What is Electronegativity?
Electronegativity is a fundamental chemical property that describes the tendency of an atom to attract shared electrons toward itself when it forms a chemical bond with another atom. The concept was first introduced by Linus Pauling in 1932, and it remains one of the most important concepts in understanding chemical bonding, molecular polarity, and reactivity.
When two atoms bond together by sharing electrons (covalent bond), the electrons are not always shared equally. The atom with higher electronegativity pulls the shared electron density closer to itself, creating a partial negative charge on that atom and a partial positive charge on the other atom. This uneven distribution of electron density is the basis of bond polarity.
Electronegativity is a dimensionless quantity, not an absolute measurable energy. Rather, it is a relative scale that allows chemists to compare how strongly different atoms attract bonding electrons. The most commonly used scale is the Pauling scale, on which fluorine -- the most electronegative element -- is assigned a value of 3.98, and cesium -- one of the least electronegative -- has a value of 0.79.
Electronegativity Scales
Several different electronegativity scales have been proposed over the decades. While they each use different methods of calculation, they generally agree on the relative ordering of elements. The three most widely used scales are:
1. Pauling Scale
The Pauling scale is the oldest and most widely used electronegativity scale. Linus Pauling developed it in 1932 by analyzing bond dissociation energies. He observed that the bond energy of a heteronuclear diatomic molecule A-B is usually greater than the average of the bond energies of A-A and B-B. He attributed this extra stabilization energy to the ionic resonance energy arising from the electronegativity difference.
Pauling defined the electronegativity difference between atoms A and B using the formula:
where ΔEd = D(A-B) - ½[D(A-A) + D(B-B)], and D represents bond dissociation energies in kJ/mol. Hydrogen was originally set as the reference with χ = 2.20, and all other values were determined relative to it.
The Pauling scale ranges from about 0.79 (cesium) to 3.98 (fluorine). Noble gases are generally not assigned Pauling electronegativity values because they rarely form bonds under normal conditions, although krypton and xenon have been assigned values based on their known compounds.
2. Mulliken Scale
Robert S. Mulliken proposed an alternative scale in 1934 based on the arithmetic mean of the first ionization energy (IE) and the electron affinity (EA) of an atom:
This approach has a clear physical meaning: an atom that has both a high ionization energy (hard to remove electrons) and a high electron affinity (strongly attracts electrons) will be highly electronegative. Mulliken values are typically expressed in electron volts (eV) and can be converted to the Pauling scale using the relationship: χPauling ≈ 0.359 × χMulliken1/2 + 0.744. The advantage of the Mulliken scale is that it is based on measurable atomic properties and can be calculated for different valence states of an element.
3. Allred-Rochow Scale
In 1958, A. Louis Allred and Eugene G. Rochow proposed a scale based on electrostatic force. They defined electronegativity as the electrostatic force exerted by the nucleus on the valence electrons, taking into account the shielding effect of inner electron shells:
where Zeff is the effective nuclear charge (calculated using Slater's rules) and rcov is the covalent radius in angstroms.
The Allred-Rochow scale tends to give values that are numerically similar to the Pauling scale but can be calculated more easily for elements with less experimental bonding data. It is particularly useful for comparing electronegativity values across the periodic table systematically.
Electronegativity Trends in the Periodic Table
Electronegativity follows two clear trends in the periodic table:
- Across a period (left to right): Electronegativity generally increases. As you move across a period, the nuclear charge increases while the shielding effect remains relatively constant (electrons are added to the same shell). The increased nuclear charge pulls bonding electrons more strongly, resulting in higher electronegativity. For example, in Period 2: Li (0.98) → Be (1.57) → B (2.04) → C (2.55) → N (3.04) → O (3.44) → F (3.98).
- Down a group (top to bottom): Electronegativity generally decreases. As you move down a group, each successive element has an additional electron shell, increasing the atomic radius and the shielding of the nucleus from the bonding electrons. The valence electrons are farther from the nucleus and less strongly attracted. For example, in Group 17 (halogens): F (3.98) → Cl (3.16) → Br (2.96) → I (2.66).
