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Electronegativity Trends

The long form of periodic table carries elements according to the increasing order of their atomic numbers. This in turn, is actually the increasing number of outermost electrons. When we look at the arrangements though, there are certain properties of these elements which are also periodic in nature.

These are called periodic properties and are discussed and understood before going over to chemical bonding or the nature of compounds formed by any group of elements.

 

Define Electronegativity

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Consider the formation of a covalent bond between two similar atoms of a molecule like Hydrogen. In this molecule the electron pair participating in the formation of covalent bond is shared equally by both hydrogen atoms, that is, the electron pair lies exactly in the center of the molecule.

On the other hand, consider the formation of a covalent bond between two dissimilar atoms of a molecule like HCl. In this molecule the electron pair participating in the formation of covalent bond is not shared equally by the two atoms, the hydrogen and chlorine. The electron pair tends to lie nearer the Cl atom than the Hydrogen atom.

The reason for this unequal sharing of electron pair is given by the fact that Cl atom has a greater tendency than the hydrogen atom to attract the shared electron pair between them towards itself. Thus, we can also say that the Cl atom has more electronegativity than the Hydrogen atom.

From the above fact, electronegativity can be defined as

" The electronegativity of a bonded atom is defined as its relative tendency (or ability) to attract the shared electron pair towards itself."

Electronegativity of an atom, A is generally expressed as $\chi _{A}$.

The concept of electronegativity was proposed by the American chemist Linus Pauling in 1932. Electronegativity is a relative quantity. It is basically written in comparing two atoms involved in the covalent bond formation. It does not have any unit.

Some of the scales suggested for measuring the electronegativity are

1. Pauling's bond energy scale


Where bond energies or the energy required to break a bond, to get neutral atoms, is employed to measure electronegativity of an element.

2. Mullikan's scale


According to this scale, electronegativity is taken as a mean difference between ionization potential of an element and its electron affinity.

Factors Affecting Electronegativity

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The electronegativity of an element is influenced by following factors.

1. Nuclear charge


The higher the nuclear charge, more will be the electronegativity value of an element, since the nucleus will be able to attract or pull more electrons towards itself.

2. Atomic size


Size of the atom is inversely proportional to electronegativity value. The smaller the element, higher is its nucleus's reach towards the outer shell. Thus, more will be the electronegativity. An example of this is Fluorine, the first member of the halogen series.

Fluorine is the most highest electronegative element available and is the smallest one too.

3. Screening effect or shielding effect


If the outermost electrons shields the nucleus effectively, the electronegativity of the element decreases.

Periodic Table Electronegativity Trend

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In a period


Electronegativity increases on moving in a period of the periodic table from left to right. This is due to the increase in nuclear charge as a result of which the added electrons can be held more tightly. Thus, C-N bond should be shown as Cδ+ - N δ- or C→N, the arrow head being towards the more electronegative element, N.

In the same period, on moving from left to right, the electronegativities increase with the increase in the number of outer electrons. An example of this can be shown with the help of the second period.

Element
Li Be
B
C
N
O
F
Electronegativity value
1.0 1.5 2.0 2.5 3.0 3.5 4.0

In a group


In moving down a group, from top to bottom, a new shell is added from the top element to the bottom one. Also, the nuclear charge increases from top to bottom.
The increase in the nuclear charge indicates that the electronegativity of the lower element should be more than that of the top element. This is not the case though.

The reason is because of the increase in the atomic radii as we move down the group. The increase in atomic radii increases the electron shielding effect, and it is much more than the increase in the nuclear charge. Consequently, the lower elements are less electronegative than the top elements. Thus, as we move down a group. electronegativity decreases.

The most electronegative elements, say, Fluorine for example, are present at the top right hand corner of the periodic table. while the most electropositive elements or the less electronegative elements are present at the bottom left hand corner, for example, Cs.

Trend for Electronegativity

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Electronegativity value of elements follows the below mentioned trend:
  1. Electronegativity increases as we move down from left to right across a period. This is because the atomic size decreases across the period and nuclear attraction over electrons increases.
  2. On moving down a group, the atomic size increases and the nuclear attraction over outer electrons decreases. Consequently, electronegativity decreases as we move from top to bottom in a group.
  3. On comparing the elements based on their group, Halogens are the most electronegative elements in their respective periods.
  4. Alkali metals, have the least electronegativity value in their periods. The last group in the periodic table, the 18th group consists of Noble gases. They have zero electronegativity value because of their completely filled outer shells.
  5. Fluorine, the smallest element has the highest electronegativity value among all elements and Helium has the lowest electronegativity value.
  6. The 15, 16, and 17th group of elements are comparatively more electronegative due to the almost filled outermost shells.
  7. The first and second group, alkali metals and alkaline earth metals are least electronegative. These are electropositive in nature.
  8. Electronegativity value is only applicable for bonded atoms, and not for the isolated gaseous atoms.
  9. It is usually calculated from experimentally determined values of bond energies compared to a reference.
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