Electronegativity is one of the most important periodic properties in chemistry that explains the tendency of an atom to attract the shared pair of electrons towards itself in a chemical bond. The concept of electronegativity was introduced to understand the unequal sharing of electrons in covalent compounds and the nature of chemical bonding. It helps in predicting bond polarity, ionic and covalent character, metallic and non-metallic properties, and molecular behavior. The study of Pauling scale, Mulliken scale, periodic trends, and factors affecting electronegativity is extremely important for board examinations, JEE Main, JEE Advanced, and NEET.
What is Electronegativity ?
Electronegativity provides a qualitative measure of the ability of an atom in a chemical compound to attract the shared pair of electrons towards itself. We know that a covalent bond is formed by mutual sharing of electrons between two atoms. However, all covalent bonds are not similar. In order to understand the nature of a covalent bond between atoms, a new concept known as electronegativity was introduced by Pauling.
Electronegativity is defined as:
the tendency of an atom of an element to attract the shared pair of electrons towards itself in a covalent bond.
Greater the electronegativity of an atom, greater will be its tendency to attract the shared pair of electrons towards itself. Fluorine atom has the greatest power of attracting electrons and is the most electronegative element.
It must be remembered that unlike other atomic properties such as ionisation enthalpy, electron gain enthalpy which are related to individual gaseous atoms, the electronegativity is related to atoms in the bonded state.
Electronegativity values of different elements are not measured but are derived indirectly by different methods. Consequently, a number of relative scales of electronegativity such as Pauling scale, Mulliken–Jaffe scale, Allred–Rochow scale have been proposed. Among these, the most commonly used scale is that proposed by Pauling which is based upon the values of bond enthalpies of different bonds.
Pauling Scale of Electronegativity
In 1932, Linus Pauling was the first to propose a scale of electronegativity. This scale was based on the bond enthalpies in heteronuclear bonds. He suggested that if two atoms A and B had the same electronegativity values, then the bond enthalpy of A—B bond would be equal to the geometric mean of the bond enthalpies of A—A and B—B bonds. For example, the bond enthalpy of A—B molecule may be written as:
$$
E_{AB} = \sqrt{E_{AA} \times E_{BB}}
$$
This relation is based on the assumption of pure covalent bonds in A₂, B₂ and AB molecules. However, Pauling observed that for most of the $A—B$ bonds, the actual bond enthalpy is more than the geometric mean of bond enthalpies of $E_{AA}$ and $E_{BB}$. This means that the two participating atoms have different tendencies to attract the shared pair of electrons i.e., have different electronegativities.
The difference between the actual bond enthalpy of $A—B$ ($E_{AB}$) and the geometric mean bond enthalpy $[\sqrt{E_{AA} \times E_{BB}}]$ is known as excess bond enthalpy ($\Delta E$). It is a measure of difference in electronegativities of two atoms $A$ and $B$.
By calculating $\Delta E$ for different bonds, Pauling suggested the following relationship between the electronegativities of two atoms $\chi_A$ and $\chi_B$ and the excess bond enthalpy, $\Delta E$ :
$$
| \chi_A – \chi_B |= 0.12 \sqrt{\Delta E} \quad (\text{kJ mol}^{-1})
$$
It follows from the above relation that if $E(\text{A–B})$ is markedly different from the geometric mean of the covalent ($A–A$) and ($B–B$) bonds, then there is large electronegativity difference between the two atoms A and B.
Assigning an arbitrary value of one of the element, the electronegativity values of other elements can be easily calculated. Pauling assigned the value of 4.0 to the most electronegative element fluorine. The Pauling electronegativity values are given in Table below for normal elements. These are most commonly used values.