The element with the highest electronegativity is fluorine (F) at 3.98, located in the top-right corner of the periodic table (excluding noble gases). The elements with the lowest electronegativity are the alkali metals in the bottom-left, with cesium (Cs) at 0.79 and francium (Fr) at approximately 0.70.
Electronegativity Trends - Periodic Table Overview
Electronegativity Chart - Pauling Scale Values
The following comprehensive table lists the Pauling electronegativity values for the most commonly encountered elements in chemistry. These values are used by this calculator to determine bond types.
| Z | Symbol | Element | EN (Pauling) | Group | Period |
|---|---|---|---|---|---|
| 1 | H | Hydrogen | 2.20 | 1 | 1 |
| 3 | Li | Lithium | 0.98 | 1 | 2 |
| 4 | Be | Beryllium | 1.57 | 2 | 2 |
| 5 | B | Boron | 2.04 | 13 | 2 |
| 6 | C | Carbon | 2.55 | 14 | 2 |
| 7 | N | Nitrogen | 3.04 | 15 | 2 |
| 8 | O | Oxygen | 3.44 | 16 | 2 |
| 9 | F | Fluorine | 3.98 | 17 | 2 |
| 11 | Na | Sodium | 0.93 | 1 | 3 |
| 12 | Mg | Magnesium | 1.31 | 2 | 3 |
| 13 | Al | Aluminum | 1.61 | 13 | 3 |
| 14 | Si | Silicon | 1.90 | 14 | 3 |
| 15 | P | Phosphorus | 2.19 | 15 | 3 |
| 16 | S | Sulfur | 2.58 | 16 | 3 |
| 17 | Cl | Chlorine | 3.16 | 17 | 3 |
| 19 | K | Potassium | 0.82 | 1 | 4 |
| 20 | Ca | Calcium | 1.00 | 2 | 4 |
| 21 | Sc | Scandium | 1.36 | 3 | 4 |
| 22 | Ti | Titanium | 1.54 | 4 | 4 |
| 23 | V | Vanadium | 1.63 | 5 | 4 |
| 24 | Cr | Chromium | 1.66 | 6 | 4 |
| 25 | Mn | Manganese | 1.55 | 7 | 4 |
| 26 | Fe | Iron | 1.83 | 8 | 4 |
| 27 | Co | Cobalt | 1.88 | 9 | 4 |
| 28 | Ni | Nickel | 1.91 | 10 | 4 |
| 29 | Cu | Copper | 1.90 | 11 | 4 |
| 30 | Zn | Zinc | 1.65 | 12 | 4 |
| 31 | Ga | Gallium | 1.81 | 13 | 4 |
| 32 | Ge | Germanium | 2.01 | 14 | 4 |
| 33 | As | Arsenic | 2.18 | 15 | 4 |
| 34 | Se | Selenium | 2.55 | 16 | 4 |
| 35 | Br | Bromine | 2.96 | 17 | 4 |
| 36 | Kr | Krypton | 3.00 | 18 | 4 |
| 37 | Rb | Rubidium | 0.82 | 1 | 5 |
| 38 | Sr | Strontium | 0.95 | 2 | 5 |
| 39 | Y | Yttrium | 1.22 | 3 | 5 |
| 40 | Zr | Zirconium | 1.33 | 4 | 5 |
| 41 | Nb | Niobium | 1.60 | 5 | 5 |
| 42 | Mo | Molybdenum | 2.16 | 6 | 5 |
| 44 | Ru | Ruthenium | 2.20 | 8 | 5 |
| 45 | Rh | Rhodium | 2.28 | 9 | 5 |
| 46 | Pd | Palladium | 2.20 | 10 | 5 |
| 47 | Ag | Silver | 1.93 | 11 | 5 |
| 48 | Cd | Cadmium | 1.69 | 12 | 5 |
| 49 | In | Indium | 1.78 | 13 | 5 |
| 50 | Sn | Tin | 1.96 | 14 | 5 |
| 51 | Sb | Antimony | 2.05 | 15 | 5 |
| 52 | Te | Tellurium | 2.10 | 16 | 5 |
| 53 | I | Iodine | 2.66 | 17 | 5 |
| 54 | Xe | Xenon | 2.60 | 18 | 5 |
| 55 | Cs | Cesium | 0.79 | 1 | 6 |
| 56 | Ba | Barium | 0.89 | 2 | 6 |
| 57 | La | Lanthanum | 1.10 | 3 | 6 |
| 72 | Hf | Hafnium | 1.30 | 4 | 6 |
| 73 | Ta | Tantalum | 1.50 | 5 | 6 |
| 74 | W | Tungsten | 2.36 | 6 | 6 |
| 75 | Re | Rhenium | 1.90 | 7 | 6 |
| 76 | Os | Osmium | 2.20 | 8 | 6 |
| 77 | Ir | Iridium | 2.20 | 9 | 6 |
| 78 | Pt | Platinum | 2.28 | 10 | 6 |
| 79 | Au | Gold | 2.54 | 11 | 6 |
