Shielding effect in the context of "Electropositive"

Electronegativity, symbolized as χ, is the tendency for an atom of a given chemical element to attract shared electrons (or electron density) when forming a chemical bond. An atom's electronegativity is affected by both its atomic number and the distance at which its valence electrons reside from the charged nucleus. The higher the associated electronegativity, the more an atom or a substituent group attracts electrons. Electronegativity serves as a simple way to quantitatively estimate the bond energy, and the sign and magnitude of a bond's chemical polarity, which characterizes a bond along the continuous scale from covalent to ionic bonding. The loosely defined term electropositivity is the opposite of electronegativity: it characterizes an element's tendency to donate valence electrons.

On the most basic level, electronegativity is determined by factors like the nuclear charge (the more protons an atom has, the more "pull" it will have on electrons) and the number and location of other electrons in the atomic shells (the more electrons an atom has, the farther from the nucleus the valence electrons will be, and as a result, the less positive charge they will experience—both because of their increased distance from the nucleus and because the other electrons in the lower energy core orbitals will act to shield the valence electrons from the positively charged nucleus).

A dianion is an anion with a net charge of −2. While there exist many stable molecular dianions, such as BeF₄ and MgF₄, thus far no stable atomic dianion has been found: Electron shielding and other quantum mechanical effects tend to make the addition of another electron to an atomic anion unstable.

The most heavily studied atomic dianion is H, usually as a short-lived resonance between an electron and a hydrogen ion. In 1976, its half-life was experimentally measured to be 23 ± 4 nanoseconds.

In atomic physics, the effective nuclear charge of an electron in a multi-electron atom or ion is the number of elementary charges (${\displaystyle e}$) an electron experiences by the nucleus. It is denoted by Z_eff. The term "effective" is used because the shielding effect of negatively charged electrons prevent higher energy electrons from experiencing the full nuclear charge of the nucleus due to the repelling effect of inner layer. The effective nuclear charge experienced by an electron is also called the core charge. It is possible to determine the strength of the nuclear charge by the oxidation number of the atom. Most of the physical and chemical properties of the elements can be explained on the basis of electronic configuration. Consider the behavior of ionization energies in the periodic table. It is known that the magnitude of ionization potential depends upon the following factors:

1. The size of an atom
2. The nuclear charge; oxidation number
3. The screening effect of the inner shells
4. The extent to which the outermost electron penetrates into the charge cloud set up by the inner lying electron

In the periodic table, effective nuclear charge decreases down a group and increases left to right across a period.