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

Neutral atoms of different elements have differing abilities to gain or loose electrons. The magnitude of these properties can be experimentally determined for each element and are the electon affinity and ionization energy of the element. The electron affinity is a measure of the energy released (or gained in some cases) when one electron is added to the atom, and the ionization energy is the energy needed to remove an electron from the atom. Atoms with a greater "pull" on their electrons have a high ionization energy and high electron affinity and tend to form monatomic ions with a negative charge. These tend to be the atoms of the non metal elements. Atoms with a weaker "pull" have a low ionization energy and low electron affinity, form ions with a positive charge, and are the metallic elements.

Since electronegativity is due to the "pull" on electrons it can be seen to be related to electron affinity and ionization energy. In a covalent bond between different atoms the electrons in the bond will be more stable in the presence of the atom that has greater attraction for electrons. This leads to a distortion of the electon cloud and the bonding electrons spend more time close to that atom. The bond is said to be polarized. As might be expected atoms with the greater electron affinity and ionization energy have the greater attraction for the bonding electrons. However in electronegativity we are looking at the atoms in the context of the chemical compound it is in, not the neutral atoms themselves. One consequence is that electronegativity 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 and we cannot directly measure the electronegativity of an element. It has to be calculated as an average on a relative scale. There have been several methods for calculating electronegativity.

Pauling scale

The most common and widely used scale of 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 calcluate differences in electronegativity between atoms in a covalent bond. Hydrogen is arbitrarily assinged a value of 2.1 and all other electronegativities calculated with respect to it. This leads to a scale that runs between 0 and 4, with 4 being the most electronegative. On this scale, the most electronegative chemical element (fluorine) is given an electronegativity value of 3.98 (textbooks often state this value to be 4.0); the least electronegative element (francium) has a value of 0.7, and 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 down each group and increases across the periods. This pattern follows the general trends in electron affinity and ionization energy. Moving across a period, non-metals tend have higher electron affinities and ionization energies, and moving down a group these properties tend to decrease. Therefore, the most electronegative atoms can be found in the upper, right hand side of the periodic table, and the least electronegative elements can be found at the bottom left.

→ 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

The Mulliken scale was proposed by Robert S. Mulliken in 1934. On the Mulliken scale, numbers are obtained by averaging ionization potential and electron affinity. Consequently, the Mulliken electronegativities are expressed directly in energy units, usually electron volts.

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