Electronegativity

  • 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, symbol χ, is a concept that describes the tendency of an atom to attract a shared pair of electrons (or electron density) towards itself.[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 number, the more an atom or a substituent group attracts electrons towards itself.

    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/location of other electrons present 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 opposite of electronegativity is electropositivity: a measure of an element's ability to donate electrons.

    the term "electronegativity" was introduced by jöns jacob berzelius in 1811,[2] though the concept was known even 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.

    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.[4] properties of a free atom include ionization energy and electron affinity. it is to be expected that the electronegativity of an element will vary with its chemical environment,[5] 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 in the periodic table (= 0.79), while fluorine is most electronegative (= 3.98). francium and caesium were originally both assigned 0.7; caesium's value was later refined to 0.79, but no experimental data allows a similar refinement for francium. however, francium's ionization energy is known to be slightly higher than caesium's, in accordance with the relativistic stabilization of the 7s orbital, and this in turn implies that francium is in fact more electronegative than caesium.[6]

  • electronegativities of the elements
  • methods of calculation
  • correlation of electronegativity with other properties
  • trends in electronegativity
  • group electronegativity
  • electropositivity
  • see also
  • references
  • bibliography
  • external links

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, symbol χ, is a concept that describes the tendency of an atom to attract a shared pair of electrons (or electron density) towards itself.[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 number, the more an atom or a substituent group attracts electrons towards itself.

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/location of other electrons present 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 opposite of electronegativity is electropositivity: a measure of an element's ability to donate electrons.

The term "electronegativity" was introduced by Jöns Jacob Berzelius in 1811,[2] though the concept was known even 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.

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.[4] Properties of a free atom include ionization energy and electron affinity. It is to be expected that the electronegativity of an element will vary with its chemical environment,[5] 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 in the periodic table (= 0.79), while fluorine is most electronegative (= 3.98). Francium and caesium were originally both assigned 0.7; caesium's value was later refined to 0.79, but no experimental data allows a similar refinement for francium. However, francium's ionization energy is known to be slightly higher than caesium's, in accordance with the relativistic stabilization of the 7s orbital, and this in turn implies that francium is in fact more electronegative than caesium.[6]