Does 5s or 4d orbital fill first?

While writing out the electronic configurations, we usually write them in the order of 1s 2s 2p 3s 3p 3d 4s 4p 4d 4f 5s 5p 5d 5f and so on. This method of writing the electronic configuration only represents the traditional increasing order of shells ( 1 2 3 4 or K L M N ), and nested within each shell are the subshells ( s p d f).

This is only the usual method of writing the electronic configurations; it does not reflect the order of increasing energy of orbitals. The order of energy is given by the Aufbau's Principle or the building-up priciple.

In essence, according to the rule, for comparing energies of orbitals of all elements, except hydrogen, we compare the values of the sum of (n+l) ; where n is the priciple quantum number ( corresponding to K=1 L=2 N=3 etc. ) and l is the azimuthal quantum number ( corresponding to s=0 p=1 d=2 etc. ). Greater the value of this sum, more is the (relative) energy associated with that orbit. Moreover, if this sum is same for two orbitals, then we compare the value of n for those orbitals.

So, for example, you mentioned the anomaly in the energy order of 4s and 3d. Using this rule, for 4s we have n=4 & l=0, giving a sum of 4, while for 3d we have n=3 & l=2, giving a sum of 5. This gives us a greater energy for 3d, which shows us that it confirms to the rule rather than being anomalous.

Further, consider the case of 4s and 3p orbitals; the n+l sum is same for both (4 in each case), but since 4s has a greater value of n than 3p , it is more energetic.

Now, electronic configurations are important in the order 1s 2s 2p 3s 3p 3d 4s 4p 4d 4f 5s 5p 5d 5f when we are considering cases of ionization ; so, if you want to make a +2 ion, of say, Fe, then the two electron will be removed from the outermost orbitals, i.e. 4s. However, in cases where we are concerned with the actual energies, we must subscribe to the Aufbau's Principle.

The electron configuration of an element is a list of the atomic orbitals which are occupied by electrons, and how many electrons are in each of those orbitals.  The rules for writing electron configurations are known as the aufbau principle (German for "building up"):

• Each electron that is added to an atom is placed in the lowest-energy orbital that is available.  The orbitals are filled in the order:

1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f

• Each orbital can hold no more than two electrons.  Two electrons in the same orbital must have opposite spins (the Pauli exclusion principle).

• If two or more orbitals are available at the same energy level (degenerate orbitals), one electron is placed in each orbital until the available orbitals are occupied by one electron; any additional electrons are placed in the half-filled orbitals.

Electron configurations are written as a list of orbitals which are occupied, followed by a superscript to indicate how many electrons are in those orbitals.

H1s1He1s2Li1s2 2s1Be1s2 2s2B1s2 2s2 2p1C1s2 2s2 2p2N1s2 2s2 2p3O1s2 2s2 2p3F1s2 2s2 2p4Ne1s2 2s2 2p5Na1s2 2s2 2p6 3s1

Electron configurations in which all of the electrons are in their lowest-energy configurations are known as ground state configurations.  If an electron absorbs energy, it can move into a higher-energy orbital, producing an excited state configuration.

For atoms with a large number of electrons, the complete electron can be very cumbersome, and not very informative.  For instance, the complete configuration of the element radium is

Ra:  1s2 2s2 2p6 3s2 3p6 3d10 4s2 4p6 4d10 4f14 5s2 5p6 5d10 6s2 6p6 7s2

(With a description like that, you'd be radioactive too!)  Since everything up to the 6p6 is the same electron configuration as the noble gas radon, the configuration can be abbreviated as

Ra:  [Rn] 7s2

Abbreviated electron configurations are always based on the nearest preceding noble gas.

Electron configurations can be written directly from the periodic table, without having to memorize the aufbau scheme, using the following patterns:

Half-filled and filled subshells are especially stable, leading to some anomalous electron configurations:

Predicted configurationActual configurationCr[Ar] 3d4 4s2[Ar] 3d5 4s1Cu[Ar] 3d9 4s2[Ar] 3d10 4s1Ag[Kr] 4d9 5s2[Kr] 4d10 5s1Au[Xe] 4f14 5d9 6s2[Xe] 4f14 5d10 6s1

In the case of chromium, an electron from the 4s orbital moves into a 3d orbital, allowing each of the five 3d orbitals to have one electron, making a half-filled set of orbitals.  In the case of copper, silver and gold, an electron from the highest-occupied s orbital moves into the d orbitals, thus filling the d subshell.  Many anomalous electron configurations occur in the heavier transition metals and inner transition metals, where the energy differences between the s, d, and f subshells is very small.

Which will be filled first 5s or 4d subshell and why?

Which orbital has more energy 4d or 5s?

Do d orbitals fill before s?

Which orbital should be filled first?