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Electronegativity

"Electronegative" redirects here. For the Nightfall EP, see Electronegative (EP).
A water molecule is put into a see-through egg shape, which is color-coded by electrostatic potential. A concentration of red is near the top of the shape, where the oxygen atom is, and gradually shifts through yellow, green, and then to blue near the lower-right and lower-left corners of the shape where the hydrogen atoms are.
Electrostatic potential map of a water molecule, where the oxygen atom has a more negative charge (red) than the positive (blue) hydrogen atoms

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.[1] 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).

The term "electronegativity" was introduced by Jöns Jacob Berzelius in 1811,[2] though the concept was known before that and was studied by many chemists including Avogadro.[2] In spite of its long history, an accurate scale of electronegativity was not developed until 1932, when Linus Pauling proposed an electronegativity scale which depends on bond energies, as a development of valence bond theory.[3] It has been shown to correlate with a number of other chemical properties. Electronegativity cannot be directly measured and must be calculated from other atomic or molecular properties. Several methods of calculation have been proposed, and although there may be small differences in the numerical values of the electronegativity, all methods show the same periodic trends between elements.[4]

The most commonly used method of calculation is that originally proposed by Linus Pauling. This gives a dimensionless quantity, commonly referred to as the Pauling scale (χr), on a relative scale running from 0.79 to 3.98 (hydrogen = 2.20). When other methods of calculation are used, it is conventional (although not obligatory) to quote the results on a scale that covers the same range of numerical values: this is known as an electronegativity in Pauling units.

As it is usually calculated, electronegativity is not a property of an atom alone, but rather a property of an atom in a molecule.[5] Even so, the electronegativity of an atom is strongly correlated with the first ionization energy. The electronegativity is slightly negatively correlated (for smaller electronegativity values) and rather strongly positively correlated (for most and larger electronegativity values) with the electron affinity.[6] It is to be expected that the electronegativity of an element will vary with its chemical environment,[7] but it is usually considered to be a transferable property, that is to say that similar values will be valid in a variety of situations.

Caesium is the least electronegative element (0.79); fluorine is the most (3.98).

  1. ^ IUPAC, Compendium of Chemical Terminology, 2nd ed. (the "Gold Book") (1997). Online corrected version: (2006–) "Electronegativity". doi:10.1351/goldbook.E01990
  2. ^ a b Jensen, W.B. (1996). "Electronegativity from Avogadro to Pauling: Part 1: Origins of the Electronegativity Concept". Journal of Chemical Education. 73 (1): 11–20. Bibcode:1996JChEd..73...11J. doi:10.1021/ed073p11.
  3. ^ Pauling, L. (1932). "The Nature of the Chemical Bond. IV. The Energy of Single Bonds and the Relative Electronegativity of Atoms". Journal of the American Chemical Society. 54 (9): 3570–3582. doi:10.1021/ja01348a011.
  4. ^ Sproul, Gordon D. (2020-05-26). "Evaluation of Electronegativity Scales". ACS Omega. 5 (20): 11585–11594. doi:10.1021/acsomega.0c00831. PMC 7254809. PMID 32478249.
  5. ^ Pauling, Linus (1960). Nature of the Chemical Bond. Cornell University Press. pp. 88–107. ISBN 978-0-8014-0333-0.
  6. ^ "PubChem ElectroNegativity," Downloaded, line-graphed, and correlated 'Electronegativity' with 'ElectronAffinity', showing a rather strong positive correlation of '0.712925965'. (Accessed linked site on 2023-09-16.)
  7. ^ Greenwood, N. N.; Earnshaw, A. (1984). Chemistry of the Elements. Pergamon. p. 30. ISBN 978-0-08-022057-4.

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