Table. Electronegativities of some elements (Pauling Scale)
| Group → | 1 | 2 | 13 | 14 | 15 | 16 | 17 |
|---|---|---|---|---|---|---|---|
| Period | |||||||
| 1 | H 2.1 | ||||||
| 2 | Li 1.0 | Be 1.5 | B 2.0 | C 2.5 | N 3.0 | O 3.5 | F 4.0 |
| 3 | Na 0.9 | Mg 1.2 | Al 1.5 | Si 1.8 | P 2.1 | S 2.5 | Cl 3.0 |
| 4 | K 0.8 | Ca 1.0 | Ga 1.6 | Ge 1.8 | As 2.0 | Se 2.4 | Br 2.8 |
| 5 | Rb 0.8 | Sr 1.0 | In 1.7 | Sn 1.7 | Sb 1.9 | Te 2.1 | I 2.4 |
| 6 | Cs 0.7 | Ba 0.9 | Tl 1.8 | Pb 1.7 | Bi 2.0 | Po 1.9 | At 2.2 |
Important Observations from the Table :
- Fluorine (F) has the highest electronegativity value (4.0).
- Cesium (Cs) has the lowest electronegativity value (0.7).
- Electronegativity increases from left to right across a period.
- Electronegativity decreases from top to bottom in a group.
The main disadvantage of Pauling scale is that the bond enthalpies are not known with good degree of accuracy for many elements. However, the values are relative values.
Mulliken Scale of Electronegativity
Robert Mulliken suggested another useful scale of electronegativity in terms of ionization enthalpy and electron gain enthalpy. Mulliken suggested that the tendency of an atom to attract the shared pair of electrons towards itself in a bond is the average of the ionization enthalpy ($\Delta_i H$) and electron gain enthalpy ($\Delta_{eg} H$). Therefore, Mulliken electronegativity is given as :
$$
\chi_M = \frac{\Delta_i H + | \Delta_{eg} H|} {2}
$$
where, $\Delta_i H$ = Ionization enthalpy
and $\Delta_{eg} H$ = Electron gain enthalpy.
The physical picture of Mulliken is reasonable because the tendency of an atom to attract the shared pair of electrons in a bond should be the average of the tendency of an atom to hold its own electrons ($\Delta_i H$) and its tendency to attract the additional electron ($\Delta_{eg} H$).
In addition, two more electronegativity scales — Allred and Rochow scale and Sanderson’s scale of electronegativity — were proposed.
“Learn more about What is Ionization Enthalpy? Definition, Units, Factors and Successive IE“
“Learn more about Ionization Enthalpy Trends Along a Period and Down a Group“
Factors Affecting Electronegativity
Electronegativity of an atom is not a fixed quantity but depends upon the following factors :
1. Oxidation state
In general, the electronegativity increases as the positive oxidation state of the atom increases. This is because with the increase in positive oxidation state, the tendency to attract the electrons will increase. For example, Pauling electronegativity value for Fe(II) is 1.83 whereas it is about 1.96 for Fe(III). Some common examples are :
Examples :
Tl(I): 1.62, Tl(III): 2.04, Sn(II): 1.80, Sn(IV): 1.96, Cu(I): 1.90, Cu(II): 2.00,
For anions, however, the electronegativity decreases with the increasing negative charge of the ion. This is due to the fact that a more negatively charged ion will attract less electrons than a less negatively charged (or neutral) ion.
2. Type of hybridisation
The type of hybridisation also affects the electronegativity of an atom. s-orbitals are nearer to the nucleus than p, d and f-orbitals because of their higher penetration power. Therefore, s-orbitals will have greater electron attracting power or electronegativity. In other words, the electronegativity increases with the increasing s-character of the hybrid orbitals for carbon.
Therefore, the electronegativity increases with the increasing s-character of the hybrid orbitals. For Carbon :
| Hydrocarbon | Hybridisation | s-character | Electronegativity |
|---|---|---|---|
| Methane, CH₄ | sp³ | 25% | 2.48 |
| Ethylene, C₂H₄ | sp² | 33% | 2.75 |
| Acetylene, C₂H₂ | sp | 50% | 3.29 |
The s-character increases in the order CH4, C2H4 and C2H2 and the electronegativity of carbon in these compounds also increases. Similarly, for nitrogen atom, the values of electronegativity are 3.68, 3.94 and 4.67 for sp³, sp² and sp hybridisation respectively.
