Electron configuration
In certain conditions, electrons are able to move from one configuration to another by the emission or absorption of a quantum of energy, in the form of a photon.
[2] Exhaustive technical details about the complete quantum mechanical theory of atomic spectra and structure can be found and studied in the basic book of Robert D.
The first notation follows the order based on the Madelung rule for the configurations of neutral atoms; 4s is filled before 3d in the sequence Ar, K, Ca, Sc, Ti.
In a sodium-vapor lamp for example, sodium atoms are excited to the 3p level by an electrical discharge, and return to the ground state by emitting yellow light of wavelength 589 nm.
Usually, the excitation of valence electrons (such as 3s for sodium) involves energies corresponding to photons of visible or ultraviolet light.
The excitation of core electrons is possible, but requires much higher energies, generally corresponding to X-ray photons.
Niels Bohr (1923) incorporated Langmuir's model that the periodicity in the properties of the elements might be explained by the electronic structure of the atom.
[8] His proposals were based on the then current Bohr model of the atom, in which the electron shells were orbits at a fixed distance from the nucleus.
The vast store of knowledge of chemical properties and relationships, such as is summarized by the Periodic Table, should serve as a better foundation for a theory of atomic structure than the relatively meager experimental data along purely physical lines...
[10] However neither Bohr's system nor Stoner's could correctly describe the changes in atomic spectra in a magnetic field (the Zeeman effect).
Pauli hypothesized successfully that the Zeeman effect can be explained as depending only on the response of the outermost (i.e., valence) electrons of the atom.
This rule was first stated by Charles Janet in 1929, rediscovered by Erwin Madelung in 1936,[12] and later given a theoretical justification by V. M. Klechkowski:[14] This gives the following order for filling the orbitals: In this list the subshells in parentheses are not occupied in the ground state of the heaviest atom now known (Og, Z = 118).
In a hydrogen-like atom, which only has one electron, the s-orbital and the p-orbitals of the same shell have exactly the same energy, to a very good approximation in the absence of external electromagnetic fields.
(However, in a real hydrogen atom, the energy levels are slightly split by the magnetic field of the nucleus, and by the quantum electrodynamic effects of the Lamb shift.)
After calcium, most neutral atoms in the first series of transition metals (scandium through zinc) have configurations with two 4s electrons, but there are two exceptions.
In this case, the usual explanation is that "half-filled or completely filled subshells are particularly stable arrangements of electrons".
[15] The apparent paradox arises when electrons are removed from the transition metal atoms to form ions.
[16] More recently Scerri has argued that contrary to what is stated in the vast majority of sources including the title of his previous article on the subject, 3d orbitals rather than 4s are in fact preferentially occupied.
Indeed, visible light is already enough to excite electrons in most transition metals, and they often continuously "flow" through different configurations when that happens (copper and its group are an exception).
There are several more exceptions to Madelung's rule among the heavier elements, and as atomic number increases it becomes more and more difficult to find simple explanations such as the stability of half-filled subshells.
It is possible to predict most of the exceptions by Hartree–Fock calculations,[21] which are an approximate method for taking account of the effect of the other electrons on orbital energies.
In general, these relativistic effects[23] tend to decrease the energy of the s-orbitals in relation to the other atomic orbitals.
In practice the configurations predicted by the Madelung rule are at least close to the ground state even in these anomalous cases.
Electron configurations beyond this are tentative and predictions differ between models,[34] but Madelung's rule is expected to break down due to the closeness in energy of the 5g, 6f, 7d, and 8p1/2 orbitals.
Every system has the tendency to acquire the state of stability or a state of minimum energy, and so chemical elements take part in chemical reactions to acquire a stable electronic configuration similar to that of its nearest noble gas.
From Hund's rules, these electrons have parallel spins in the ground state, and so dioxygen has a net magnetic moment (it is paramagnetic).
The electronic configuration of polyatomic molecules can change without absorption or emission of a photon through vibronic couplings.
The most widespread application of electron configurations is in the rationalization of chemical properties, in both inorganic and organic chemistry.
This approach is taken further in computational chemistry, which typically attempts to make quantitative estimates of chemical properties.
In this case, it is necessary to supplement the electron configuration with one or more term symbols, which describe the different energy levels available to an atom.
