Electronegativity# - Biology


Molecules are collections of atoms that are associated with one another through bonds. It is reasonable to expect—and the case empirically—that different atoms will exhibit different physical properties, including abilities to interact with other atoms. While several methods for measuring electronegativity have been developed, the one most commonly taught to biologists is the one created by Linus Pauling.

A description of how Pauling electronegativity can be calculated is beyond the scope of BIS2A. What is important to know, however, is that electronegativity values have been experimentally and/or theoretically determined for nearly all elements in the periodic table. The values are unitless and are reported relative to the standard reference, hydrogen, whose electronegativity is 2.20. The larger the electronegativity value, the greater tendency an atom has to attract electrons. Using this scale, the electronegativity of different atoms can be quantitatively compared. For instance, by using Table 1 below, you could report that oxygen atoms are more electronegative than phosphorous atoms.

Table 1. Pauling electronegativity values for select elements of relevance to BIS2A as well as elements at the two extremes (highest and lowest) of the electronegativity scale.

Attribution: Marc T. Facciotti (original work)

The utility of the Pauling electronegativity scale in BIS2A is to provide a chemical basis for explaining the types of bonds that form between the commonly occurring elements in biological systems and to explain some of the key interactions that we observe routinely. We develop our understanding of electronegativity-based arguments about bonds and molecular interactions by comparing the electronegativities of two atoms. Recall, the larger the electronegativity, the stronger the "pull" an atom exerts on nearby electrons.

We can consider, for example, the common interaction between oxygen (O) and hydrogen (H). Let us assume that O and H are interacting (forming a bond) and write that interaction as O-H, where the dash between the letters represents the interaction between the two atoms. To understand this interaction better, we can compare the relative electronegativity of each atom. Examining the table above, we see that O has an electronegativity of 3.44, and H has an electronegativity of 2.20.

Based on the concept of electronegativity as we now understand it, we can surmise that the oxygen (O) atom will tend to "pull" the electrons away from the hydrogen (H) when they are interacting. This will give rise to a slight but significant negative charge around the O atom (due to the higher tendency of the electrons to be associated with the O atom). This also results in a slight positive charge around the H atom (due to the decrease in the probability of finding an electron nearby). Since the electrons are not distributed evenly between the two atoms AND, by consequence, the electric charge is also not evenly distributed, we describe this interaction or bond as polar. There are two poles in effect: the negative pole near the oxygen and the positive pole near the hydrogen.

To extend the utility of this concept, we can now ask how an interaction between oxygen (O) and hydrogen (H) differs from an interaction between sulfur (S) and hydrogen (H). That is, how does O-H differ from S-H? If we examine the table above, we see that the difference in electronegativity between O and H is 1.24 (3.44 - 2.20 = 1.24) and that the difference in electronegativity between S and H is 0.38 (2.58 – 2.20 = 0.38). We can therefore conclude that an O-H bond is more polar than an S-H bond. We will discuss the consequences of these differences in subsequent chapters.

Figure 2. The periodic table with the electronegativities of each atom listed.

Attribution: By DMacks ( [CC BY-SA 3.0 (], via Wikimedia Commons

An examination of the periodic table of the elements (Figure 2) illustrates that electronegativity is related to some of the physical properties used to organize the elements into the table. Certain trends are apparent. For instance, those atoms with the largest electronegativity tend to reside in the upper right hand corner of the periodic table, such as Fluorine (F), Oxygen (O) and Chlorine (Cl), while elements with the smallest electronegativity tend to be found at the other end of the table, in the lower left, such as Francium (Fr), Cesium (Cs) and Radium (Ra).

The main use of the concept of electronegativity in BIS2A will therefore be to provide a conceptual grounding for discussing the different types of chemical bonds that occur between atoms in nature. We will focus primarily on three types of bonds: Ionic Bonds, Covalent Bonds and Hydrogen Bonds.

What if two atoms of equal electronegativity bond together?

Consider a bond between two atoms, A and B. If the atoms are equally electronegative, both have the same tendency to attract the bonding pair of electrons, and so it will be found on average half way between the two atoms:

To get a bond like this, A and B would usually have to be the same atom. You will find this sort of bond in, for example, H2 or Cl2 molecules. Note: It's important to realize that this is an average picture. The electrons are actually in a molecular orbital, and are moving around all the time within that orbital. This sort of bond could be thought of as being a "pure" covalent bond - where the electrons are shared evenly between the two atoms.

