Electron Configurations Using Subshell Notation
All electrons in an atom are placed into shells according to electron energies. These shells are labeled with the integers, starting from one up to infinity, with one being the lowest energy. While there are an infinite number of electron configurations for any atom, a method of describing an atom or an ion is based on the lowest energy state called the ground state (any other configuration with electron(s) in higher energy states are called excited states). The ground state is determined by always placing each electron into the lowest energy subshell.
The order for filling subshells is (sometimes called the Aufbau filling order):
1s 2s 2p 3s 3p 4s 3d 4p 5s 4d 5p 6s 4f 5d 6p 7s 5f 6d 7p
Each subshell has a maximum number of electrons that it can hold. The s-type subshells can hold a maximum of two electrons, the p-type subshells can hold a maximum of six electrons, the d-type subshells can hold a maximum of ten electrons, and the f-type subshells can hold a maximum of fourteen electrons.
To write an electron configuration using the subshell notation, a combination of the subshell followed by a superscript indicating the number of electrons in that subshell is used. Thus, for the first two elements, we would write their electron configurations as:
H: 1s1 and He: 1s2
For the next element, Li, we can’t put a third electron into the 1s subshell because it is full. Thus, we would need to got to the next available subshell – the 2s.
The filling of subshells would continue to build upon the previous element and fill subshells completely before going on to the next subshell. You can see that, however, this would get to be quite a chore when we reach larger elements like lead (Pb). Thus, it is preferential to use a shorthand method that utilizes the configuration of the noble gases (group 8A). To do a shorthand configuration for any element, count backwards from that element until you reach a noble gas. Write that element in brackets. Then, continue forward with next subshell(s) – see the attached version of the periodic chart that shows the subshell order with respect to the elements.
For example, if we wanted to do the shorthand configuration for sodium (Na), you would count back one element to neon (Ne). Put this element symbol in brackets and then, noting that the next correct subshell is 3s, include the rest of the electrons as we did with the smaller elements.
Na: [Ne] 3s1
One of the reasons for doing this is also to distinguish between core and valence electrons. Core electrons are held very tightly by the atom and do not interact when bonds are formed to make compounds. These include the noble gas configuration plus any completely filled d-subshells. The valence, or outer shell, electrons are the ones that interact with each other when bonds formed are therefore very important. These are always the outermost shell of electrons plus any unfilled d-subshell. Note that for any main group element, the number of valence electrons always equals the group number! For sodium, it has ten core electrons (neon has ten electrons) and one valence electron (it is in group 1A).
SUBSHELLS and the PERIODIC CHART