Yes, we lied to you at GCSE - electrons don't really orbit the nucleus! Why did we do it? Well, imagine we taught you about what really happens in year 10 we probably would've brought you to tears!
The good news is that there are still shells; but these shells are divided into sub-shells which contain one, three or five orbitals (at A Level) and they don't follow the 2,8,8 rule. Described simply, an orbital is a region of space within an atom which can hold up to two electrons of opposite spin. From a mathematical perspective, the total probability of finding one or more electrons anywhere within a specific orbital is always equal to one.
What is spin I hear you ask? Spin is a property of an electron that allows it to pair with another of opposite spin - i.e. an electron with 'up-spin' can pair with one with 'down-spin'. It doesn't play that big part at A Level, so there's no real need to go into it beyond what has just been said.
There are three different types of orbital which you need to be aware of at A Level:
- s orbitals - spherical in shape, with an area of zero electron density at the centre (where you'd find the nucleus). s orbitals are so named after their historical name; 'sharp'.
- p orbitals - comprised of two separate balloon-shaped lobes, with an area of zero electron density in the centre (again, where you'd find the nucleus). These orbitals sit along the x, y and z axes. p orbitals are so named after their historical name; 'principal'.
- d orbitals - each is slightly different in shape or orientation, but all with an area of zero electron density at the centre (...where you'd find the nucleus). d orbitals are so named after their historical name; 'diffuse'.
At A Level, most exam boards will expect you to be able to draw s and p orbitals (with axes), but not d orbitals - so don't worry about them too much.
As previously mentioned, orbitals are arranged into sub-shells. where each type of sub-shell is comprised of a specific number of certain types of orbitals:
- s sub-shell; contains one s orbital, which can accommodate up to two electrons of opposite spin.
- p sub-shell; contains three p orbitals, each accommodating up to two electrons each, so six in total.
- d sub-shell; contains five d orbitals, each capable of accommodating up to two electrons each (of opposite spin) and hence a maximum occupancy of ten electrons.
The principal quantum number, often referred to as 'n', denotes the shell number. If an element has a principle quantum number of 1, that means that its valence (outermost) electrons are within the first shell. Each shell can hold a different number of electrons (maximum occupancy):
- n = 1; first shell - contains 1 x s sub-shell (2) and therefore the maximum occupancy is 2.
- n = 2; second shell - contains 1 x s sub-shell (2) and 1 x p sub-shell (6), therefore the maximum occupancy is 8.
- n = 3; third shell - contains 1 x s sub-shell (2), 1 x p sub-shell (6) and 1 x d sub-shell (10), therefore the maximum occupancy is 18.
With regards to the fourth shell (n = 4), you only need to be aware of the s sub-shell within it at A Level.
The orbitals inside an atom get progressively larger, so an s orbital in the third shell (3s) is larger than an s orbital in the first shell (1s). The sub-shells are arranged in terms of increasing energy, and rather bizarrely, the s sub-shell in the fourth shell (4s) actually has a slightly lower energy than the d sub-shell in the third shell (3d), as illustrated below:
So, that summarises the way in which electrons are arranged at A Level. In our next blog, we will start to unpick how these orbitals are filled, the common way to represent electron configurations, how to abbreviate them and the exceptions to the rules.