Electronegativity# - Biology

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From Backwater to Center Stage: Using Electronegativity as a Central Concept for Understanding Chemical Principles in Biology Classes

* Roger Sauterer is Associate Professor of Biology at Jacksonville State University, 700 Pelham Road North, Jacksonville, AL 36265 e-mail: [email protected]

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Understanding basic chemical concepts, including bonding, polar and nonpolar molecules, and hydrogen bonds is difficult for many biology students, who often have minimal chemistry backgrounds. The concept of electronegativity is introduced at the beginning of the chemical foundations part of a biology course as a central integrative concept. By using the electronegativity concept and an associated line graph, students gain an understanding of why ionic and covalent bonds form and which atoms form them, why atoms form polar and nonpolar covalent bonds, and what chemical groups can form hydrogen bonds. Positive student reviews indicate that this is an effective method for introducing chemical principles.


The electronegativity of an atom is a measure of its affinity for electrons.

The atoms of the various elements differ in their affinity for electrons.

This image distorts the conventional periodic table of the elements so that the greater the electronegativity of an atom, the higher its position in the table.

Although fluorine (F) is the most electronegative element, it is the electronegativity of runner-up oxygen (O) that is exploited by life.

  • The shuttling of electrons between carbon (C) and oxygen (O) atoms powers life.
    1. Moving electrons against the gradient (O to C) — as occurs in photosynthesis — requires energy (and stores it).
    2. Moving electrons down the gradient (C to O) — as occurs in cellular respiration — releases energy.
  • The relative electronegativity of two interacting atoms also plays a major part in determining what kind of chemical bond forms between them.

Example 1: Sodium (Na) and Chlorine (Cl) = Ionic Bond

  • There is a large difference in electronegativity, so
  • the chlorine atom takes an electron from the sodium atom
  • converting the atoms into ions (Na + ) and (Cl − ).
  • These are held together by their opposite electrical charge forming ionic bonds.
  • Each sodium ion is held by 6 chloride ions while each chloride ion is, in turn, held by 6 sodium ions.
  • Result: a crystal lattice (not molecules) of common table salt (NaCl).

Example 2: Carbon (C) and Hydrogen (H) = Covalent Bond

  • There is only a small difference in electronegativity, so
  • the two atoms share the electrons.
  • Result: a covalent bond (depicted as C:H or C-H).
  • The atoms are held together by their mutual affinity for their shared electrons.
  • An array of atoms held together by covalent bonds forms a true molecule.

Example 3: Hydrogen (H) and Oxygen (O) = Polar Covalent Bond

  • Here there is a moderate difference in electronegativity, causing
  • the oxygen atom to pull the electron of the hydrogen atom closer to itself.
  • Result: a polar covalent bond.
  • Oxygen does this with 2 hydrogen atoms to form a molecule of water.

Molecules, like water, with polar covalent bonds

  • are themselves polar that is, have partial electrical charges across the molecule
  • may be attracted to each other (as occurs with water molecules) [Link to schematic]
  • are good solvents for polar and/or hydrophilic compounds. [Link to a schematic of how water molecules dissolve a crystal of table salt (NaCl)]
  • may form hydrogen bonds.
  • You are here:  
  • Home
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  • Kimball's Biology (supplemental textbook for Biol-58x Sequence)
  • The Chemical Basis of Life
  • Electronegativity

This work is licensed under a Creative Commons Attribution-NonCommercial 3.0 Unported License.

Electronegativity in the periodic table

The Pauling scale is most commonly used to measure electronegativity. According to the scale Fluorine is the most electronegative element and its value is 4.0. Caesium and Francium are the least electronegative element and its value is 0.7.

Here we can see that the electronegativity increases across the period and it decreases down the group. This can be explained with atomic number and the distance of valence electrons form nucleus.

Electronegativity increases across the period

Moving from left to right across a period, the number of protons and electrons increases while the number of energy shells stay same. Thus because of more attraction between increasing number of positive nucleus and negative electrons, the atomic radius decreases and the electronegativity (attraction of nucleus towards electrons) increases across the period.

Electronegativity decreases down the group

Moving down a group, the number of energy shells also increases with the increase of protons and electrons. So because of shielding effects of electrons in the increased inner shells, the attraction between electrons and nucleus (electronegativity) is reduced down the group.

Diagonal relationships in the periodic table

A diagonal relationship is said to exist between certain pairs of diagonally (not side by side) adjacent elements in second and third period in the periodic table.

For example, lithium and magnesium, beryllium and aluminium, boron and silicon have similar properties. Boron and silicon both act as a semiconductors. This can be explained in terms of electronegativity.

We already have seen that electronegativity increases across the period in periodic table. So, for example, the electronegativity of boron is more than beryllium.

Again it decreases down the period, so the electronegativity of aluminium is less than boron. Thus berylium and aluminium both have less electronegativity than boron. This way the diagonal elements have similar (in this case same) electronegativities and similar properties than others.