| 80 | Hg | Mercury | 2.00 | 12 | 6 |
| 81 | Tl | Thallium | 1.62 | 13 | 6 |
| 82 | Pb | Lead | 2.33 | 14 | 6 |
| 83 | Bi | Bismuth | 2.02 | 15 | 6 |
| 84 | Po | Polonium | 2.00 | 16 | 6 |
| 85 | At | Astatine | 2.20 | 17 | 6 |
| 87 | Fr | Francium | 0.70 | 1 | 7 |
| 88 | Ra | Radium | 0.90 | 2 | 7 |
How to Determine Bond Type from Electronegativity
The electronegativity difference between two bonded atoms is the primary method for predicting the type of chemical bond that will form between them. This calculator uses the widely accepted threshold values on the Pauling scale:
| EN Difference | Bond Type | Electron Sharing | Example |
|---|---|---|---|
| < 0.4 | Nonpolar Covalent | Electrons shared equally or nearly equally | H-H (0.0), C-H (0.35) |
| 0.4 to 1.7 | Polar Covalent | Electrons shared unequally; partial charges develop | H-Cl (0.96), O-H (1.24), C-O (0.89) |
| ≥ 1.7 | Ionic | Electrons effectively transferred from one atom to the other | Na-Cl (2.23), K-F (3.16), Ca-O (2.44) |
Percentage Ionic Character
Rather than treating bonds as strictly one type or another, chemists often quantify the degree of ionic character using Pauling's empirical formula:
where Δχ is the electronegativity difference between the two atoms. This formula yields a smooth curve: a difference of 0 gives 0% ionic character, a difference of 1.0 gives about 22%, and a difference of 1.7 gives about 51%. A fully ionic bond would require a very large (theoretically infinite) electronegativity difference.
This percentage is particularly useful because it acknowledges that real chemical bonds exist on a spectrum. A bond with 30% ionic character still has significant covalent character. The formula is an approximation, but it provides useful insight into the nature of the bond.
Bond Polarity and Dipole Moments
When a bond is polar (i.e., the two atoms have different electronegativities), the uneven distribution of electron density creates a dipole moment. The dipole moment (μ) is a vector quantity that points from the partial positive charge toward the partial negative charge, and its magnitude depends on both the charge separation and the bond length:
where δ is the magnitude of the partial charge (in coulombs) and d is the distance between the charges (in meters). Dipole moments are typically measured in debyes (D), where 1 D = 3.336 × 10-30 C·m.
In polyatomic molecules, the overall molecular dipole moment is the vector sum of all individual bond dipole moments. This means that a molecule can have polar bonds but still be nonpolar overall if the bond dipoles cancel each other out due to molecular symmetry. Carbon dioxide (CO2) is a classic example: each C=O bond is polar, but because the molecule is linear and symmetric, the two bond dipoles cancel, resulting in zero net dipole moment. Water (H2O), by contrast, has a bent geometry, so the two O-H bond dipoles do not cancel, giving water a significant dipole moment of 1.85 D.
