Electron configuration

Heisenberg's Uncertainty Principle basically states that we cannot determine definitely where an electron is and where it is heading to, ie we cannot define the position and momentum of an electron with 100% certainty. Further information can be found on the web, including: here and here, which also goes into energy levels.

Electron simulation

RECAP

At a more general level, and in introductory chemistry you would have learnt that electrons fill shells and atoms are only stable when they have a full shell. The first shell can contain 2 electrons, the second shell can contain 8 electrons, as can the third shell. This then became an 'untalked' about issue for the fourth shell. For example:

electron configuration

At this stage (NCEA L3 Chemistry) we need to realise that electrons do not travel in orbits (see Heisenberg Uncertainty Princple) like planets; instead they fill energy levels, which equate to the number of shells around the nucleus, ie the period the element is in. Even more importantly these areas that electrons fill contain sub orbitals, or areas that electrons are likely to be found in. We notate these sub levels with a main energy level (shell) with the letter s, p, d and f. At Level 3 NCEA we do not touch on the shapes of these sub orbitals (for want of a better phrase) and we do not look at f sub orbitals, however here are the key points:

Orbital shapes

For a more in depth discussion on electron orbitals go here

The table below shows a selection of the elements electron arrangements, from the animation above.

elements in period 3 and above

Elements in period 3 and above can promote electrons and fill d orbitals . Elements in period 4 after calcium have electrons in the 3d orbitals in their ground state. It is important to remember here that electrons fill into the 3d orbitals before the 4s and when the transition metals form ions they lose electrons out of the 4s orbit first.

Another important point to remember, thinking about Hund's Law, is that the most stable state, with maximum repulsion and charge separation is when a set of orbitals are full, followed by a half full orbitals. This will be made clearer in due course.

The slide show shows the electron configurations for Titanium and some of Titanium's ions as well as Chromium and the Chromium (III) ion.

COPPER

Copper is also an element, like Chromium, that you need to be aware of regarding its electron arrangement. Copper has an atomic number of 29, so has 9 valence electrons. If you just count across the periodic table and fill in electrons the electron arrangement would look like:

However bearing in mind full orbits are the most stable configuration, followed by half full orbits this is not the most stable electron arrangement. It is more stable if an electron is promoted from the 4s to the 3d, giving the following, correct electron configuration:

When Cu forms ions, like all transition metals it will lose electrons out of the 4s orbits first. Below are the electron configurations for the Cu(I) and Cu(II) ions:

Below are two videos. One details the electron configuration the of Copper (II) ion. The second shows the electron configurations of some other ions.