3. Nature of the substituents
The electronegativity of a group varies with the nature of the substituents due to the inductive effect of the substituent group.
For example : CH₃ = 2.30, CCl₃ = 3.30, CF₃ = 3.35.
In these cases, the electronegativites of these groups will be the electronegativity of carbon as it is adjusted by the presence of substituents (3H, 3Cl or 3F atoms).
Periodic Variation of Electronegativity
(a) In a period
The electronegativity generally increases on moving across a period from left to right (e.g., from Li to F). This is primarily due to the decrease in atomic size and increase in effective nuclear charge. As a result of increase in effective nuclear charge, the attraction for the outer electrons and the nucleus increases in a period and therefore, electronegativity also increases.
(b) In a group
Electronegativity generally decreases from top to bottom in a group as atomic size increases and the bonding electrons move farther away from the nucleus. This trend is similar to that of ionization enthalpy.
It is clear from Table above that the electronegativity generally increases on moving across a period from left to right and decreases on moving down a group.
Thus, the alkali metals (Group 1) have the lowest electronegativities and the halogens (Group 17) have the highest electronegativities.
“Learn more about Atomic Radii, Ionic Radii and Isoelectronic Ions Trends Conceptual Questions and Answers of Periodic Table“
Electronegativity and Metallic and Non-metallic Character
Non-metallic elements have a strong tendency to gain electrons. Therefore, electronegativity is directly related to the non-metallic properties of elements. Alternatively, electronegativity is inversely related to the metallic properties of the elements.
The increase in electronegativity along a period is accompanied by an increase in non-metallic properties (or decrease in metallic properties).
The decrease in electronegativity down a group is accompanied by a decrease in non-metallic properties (or increase in metallic properties) of the elements.
Thus, fluorine (with the highest electronegativity of 4.0) is the most non-metallic element, while cesium (with the lowest electronegativity of 0.7) is the most metallic element.
The elements with high electronegativities (on the right-hand side of the periodic table) are non-metallic elements, while those having low electronegativities (on the left-hand side) are metallic elements. In general :
Elements with electronegativity 2.0 or greater are generally non-metals.
Elements with electronegativity less than 2.0 are generally metals.
The electronegativity helps in predicting the polar or non-polar bonds in molecules.
Differences between Electron Gain Enthalpy and Electronegativity
The main points of difference between electron gain enthalpy and electronegativity are as follows :
| Electron Gain Enthalpy | Electronegativity |
|---|---|
| It is the tendency of an isolated gaseous atom to attract an electron. | It is the tendency of an atom in a chemical compound (bonded state) to attract the shared pair of electrons. |
| It is the property of isolated atoms. | It is the property of atoms in the bonded state. |
| It is the absolute electron attracting power of an atom. | It is the relative electron attracting power of an atom. |
| It can be experimentally measured. | It cannot be measured experimentally; it is derived indirectly. |
| It has units such as kJ mol⁻¹ or eV/atom. | It is a dimensionless quantity (no units). There are only scales for comparison. |
| The electron gain enthalpy of an atom is constant. | The electronegativity of an atom is not constant. It depends upon the oxidation state of an atom, hybridisation state of the atom and the nature of substituents attached to it. |
“Learn more about Electron Gain Enthalpy : Definition, Units, Factors, Trends, Successive ΔegH“
FAQs Frequently Asked Short Questions and Answers Based on Electronegativity
What is electronegativity?
Electronegativity is the tendency of an atom in a chemical compound to attract the shared pair of electrons towards itself. It is mainly associated with covalent bonding where electrons are shared between atoms. The greater the electronegativity of an atom, the stronger will be its attraction for bonding electrons. Electronegativity helps in understanding the nature of chemical bonds and polarity of molecules.