The magnitude of the dipole moment generally increases with increasing electronegativity difference, but bond length also plays a significant role. For example, HF has a larger electronegativity difference (1.78) than HCl (0.96), and HF also has a shorter bond length, yet HF's dipole moment (1.82 D) is actually close to that of HCl (1.08 D) because of the much shorter bond distance.
Electropositivity
Electropositivity (also called metallicity) is the inverse concept of electronegativity. It describes the tendency of an atom to donate or lose electrons when forming a chemical bond. Electropositive elements readily lose electrons to form positive ions (cations).
The most electropositive elements are the alkali metals (Group 1) and alkaline earth metals (Group 2), which are found on the far left side of the periodic table. Cesium (Cs) and francium (Fr) are the most electropositive elements. These elements have very low ionization energies and very low electronegativities, meaning they readily give up their valence electrons.
Electropositivity trends are exactly opposite to electronegativity trends:
- Across a period (left to right): Electropositivity decreases.
- Down a group (top to bottom): Electropositivity increases.
In ionic bonding, the electropositive element donates its electron(s) to the electronegative element. The greater the difference in electropositivity/electronegativity between two elements, the more ionic the bond tends to be.
Notable Elements
Fluorine - The Most Electronegative Element (3.98)
Fluorine holds the distinction of being the most electronegative element on the Pauling scale with a value of 3.98. This extreme electron-attracting ability is due to its small atomic radius and high effective nuclear charge. Fluorine's high electronegativity makes it form extremely polar bonds with almost every other element. The H-F bond, for instance, has an electronegativity difference of 1.78, giving it about 55% ionic character despite being classified as a covalent bond.
Fluorine's electronegativity also explains why it is the strongest oxidizing agent among the elements and why fluorine compounds tend to have unusual properties. The C-F bond is one of the strongest single bonds in organic chemistry (bond dissociation energy of approximately 485 kJ/mol), which is why fluorinated compounds (such as PTFE/Teflon) are exceptionally stable and chemically resistant.
Cesium - Among the Least Electronegative Elements (0.79)
Cesium (Cs) has one of the lowest Pauling electronegativity values at 0.79 (francium at approximately 0.70 is slightly lower but is extremely rare and radioactive). Cesium's low electronegativity results from its large atomic radius (260 pm) and strong shielding of the nucleus by five inner electron shells. The single 6s valence electron is weakly held and easily donated in bonding.
When cesium bonds with fluorine, the electronegativity difference is 3.98 - 0.79 = 3.19, one of the largest possible differences between any two elements. CsF has approximately 93% ionic character, making it one of the most ionic compounds known. Cesium is so electropositive that it reacts explosively with water and even with ice at temperatures as low as -116 degrees Celsius.
Oxygen - The Second Most Electronegative (3.44)
Oxygen is the second most electronegative element with a Pauling value of 3.44. Its high electronegativity is responsible for many of water's unique properties, including its high boiling point, surface tension, and ability to act as a universal solvent for ionic and polar compounds. The electronegativity difference between O and H (1.24) creates the polar O-H bonds that enable hydrogen bonding, which is critical to the structure of water, proteins, and DNA.
Electron Affinity vs Electronegativity
While both electron affinity and electronegativity relate to an atom's interaction with electrons, they are fundamentally different properties:
| Property | Electronegativity | Electron Affinity |
|---|---|---|
| Definition | Tendency to attract shared electrons in a chemical bond | Energy change when a neutral atom gains an electron in the gas phase |
| Context | Property of an atom within a molecule (bonded state) | Property of an isolated atom (gas phase) |
| Units | Dimensionless (Pauling scale) or eV (Mulliken scale) | kJ/mol or eV |
| Highest for | Fluorine (3.98) | Chlorine (349 kJ/mol) |
| Noble gases | Generally not assigned values | Near zero or negative (endothermic) |
| Measurement | Derived from bond energies or atomic properties | Directly measured experimentally |
A notable discrepancy exists between the two properties for fluorine and chlorine. Fluorine has the highest electronegativity (3.98) but not the highest electron affinity. Chlorine actually has a higher electron affinity (349 kJ/mol) than fluorine (328 kJ/mol). This is because fluorine's extremely small size causes strong electron-electron repulsion when an additional electron is added to the already compact 2p orbital. In a bond, however, the shared electron density is distributed between two atoms, reducing this repulsion issue, so fluorine's ability to attract bonding electrons is still the strongest of all elements.
