Electronegativity

Electronegativity is one of the fundamental concepts for the understanding of chemical bonding in chemistry. It provides guidelines for qualitatively understanding the difference between ionic bonding and covalent bonding, and is especially important for bonding between different types of atoms. The type of bond formed is largely determined by the difference in electronegativity between the atoms involved. It also allows us to estimate the polarity of a chemical bond and, when taken together with molecular geometry, the polarity of a molecule. Most chemical reactions have to do with polarity in some way, so electronegativity lies at the heart of chemistry. Given its importance, however, it is a difficult topic to pin down and there have been several approaches to deriving electronegativities of the elements. The first definition was suggested by Pauling and has not been improved upon since. In this definition electronegativity is defined as the ability of an atom in a molecule to attract electrons to itself. The opposite of electronegativity is termed electropositivity.

Basic Concepts

The neutral atoms of different elements have differing abilities to gain or lose electrons. These properties are known as the electon affinity and ionization energy of a given element, and they can be quantitated experimentally. Electron affinity of an element is a measure of the energy released (or gained, in some cases) when one electron is added to an atom of that element. Ionization energy is the energy needed to remove an electron from an atom of that element. Atoms that attract electrons more strongly have relatively higher ionization energy and electron affinity, and they tend to form monatomic ions with a negative charge. They tend to be the atoms of nonmetals. Atoms that attract electrons more weakly have lower ionization energy and electron affinity, and they form ions with a positive charge. They tend to be the atoms of metallic elements.

Given that electronegativity is based on the degree to which an atom attracts electrons, it can be seen as related to electron affinity and ionization energy. In a covalent bond between two atoms of two different elements, the electrons in the bond will be more stable when closer to the atom with greater attraction for electrons. Consequently, the electron cloud surrounding the two atoms becomes distorted, and the bond is said to be "polarized."

As might be expected, atoms with greater electron affinity and ionization energy have stronger attraction for the bonding electrons. In the case of electronegativity, however, the atoms are considered within the context of the chemical compound they are in, not as isolated atoms. Electronegativity, therefore, is not a property of the atom itself, though we tend to treat it as such. Rather, it depends on the state of the atom in the molecule. Consequently, the electronegativity of an element cannot be measured directly—it has to be calculated as an average, on a relative scale. Several methods have been proposed for calculating electronegativity.

Pauling scale

The most common and widely used scale for electronegativities is the Pauling scale, devised in 1932 by Linus Pauling. This is the scale commonly presented in general chemistry textbooks. Pauling based his scale on thermochemical data, particularly bond energies, which allowed him to calculate differences in electronegativity between atoms in a covalent bond. Hydrogen is arbitrarily assigned a value of 2.1, and all other electronegativities are calculated with respect to that. On this basis, the Pauling scale runs between 0 and 4, with 4 being the most electronegative. The most electronegative chemical element is fluorine, and its electronegativity on this scale is 3.98 (often rounded off to 4.0). The least electronegative element is francium, with a value of 0.7. The remaining elements have values in between.

Electronegativity trends

The trends in electronegativities of the elements are shown in the table below. In general, the degree of electronegativity decreases for the elements going down each group, and it increases across each period (from left to right). This pattern follows the general trends for the values of electron affinity and ionization energy. Moving across a period, nonmetals tend to have higher electron affinities and ionization energies; and moving down a group, the values for these properties tend to decrease. The most electronegative atoms are therefore clustered in the upper, right-hand corner of the periodic table (excluding the noble gases in group 18), and the least electronegative elements are located at the bottom left of the table.