Linus Pauling’s electronegativity scale was inspired by Biology. In the early 1930s, chromosomal genes were being mapped out by measuring how frequently two independent traits were inherited together the idea being that the closer the genes were, the more likely that they would stay linked during genetic crossover.

Pauling tested this idea with chemical compounds, finding that bonds between similar elements were not as strong as bonds between dissimilar elements. He attributed this discovery to ionic contributions in the stronger bonds, and correlated the ionic nature of certain elements with further spreads on his electronegativity scale. For example, the bond between Lithium and Fluorine was almost one-hundred percent ionic – therefore, he placed Lithium on one end of his scale and Fluorine on the other end.

From there, Pauling assigned arbitrary values for each known element based upon their position on the accepted ‘map’ of ionic bond proclivities. Later, he explained that his calculation of each electronegativity value was an estimate of the covalent contribution to an element’s bond subtracted from the actual bond energy, as per the following formula:

In this formula, Δ is a measure of excess ionic energy – the value that Pauling used to arbitrarily assign electronegativity values to elements. Again, the higher the ionic bond energy measured within an element, the more electronegative the element was to be considered.

In terms of chemistry, Pauling’s electronegativity scale was one of his least theoretically well-founded theories. On the very same token, it was also one of his most influential ideas in that it allowed chemists to make assumptions about bonds and molecules that could give rise to new interesting and useful correlations.

Indeed, Pauling’s electronegativity scale was very practical. He used electronegativity to explain chemical bonding characteristics, including the changes in the energy of atoms that occur as electrons rearrange their placement in the atoms’ orbitals. By comparing these values, researchers could predict the properties of a given bond without ever needing to know the bond’s complicated wave equation from quantum mechanics.

Pauling’s faith in his scale was such that he used it to theorize that Fluorine was so electronegative, it would form compounds with an inert gas – something that, at the time, was thought to be impossible. Inert gasses simply did not bond. However, he couldn’t prove the relationship, and it frustrated Pauling. Eventually though, some thirty years later, he was proven correct by another team of scientists. In their discussions of Pauling’s “stochastic method,” biographers have shown that much of Pauling’s research followed the Fluorine example: more often that not, his intuition about chemical systems was correct, despite his inability to empirically prove his ideas with hard data.

Electronegativity data, 1930s.

In 1932, Linus Pauling published his original paper proposing a thermochemical method of assigning relative electronegativity values. He applied his system to ten nonmetallic elements. As with Berzelius’ earlier attempts at developing an electronegativity scale, Pauling failed to clearly define how he established his proposed values. (For the contemporary student, his later calculations regarding electronegativity are contained in his 37th research journal, though to most readers the journal is difficult to follow due to it being more of a stream-of-conscious study as opposed to a series of well-explained experimental argument.)

All of this noted, and despite many additional attempts at determining a rigorous electronegativity scale in the years following his work, Pauling’s 1932 scale is still the one most-commonly in use today.

Electronegativity Definition

Electronegativity is a chemical property that measures how likely an atom is to attract a shared pair of electrons towards itself in a covalent bond.

Electronegativity is important because it makes bonding between atoms possible. The higher the electronegativity, the greater an atom’s propensity to attract electrons.

Atoms form molecular compounds by combining with other atoms. Electronegativity determines how the bonds between atoms exist. The greater the difference between the electronegativity values of different atoms, the more polar the chemical bond formed between them is.

Electronegativity is not stagnant - it can depend on an atom’s environment. That being said, most atoms display similar electronegativity behavior no matter their environments, so there are common scales used to calculate electronegativity.

Electronegativity and Bond Polarity

Although we defined covalent bonding as electron sharing, the electrons in a covalent bond are not always shared equally by the two bonded atoms. Unless the bond connects two atoms of the same element, as in H2, there will always be one atom that attracts the electrons in the bond more strongly than the other atom does, as in HCl, shown in Figure (PageIndex<1>). A covalent bond that has an equal sharing of electrons (Figure (PageIndex<1a>)) is called a nonpolar covalent bond. A covalent bond that has an unequal sharing of electrons, as in Figure (PageIndex<1b>), is called a polar covalent bond.

Figure (PageIndex<1>) Polar versus Nonpolar Covalent Bonds. (a) The electrons in the covalent bond are equally shared by both hydrogen atoms. This is a nonpolar covalent bond. (b) The chlorine atom attracts the electrons in the bond more than the hydrogen atom does, leading to an imbalance in the electron distribution. This is a polar covalent bond.

The distribution of electron density in a polar bond is uneven. It is greater around the atom that attracts the electrons more than the other. For example, the electrons in the H&ndashCl bond of a hydrogen chloride molecule spend more time near the chlorine atom than near the hydrogen atom. Note that the shaded area around Cl in Figure (PageIndex<1b>) is much larger than it is around H.