Who introduced the concept of electronegativity?
The concept of electronegativity was introduced by Linus Pauling to explain the unequal sharing of electrons in covalent bonds. He proposed that different atoms have different tendencies to attract shared electrons. This concept became very important in chemical bonding and periodic properties. Pauling also developed the most widely used electronegativity scale.
Why is fluorine the most electronegative element?
Fluorine has the smallest atomic size and very high effective nuclear charge among the elements in its period. Because of this, it attracts the shared pair of electrons very strongly towards itself. Fluorine has the highest electronegativity value on the Pauling scale. Therefore, it is considered the most electronegative element.
Why is electronegativity considered a relative property?
Electronegativity is considered a relative property because it cannot be measured directly in absolute terms. Its values are obtained indirectly using different scales and comparisons between elements. The values mainly indicate the relative electron-attracting ability of atoms in compounds. Hence, electronegativity is expressed without units.
What is the Pauling scale of electronegativity?
The Pauling scale is the most commonly used scale for electronegativity. It was proposed by Linus Pauling and is based on bond enthalpy values of chemical bonds. In this scale, fluorine is assigned the highest value of 4.0. The scale helps compare the electron-attracting power of different elements.
What is the Mulliken scale of electronegativity?
The Mulliken scale defines electronegativity as the average of ionization enthalpy and electron gain enthalpy of an atom. It relates electronegativity to the tendency of an atom to lose or gain electrons. This scale provides a more theoretical explanation of electron attraction. It is useful in understanding the electronic behavior of atoms.
Why is electronegativity important in chemistry?
Electronegativity is important because it helps predict the nature of chemical bonds between atoms. It explains whether a bond will be polar or non-polar and also indicates ionic character in compounds. Electronegativity is widely used to study molecular polarity, reactivity, and bond formation. It is one of the key periodic properties of elements.
What are the factors affecting electronegativity?
Electronegativity is affected by oxidation state, hybridisation, atomic size, and the nature of substituents attached to the atom. An increase in positive oxidation state generally increases electronegativity. Greater s-character in hybrid orbitals also increases electronegativity. Inductive effects of substituents may further modify electron-attracting ability.
How does oxidation state affect electronegativity?
As the positive oxidation state of an atom increases, its electronegativity generally increases. This happens because the effective nuclear attraction on electrons becomes stronger. In contrast, negatively charged ions show lower electronegativity because they attract additional electrons less strongly. Thus, oxidation state significantly influences electronegativity values.
How does hybridisation affect electronegativity?
Electronegativity increases with increasing s-character of hybrid orbitals. Orbitals having greater s-character remain closer to the nucleus and attract electrons more strongly. Therefore, atoms in sp hybridisation are generally more electronegative than those in sp² or sp³ hybridisation. This effect is commonly observed in carbon compounds.
What is the periodic trend of electronegativity across a period?
Electronegativity generally increases from left to right across a period in the periodic table. This happens because atomic size decreases and effective nuclear charge increases. As a result, atoms attract bonding electrons more strongly. Non-metallic character also increases along a period.
What is the periodic trend of electronegativity down a group?
Electronegativity generally decreases from top to bottom in a group. The atomic size increases down the group, causing bonding electrons to move farther from the nucleus. Due to weaker attraction between the nucleus and shared electrons, electronegativity decreases. Metallic character increases down the group.
Which elements have the highest and lowest electronegativity?
Fluorine has the highest electronegativity value and is the most electronegative element. Cesium has one of the lowest electronegativity values and is highly metallic in nature. Highly electronegative elements are usually non-metals, whereas low electronegativity elements are generally metals. These trends are clearly observed in the periodic table.
How is electronegativity related to metallic and non-metallic character?