Frequently Asked Questions
Electronegativity is a measure of how strongly an atom attracts shared electrons in a chemical bond. It matters because it determines the type of bond that forms between two atoms (ionic, polar covalent, or nonpolar covalent), the polarity of molecules, and many chemical and physical properties such as solubility, boiling points, and reactivity. Understanding electronegativity is essential for predicting how elements will interact in chemical reactions and what kinds of compounds they will form.
Electronegativity describes an atom's ability to attract electrons within a chemical bond (a property of the bonded atom), while electron affinity is the energy change when an isolated gaseous atom gains an electron (a property of the free atom). They are related but distinct: chlorine has a higher electron affinity than fluorine, but fluorine has a higher electronegativity. Electronegativity is a dimensionless scale value, while electron affinity is measured in kJ/mol or eV.
Fluorine (EN = 3.98) is more electronegative than oxygen (EN = 3.44) because fluorine has a higher nuclear charge (9 protons vs. 8 for oxygen) while having the same number of electron shells. The additional proton in fluorine's nucleus pulls the bonding electrons more strongly. Additionally, fluorine needs only one more electron to achieve a noble gas configuration, whereas oxygen needs two, making fluorine's electron-attracting tendency in bonds even more pronounced.
The threshold values of 0.4 and 1.7 are approximate guidelines, not sharp boundaries. In reality, bond character transitions gradually from nonpolar covalent to polar covalent to ionic. A bond with a difference of exactly 0.4 or 1.7 is right at the boundary and could exhibit characteristics of both adjacent types. The percentage ionic character formula provides a more nuanced view. Different textbooks may use slightly different threshold values (some use 0.5 and 1.7, or 0.4 and 2.0), reflecting the approximate nature of these categories.
The Pauling scale is based on bond dissociation energies. Since noble gases (He, Ne, Ar) rarely form chemical bonds under normal conditions, there are insufficient bonding data to calculate Pauling electronegativity values for them. However, for the heavier noble gases that do form compounds -- particularly krypton (Kr = 3.00) and xenon (Xe = 2.60) -- Pauling values have been estimated based on their known fluoride compounds (KrF2, XeF2, XeF4, XeF6, etc.).
The electronegativity difference between oxygen (3.44) and hydrogen (2.20) is 1.24, creating polar covalent O-H bonds. Oxygen's high electronegativity pulls electron density toward itself, giving it a partial negative charge and leaving hydrogen with a partial positive charge. This polarity, combined with water's bent molecular geometry, gives water a net dipole moment. The polar nature of water molecules enables hydrogen bonding between them, which is responsible for water's unusually high boiling point, high heat capacity, surface tension, and its ability to dissolve ionic and polar substances.
Pauling's formula for estimating the percentage ionic character of a bond is: % Ionic Character = 100 × (1 - e-0.25Δχ2), where Δχ is the electronegativity difference between the two atoms. For example, an NaCl bond with Δχ = 2.23 gives approximately 67% ionic character. The formula produces a smooth curve from 0% (when atoms have identical electronegativities) approaching 100% for very large differences. This formula is an empirical approximation developed by Linus Pauling and has been refined by later researchers.
In organic chemistry, electronegativity is crucial for understanding reaction mechanisms, functional group behavior, and molecular properties. Electronegative atoms like oxygen, nitrogen, and halogens create polar bonds with carbon, generating electron-poor (electrophilic) and electron-rich (nucleophilic) sites that drive organic reactions. For example, in a carbonyl group (C=O), the electronegative oxygen makes the carbon electrophilic, susceptible to nucleophilic attack. Electronegativity also determines inductive effects (electron-withdrawing or electron-donating through sigma bonds), affects acidity and basicity, and influences the stability of reaction intermediates like carbocations and carbanions.