Chemistry

Arrangement of Electrons the Basics of Electron Configuration



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Electrons are sub-atomic, negatively charged particles orbiting the nuclei of atoms. Unimaginably tiny, these particles only 1/2000th the mass of the nucleus move at nearly the speed of light in extremely tight orbits at varying distances from the nuclei based on the amount of energy the electrons have.

These distances—the areas of empty space—are the orbitals, and the arrangement of orbitals determines the unique chemical properties of the atoms, and thus the elements the atoms form. Electron configuration means the particular arrangement of electrons for each element.

The simplest atom is an isotope of hydrogen, protium. It has a single proton for a nucleus, and a single electron orbiting that proton. Since an atom must have at least one proton to be considered an atom, a handy way to keep track of electrons is to know the number of protons in the nucleus.

So while hydrogen can get by with but a single proton and its lonely electron, the upper limit of electrons may not have been reached, since humans are still creating new elements in labs. The highest number of electrons possible in a naturally occurring element is uranium, with 92 protons and 92 electrons in a stable atom.

Electrons in the orbitals are arranged in orbital shells, and these shells can be divided into sublevels (1s, 2s, 3s and up to 7s, for the s orbital). The original names—sharp, principal, diffuse and fundamental—are now used as shorthand: s, p, d and f.

The s orbital is spherical, the p orbital is shaped as if two water droplets were pulling apart, the d orbital has been described as butterfly shaped, and the f orbital often defies characterization, or even attempts to draw it, having seven different bulbous shapes possible.

The shapes of these orbitals are not physically limiting spaces. Rather, they are probability areas—places in which the chance of finding an electron is extremely high. The chance of finding a low-energy electron in the s orbital is extremely high, but exactly where it is at any given moment can never be determined. The chance of finding a more energetic electron within the shape of a p orbital is extremely high, but one cannot say conclusively where the electron is.

Electrons in the s orbital, for example, could actually be moving through the nucleus itself, despite the almost universal notion that electrons orbit the nucleus, staying well away from it due to the opposite charges of nucleus (+) and electron (-).

The Aufbau principle dictates that orbitals fill from the lowest (s-orbital) level first, and no orbital can hold more than two electrons with opposite spin (paired electrons) at a time.

We can infer the energy levels of the electrons from their orbital, and within the orbital, the shells, filling in precise order as 1s, 2s, 2p, 3s, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f, 6d and 7p.

Electrons move between levels by gaining or giving off energy. So an electron at the lowest energy level can be bumped to a higher level if energy is added, through a photon of light. That electron then jumps to a higher energy level. More often than not, however, the instability of the jump results in the electron giving off a burst of energy (a photon) and returning to its previous, lower energy level.                                                        

Most chemists do not concern themselves with energy levels of electrons and their configurations, having moved past the concept to spend more time on more challenging matters.

The orbitals provide a handy reference, though, giving a basis for work on bonding, formation of ions, prediction of cations and anions, and other chemistry.

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