shells and sub-shells. I don't tend to use those terms.
A shell of electrons is all of the electrons at a particular level - the 1-level or the 2-level or 3-level or whatever.
The outermost shell is sometimes called the valence shell. This is the shell where the electrons are involved in reactions.
A sub-shell covers the slight differences in energy of various orbitals in each of these levels. So, for example, at the 2-level (the second shell), there are the sub-shells 2s and 2p. At the 3-level (the third shell), there are sub-levels 3s, 3p and 3d.
It also uses the term principle quantum number (n). This is the posh name for the over-all energy level. For example, electrons in the 1-level, have a principle quantum number of 1. In the 3-level they all have the principle quantum number of 3.
When you write electron structures like 1s22s22px1, the big numbers before the s or p are the principle quantum numbers.
Note: Quantum theory gives each electron a total of 4 quantum numbers. The others govern things like the shape of the orbital, the direction (if any) that it points in, and the spin of the electron. No two electrons in an atom can have the same 4 quantum numbers. You don't need to know any of this, except that the principle quantum number gives the over-all energy level of the electron.
The statement also mentions the term ground state of an electron. When we write electronic structures for atoms or ions, electrons always go into the lowest energy levels available. That is the ground state of the atom.
It is, however, possible to excite electrons to higher levels by putting in energy - this would be called an excited state, and the electrons will tend to fall back to their original ground state releasing energy again.
Learning outcome 1.3.9
This statement deals with free radicals.
The syllabus defines free radicals as any species with one or more unpaired electrons. Species is a usefully inclusive term in chemistry which can include atoms, ions and molecules.
One of the free radicals you will come across later in the course is a chlorine radical. A chlorine radical is just an isolated chlorine atom which has 7 electrons in its outer level. Obviously, one of them must be unpaired. In a chlorine molecule, Cl2, all the electrons are now paired.
In the atmosphere, chlorine radicals are responsible for the depletion of the ozone layer. You will also meet chlorine radicals as something produced during the reaction between methane and chlorine.
For now, just recognise that free radicals are any species with one or more unpaired electrons.
© Jim Clark 2019