Electronegativity and its measurement, shielding effect and diagonal relationship

Electronegativity

Pauling defined electronegativity as the tendency of an atom to attract the shared pair of electrons towards itself. When the covalent bond is formed electrons are shared equally between the two atoms and the atom which attract the electron towards itself gains negative charge and the other atoms gains positive charge. Atoms with smaller atomic size attractelectron more strongly than the large atom.

Electronegativity measurement on different scales

1. Pauling scale of electronegativity:

Pauling pointed out that the reactions of thetype given below are almost always exothermic, the bond formed between two atoms A and B must be stronger than the average of the single bond energies of A—A and B—B molecules.

A ₂+ B₂→ 2AB

Therefore, he based his scale on the difference between measured bond energy of a molecule A-B and the energy expected for the purely covalent bond A-B that is

ΔE= E_(A-B)- E_(A-B)cov ...........(i)

According to Pauling, the energy of the pure covalent bond (E_(A-B)) can be calculated as the geometric mean of the covalent energies of A-A and B-B molecules.

E_(A-B)cov =√(E_A-A.E_B-B)

or,E_(A-B)cov =(E_A-A.E_B-B)^{\(\frac{1}{2}\)}

If actual bond energy is equal to covalent bond energy the ΔE = 0, however, the measured bond energy is greater than the purely covalent bond energy thenΔE would be greater than zero. This excess bond energy is called is called Ionic-covalent resonance energy. Pauling evaluated the difference between the electronegativities A and B, which is given by,

\(\chi\)_A-\(\chi\)_B= 0.18 (ΔE)^{\(\frac{1}{2}\)}

Where\(\chi\)_A and\(\chi\)_Bare electronegativity of A and B respectively andΔE is expressed Kcalmol^-1

However ifΔE is expressed in ev mol^-1then the above expression becomes.

\(\chi\)_A-\(\chi\)_B = 0.208 (ΔE)^{\(\frac{1}{2}\)}

2. Mulliken scale of electronegativity:

In 1934, Mulliken gave an alternative method to calculate the electronegativity according to which electronegativity could be regarded as the arithmetic mean of the ionization energy and the electron affinity of an atom.

i.e, E..N. = \(\frac{I.E + E.A}{2}\) ev

The value in Mulliken scale is 2.8 times larger than Pauling scale. It can be converted into Pauling's scale with the unit in Kj mol^-1by

E.N = \(\frac{IE + EA}{2\times 28\times 96.48}\) = \(\frac{IE + EA}{540}\) Kj mol^-1

3. Allred and Rochow scale of electronegativity:

In 1958 Allred and Rochow considered electronegativity in a different way based on an electrostatic force between nucleus and electron. They defined electronegativity as the attractive force between a nucleus and an electron at a distance equal to the covalent radius.

F = Z_eff\(\times\) \(\frac{e^2}{r^2}\)

Where, e = change on the electron

Z_eff= effective nuclear charge

r = covalent radius

This F value may be converted to electronegativity value on the Pauling scale value using an empirical relationship.

\(\chi\) = 0.744 + 0.395\(\times\) \(\frac{Z_eff}{r^2}\)

Shielding effect

When an electron enters into extra atomic orbital, the number of electrons obviously increases and repulsive force act among these electrons but at the same time nuclear charge also increases. However, the nuclear charge is overcome by repulsive force so the outer electrons experience less attraction towards thenucleus. This is called shielding effect. Therefore , the decrease in anattractive force of nucleus for the outer electrons due to the inner shell electrons lying between the outer electrons (valence electrons) and a nucleus is called shielding effect.

The concept of shielding effect has been used to explain why the ionization potential values of the element of given group decreaseon descending down the group. The shielding effect has also been used to explain that when we proceed from alkali metal to inert gas in thesame period large increase in ionization potential is observed.

Factors influencing the shielding effect

1. Shielding effect is affected by the number inner shell electrons. If a large number of electrons are present between valence electrons and nucleus then the screening effect increases and the nuclear attraction for the valence electrons decreases. Therefore shielding effect increases from top to bottom of the periodic table.
2. Shielding effect is also influenced by thetype of orbital that the electrons exist in. The shielding effect decreases in following order in different orbitals, S> P> d> f. S-orbital electrons are closer to thenucleus than the electrons in p, d and f orbitals. Due to this, the penetration power of s-electron is maximum in comparison to p, d and f orbitals and it shields the nucleus more effectively than the p-electrons. Similarly, p-electrons screen the nuclear charge more efficiently than the d-electrons and so on. Therefore, the electrons in the larger atomic orbitals d and f shield the nuclear charge poorly and there exist contraction isknown as d-block and f-block contraction.

Diagonal relationship

From left to the right of the period, the number of electrons in the outer shell increases from one to eight. Therefore, all elements in Group 1 have one electron in their valence shell and are univalent (lose only one electron when they react). Likewise, the elements in the Group 2 have two electrons in their valence shell and are divalent.

Here, on moving from top to bottom of the group, the elements all have the same number of outer electrons and valency, but the size increases. Therefore ionization energy decreases and the metallic character increases.

On moving diagonally across the periodic table, the elements shows certain similarities. However these are not as strong as the similarities within a group but are significant in the following pairs of elements:

Li Be B C

Na Mg Al Si

Across the period, the size decreases and the charge on the ion increases which cause the increase of the polarizing power. On the other hand, on moving down a group, the size increases, and the polarizing power decreases. However, diagonally these effects are cancelled by each other, so there is no distinctive change in properties. For example, Be is smaller than Mg and Al is smaller than Mg. So, Li and Mg are nearly equal in size and their behaviour are similar.

Diagonal relationship between Li and Mg

Both Li and Mg have ahigh melting point comparing to other metals in their respective group and are harder. The reaction of Li and Mg is slower but other alkali metals react vigorously with water. The nitrides of Li (Li₃N) and Mg (Mg₃N₂) gives NH₃reacting with H₂O. Li and Mg do not form peroxide and superoxides like other metals but forms Li₂O and MgO respectively.The hydroxides of Li and Mg (LiOH) and Mg(OH)₂are sparingly soluble in water but other alkali metal hydroxides are highly soluble in water. The ions of Li and Mg both forms hydrate.

Diagonal relationship between Be and Al

1. Be and Al dissolves rapidly in H₂SO₄but slowly with HNO₃.
2. Both reacts with strong hydroxides giving H₂, berrylate and aluminate ion.
3. They both reacts with an acid and base i.e, amphoteric.
4. Their carbides react with H₂O to give CH₄.
5. BeX_2,AlX₃are covalent.

Diagonal relationship between B and Si

1. The crystals of B and Si are asemiconductor and have high M.P.
2. The oxides of B and Si are acidic.
3. The hydroxides of B and Si are weakly acidic.
4. Boranes and silanes are prepared similarly.

References

Lee, J.D. Concise Inorganic Chemistry. John Wiley and sons.Inc, 2007.

chemistry stake exchange. n.d. <http://chemistry.stackexchange.com/questions/32591/diagonal-relationships-in-the-periodic-table>.

chemwiki. n.d. <http://chemwiki.ucdavis.edu/Core/Physical_Chemistry/Physical_Properties_of_Matter/Atomic_and_Molecular_Properties/Electronegativity/Mulliken_Electronegativity>.

wikipedia. n.d. <https://en.wikipedia.org/wiki/Shielding_effect>.