Electronegativity is directly related to non-metallic character and inversely related to metallic character. Elements with high electronegativity strongly attract electrons and behave as non-metals. Elements with low electronegativity easily lose electrons and show metallic behavior. Thus, non-metals are usually more electronegative than metals.
How does electronegativity help predict bond polarity?
When two bonded atoms have different electronegativities, the shared electrons are attracted more towards the more electronegative atom. This unequal sharing creates a polar covalent bond. Larger electronegativity difference leads to greater bond polarity. Therefore, electronegativity is very useful in predicting molecular polarity.
What is the difference between electron gain enthalpy and electronegativity?
Electron gain enthalpy refers to the tendency of an isolated gaseous atom to accept an electron, whereas electronegativity refers to the attraction of shared electrons in a bonded state. Electron gain enthalpy can be measured experimentally and has units. Electronegativity is a relative property without units and depends on bonding conditions. Both properties are related to electron attraction but differ in their physical meaning.
Why does electronegativity increase across a period?
Across a period, the effective nuclear charge increases while atomic size decreases. This causes the nucleus to attract shared electrons more strongly. As a result, the electronegativity of elements increases from left to right. This trend is clearly visible from alkali metals to halogens.
Why does electronegativity decrease down a group?
Down a group, the atomic size increases and additional electron shells are added. The outer electrons become farther from the nucleus and experience weaker attraction. Therefore, the ability to attract bonding electrons decreases. This causes electronegativity to decrease from top to bottom in a group.
Why are halogens highly electronegative?
Halogens have small atomic size and high effective nuclear charge. They require only one electron to complete their outermost shell and achieve stable electronic configuration. Therefore, they strongly attract shared electrons in chemical bonds. This makes halogens highly electronegative elements.
Can electronegativity predict ionic and covalent character?
Yes, electronegativity difference between atoms helps predict the nature of chemical bonds. A large electronegativity difference generally leads to ionic character, while a small difference results in covalent character. This concept is widely used in chemical bonding and molecular structure studies. It helps explain electron transfer and electron sharing in compounds.
Important Chapter Interlinks
This section provides a complete and interconnected study of Classification of Elements and Periodicity in Properties, starting with detailed theory and notes for Class 11 Chemistry to build a strong conceptual foundation. You can explore atomic radius and its types including covalent, van der Waals, metallic, and ionic radii to understand periodic trends in atomic size. It also includes Screening Effect (Shielding Effect) : Calculation of Effective or Reduced Nuclear Charge (Slater’s Rules), which explains how inner electrons reduce the nuclear attraction on outer electrons and influence periodic trends. In addition, topics like Radius of Cation is Less and Anion is More Than Its Parent Atom, Size Variation in Isoelectronic Series help explain how ionic size changes due to gain or loss of electrons and how nuclear charge affects size in species with the same number of electrons. The causes of periodicity explain why elements show repeating properties based on electronic configuration, which is further supported by the modern periodic law and structure of the modern periodic table including groups, periods, and blocks for elements even beyond atomic number 100. The historical development is covered through Mendeleev’s periodic law and table, leading to the modern classification of elements into s, p, d, and f blocks with prediction of period, group, and block. To strengthen exam preparation, you can practice JEE Main PYQs, IMU CET PYQs and Merchant Navy sponsorship exam MCQs, and other previous year questions with solutions, along with solved examples, conceptual questions, and practice problems on the modern periodic table. Learn more in this section also to radius of cation is less and anion is more than its parent atom and size variation in Isoelectronic Series. Additionally, complete study material, mock tests, and guidance are provided under Anand Classes Chemistry notes, along with expert support from Er Neeraj Anand, making this section a comprehensive resource for competitive exam preparation. This section also includes detailed study of What is Ionization Enthalpy? Definition, Units, Factors and Successive IE and Ionization Enthalpy Trends Along a Period and Down a Group for better understanding of periodic properties and reactivity of elements.