→ Atomic radius decreases → Ionization energy increases → Electronegativity increases →
Group	1	2	3	4	5	6	7	8	9	10	11	12	13	14	15	16	17	18
Period
1	H 2.20																	He
2	Li 0.98	Be 1.57											B 2.04	C 2.55	N 3.04	O 3.44	F 3.98	Ne
3	Na 0.93	Mg 1.31											Al 1.61	Si 1.90	P 2.19	S 2.58	Cl 3.16	Ar
4	K 0.82	Ca 1.00	Sc 1.36	Ti 1.54	V 1.63	Cr 1.66	Mn 1.55	Fe 1.83	Co 1.88	Ni 1.91	Cu 1.90	Zn 1.65	Ga 1.81	Ge 2.01	As 2.18	Se 2.55	Br 2.96	Kr 3.00
5	Rb 0.82	Sr 0.95	Y 1.22	Zr 1.33	Nb 1.6	Mo 2.16	Tc 1.9	Ru 2.2	Rh 2.28	Pd 2.20	Ag 1.93	Cd 1.69	In 1.78	Sn 1.96	Sb 2.05	Te 2.1	I 2.66	Xe 2.6
6	Cs 0.79	Ba 0.89	*	Hf 1.3	Ta 1.5	W 2.36	Re 1.9	Os 2.2	Ir 2.20	Pt 2.28	Au 2.54	Hg 2.00	Tl 1.62	Pb 2.33	Bi 2.02	Po 2.0	At 2.2	Rn
7	Fr 0.7	Ra 0.9	**	Rf	Db	Sg	Bh	Hs	Mt	Ds	Rg	Uub	Uut	Uuq	Uup	Uuh	Uus	Uuo
Lanthanides	*	La 1.1	Ce 1.12	Pr 1.13	Nd 1.14	Pm 1.13	Sm 1.17	Eu 1.2	Gd 1.2	Tb 1.1	Dy 1.22	Ho 1.23	Er 1.24	Tm 1.25	Yb 1.1	Lu 1.27
Actinides	**	Ac 1.1	Th 1.3	Pa 1.5	U 1.38	Np 1.36	Pu 1.28	Am 1.13	Cm 1.28	Bk 1.3	Cf 1.3	Es 1.3	Fm 1.3	Md 1.3	No 1.3	Lr

Periodic table of electronegativity using the Pauling scale See also Periodic table

Qualitative Predictions

Electronegativities can be used to make qualitative predictions about the nature of a chemical bond between two atoms or elements. The usefull quantity here is actually the difference in electronegativity (ΔEN) between the two atoms. Bonds between atoms with a large electronegativity difference (greater than or equal to 1.7) are usually considered to be ionic, while values between 1.7 and 0.4 are considered polar covalent. Values below 0.4 are considered non-polar covalent bonds, and electronegativity differences of 0 indicate a completely non-polar covalent bond.

Electronegativity and Oxidation Number

Oxidation and reduction reactions take place through the transfer of electrons. Oxidation refers to a loss of electrons and reduction to a gain. In following the course of such a reaction it is important to be able to assign the number of electrons to an atom in the reactants and products of a chemical reaction. The most basic way chemists do this is to assign an oxidation number or oxidation state to each atom in a compound. The oxidation number signifies the number of charges an atom would have in a molecule if electrons were transfered completely. Essentially this means that the electrons in a chemical bond are considered to belong the most electronegative element. Thus the rules for assigning oxidation numbers are based on this concept of electronegativity.

Other Scales

There have been several methods proposed for determining electronegativities. Two general approaches are based either on electron affinity and ionization energy, or on size and charge. These methods give electronegativities that can be made similar in size and trends to the Pauling scale.

The Mulliken scale

This was proposed by Robert S. Mulliken in 1934 shortly after Pauling proposed his method. Mulliken suggested that the electronegativity of an atom should be an average of the electon affinity (EA_v) and ionization enrgy (IE_v) of the atom. Mulliken electronegativities, C_M, may be estimated by the following equation.

C_M = 0.168(IE_v + EA_v −1.23)

In this equation the electron affinity and ionization energy are reported in electron volts. These values must be calculated for the atom as it exists in the molecule, they are not the experimentally determined values for the neutral atom.

The Allred-Rochow scale

Allred and Rochow proposed this scale in 1958, and used the second general method, based on size and charge, to calculate electronegativities. They defined electronegativity as the electrostatic force exerted by the atomic nucleus on the valence electrons. When calculated according to the following equation they agree well with the Pauling scale.

C_R = 0.744 + 0.359Z_eff/r²

where Z_eff is the effective nuclear charge experienced by the valence electrons, and r is the distance between the electron and the nucleus (covalent radius).

Periodic tables

Standard table | Vertical table | Table with names | Names and atomic masses (large) | Names and atomic masses (small) | Names and atomic masses (text only) | Inline F-block | Elements to 218 | Electron configurations | Metals and non metals | Table by blocks | List of elements by name

Groups:   1 -  2 -  3 -  4 -  5 -  6 -  7 -  8 -  9 - 10 - 11 - 12 - 13 - 14 - 15 - 16 - 17 - 18

Periods:  1  -  2  -  3  -  4  -  5  -  6  -  7  -  8

Series:   Alkalis  -  Alkaline earths  -  Lanthanides  -  Actinides  -  Transition metals  -  Poor metals  -  Metalloids  -  Nonmetals  -  Halogens  -  Noble gases

Blocks:  s-block  -  p-block  -  d-block  -  f-block  -  g-block

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