This imbalance in electron density results in a buildup of partial negative charge (designated as &delta&minus) on one side of the bond (Cl) and a partial positive charge (designated &delta+) on the other side of the bond (H). This is seen in Figure (PageIndex<2a>). The separation of charge in a polar covalent bond results in an electric dipole (two poles), represented by the arrow in Figure (PageIndex<2b>). The direction of the arrow is pointed toward the &delta&minus end while the + tail of the arrow indicates the &delta+ end of the bond.

Figure (PageIndex<2>): (a) Unequal sharing of the bonding pair of electrons between H and Cl leads to partial positive charge on the H atom and partial negative charge on the Cl. Symbols &delta+ and &delta&ndash indicate the polarity of the H&ndashCl bond. (b) The dipole is represented by an arrow with a cross at the tail. The cross is near the &delta+ end and the arrowhead coincides with the &delta&ndash.

Any covalent bond between atoms of different elements is a polar bond, but the degree of polarity varies widely. Some bonds between different elements are only minimally polar, while others are strongly polar. Ionic bonds can be considered the ultimate in polarity, with electrons being transferred rather than shared. To judge the relative polarity of a covalent bond, chemists use electronegativity, which is a relative measure of how strongly an atom attracts electrons when it forms a covalent bond. There are various numerical scales for rating electronegativity. Figure (PageIndex<3>) shows one of the most popular&mdashthe Pauling scale.

Figure (PageIndex<3>) The electronegativity values derived by Pauling follow predictable periodic trends with the higher electronegativities toward the upper right of the periodic table. Fluorine has the highest value (4.0).

Looking Closer: Linus Pauling

Arguably the most influential chemist of the 20th century, Linus Pauling (1901&ndash94) is the only person to have won two individual (that is, unshared) Nobel Prizes. In the 1930s, Pauling used new mathematical theories to enunciate some fundamental principles of the chemical bond. His 1939 book The Nature of the Chemical Bond is one of the most significant books ever published in chemistry.

By 1935, Pauling&rsquos interest turned to biological molecules, and he was awarded the 1954 Nobel Prize in Chemistry for his work on protein structure. (He was very close to discovering the double helix structure of DNA when James Watson and James Crick announced their own discovery of its structure in 1953.) He was later awarded the 1962 Nobel Peace Prize for his efforts to ban the testing of nuclear weapons.

Linus Pauling was one of the most influential chemists of the 20th century.

In his later years, Pauling became convinced that large doses of vitamin C would prevent disease, including the common cold. Most clinical research failed to show a connection, but Pauling continued to take large doses daily. He died in 1994, having spent a lifetime establishing a scientific legacy that few will ever equal.

The polarity of a covalent bond can be judged by determining the difference in the electronegativities of the two atoms making the bond. The greater the difference in electronegativities, the greater the imbalance of electron sharing in the bond. Although there are no hard and fast rules, the general rule is if the difference in electronegativities is less than about 0.4, the bond is considered nonpolar if the difference is greater than 0.4, the bond is considered polar. If the difference in electronegativities is large enough (generally greater than about 1.8), the resulting compound is considered ionic rather than covalent. An electronegativity difference of zero, of course, indicates a nonpolar covalent bond.

Figure (PageIndex<4>): As the electronegativity difference increases between two atoms, the bond becomes more ionic.

Describe the electronegativity difference between each pair of atoms and the resulting polarity (or bond type).

  1. Carbon has an electronegativity of 2.5, while the value for hydrogen is 2.1. The difference is 0.4, which is rather small. The C&ndashH bond is therefore considered nonpolar.
  2. Both hydrogen atoms have the same electronegativity value&mdash2.1. The difference is zero, so the bond is nonpolar.
  3. Sodium&rsquos electronegativity is 0.9, while chlorine&rsquos is 3.0. The difference is 2.1, which is rather high, and so sodium and chlorine form an ionic compound.
  4. With 2.1 for hydrogen and 3.5 for oxygen, the electronegativity difference is 1.4. We would expect a very polar bond. The sharing of electrons between O and H is unequal with the electrons more strongly drawn towards O.

Describe the electronegativity (EN) difference between each pair of atoms and the resulting polarity (or bond type).

The EN difference is 1.0 , hence polar. The sharing of electrons between C and O is unequal with the electrons more strongly drawn towards O.

The EN difference is greater than 1.8, hence ionic.

Identical atoms have zero EN difference, hence nonpolar.

The EN difference is greater than 1.8, hence ionic.

Meaning and definition of electronegativity :

The attraction of an atom for the electrons of a covalent bond.

For the term electronegativity may also exist other definitions and meanings, the meaning and definition indicated above are indicative not be used for medical and legal or special purposes.

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