Lesson Explainer: Electronegativity Chemistry

In this explainer, we will learn how to explain the chemical property of electronegativity.

Electronegativity measures the tendency of an atom to attract a bonding pair of electrons. Some atoms have relatively high electronegativity values, and they are prone to withdrawing a large amount of electron density from a bonding pair of electrons. Other atoms have much lower electronegativity values, and they are less prone to withdrawing a significant amount of electron density from a bonding pair of electrons. The following image shows how the highly electronegative fluorine atom withdraws most of the electron density (red-to-blue colored cloud) from a hydrogen–fluoride covalent bond.

Definition: Electronegativity

Electronegativity measures the tendency of an atom to attract a bonding pair of electrons from a chemical bond.

Linus Pauling proposed an electronegativity scale for the periodic table elements that depends on bond dissociation energies. Chemists can use bond dissociation data for different homonuclear and heteronuclear compounds to determine differences in electronegativity values. The process is somewhat convoluted and it is beyond the scope of this explainer. We just need to remember that Pauling found a clever way to use bond dissociation data to determine differences in electronegativity values. The following figure shows the periodic table in terms of Pauling scale electronegativity numbers. It is important to note that data is not available for some noble gas elements and many of the most recently discovered elements. Data is not available for any of the three lightest noble gas atoms and also data is not available for the heaviest and least stable elements that make up the bottommost row of the periodic table.

Electronegativity values generally increase as we move from the left-hand side of the periodic table to the right-hand side and it also increases as we move from the bottom of the periodic table to the top. Group 17 elements tend to have a relatively high electronegativity number and group 1 elements tend to have a relatively low electronegativity number. Period 1 and period 2 elements tend to have a relatively high electronegativity value and period 6 elements tend to have a relatively low electronegativity value.

Francium and cesium are found in the bottom-left corner of the periodic table and they have the lowest electronegativity values of all the chemical elements. Cesium currently has the lowest recorded electronegativity value and it is predicted that francium has an electronegativity value that is just as low or slightly higher. This is challenging to prove outright because scientists cannot isolate enough francium to definitively determine its electronegativity value. Its electronegativity can merely be estimated through the use of known ionization energy data. Fluorine is found close to the top right-hand corner of the periodic table and it currently has the highest recorded Pauling electronegativity number. Helium and neon might have exceptionally high Pauling electronegativity numbers, but this has not been proved outright. Scientists cannot obtain the bond dissociation data they need to determine the Pauling electronegativity of helium, neon, or argon. This inability to quantify the Pauling electronegativity number of the three lightest noble gas atoms is explained in the next couple of paragraphs.

Example 1: Identifying Which Elements Have the Highest Electronegativity Numbers

Which letter corresponds to the elements in the periodic table that are the most electronegative?

Answer

Electronegativity values measure the tendency of an atom to attract a bonding pair of electrons from a chemical bond. Pauling electronegativity numbers generally increase as we move from the left-hand side of the periodic table to the right-hand side and as we move from the bottom of the periodic table to the top. The figure uses four letters to represent what could be called the four corners of the periodic table of elements. A and B are used to represent the top-left and top-right corners of the periodic table and C and D are used to represent the bottom-left and bottom-right corners. The letter B must show the location of the elements that have the highest Pauling electronegativity numbers, because it is both high up the periodic table and far to the right side of the periodic table.

The Pauling electronegativity values quantify the tendency of an atom to attract a bonding pair of electrons. Most atoms have an incomplete valence shell and they withdraw bonding pairs of electrons to have eight valence shell electrons and the same electron configuration as the nearest noble gas.

The noble gas elements already have stable electron configurations and they do not tend to form polyatomic compounds at room temperature and atmospheric pressure. Helium tends to remain completely unbonded if the pressure is kept below 1 gigapascal. Neon tends to only make the so-called van der Waals molecules. Argon tends to only make unstable diatomic hydrides through intense irradiation processes. The three lightest noble gases do not form conventional chemical bonds under standard conditions and chemists cannot get the bond dissociation data they need to quantify their Pauling electronegativity values.

Example 2: Understanding Why Some Noble Gas Elements Do Not Have Electronegativity Values

Why do no values exist on the Pauling scale of electronegativity for argon, neon, and helium?

  1. These noble gases have too great an electron density.
  2. These noble gases are synthetic elements and do not exist in quantities that are large enough to be measured.
  3. These noble gases require too much energy to be ionized.
  4. These noble gases do not form bonds; therefore, no bond dissociation energy data is available.
  5. These noble gases exist as electrically neutral atoms.

Answer

Linus Pauling proposed an electronegativity scale for the chemical elements that is based on the value of bond dissociation energies. Chemists have to obtain the bond dissociation data for different homonuclear and heteronuclear compounds if they want to determine differences in electronegativity values. The differences in electronegativity values can then be processed to determine the Pauling electronegativity values for individual chemical elements. The three lightest noble gases do not form conventional chemical bonds under standard conditions and chemists cannot get the bond dissociation data they need to quantify their Pauling electronegativity values. This explains why some of the noble gas elements do not have Pauling electronegativity numbers when they are shown in normal periodic tables or in Pauling electronegativity tables. These statements are summarized quite succinctly in answer D.

Electronegativity values tend to increase with atomic number across a period. Elements on the right-hand side of a period have a higher atomic number and a higher effective nuclear charge (𝑍)eff. This means that they tend to have smaller atoms. Smaller atoms are more effective at β€œpulling” valence electrons from chemical bonds.

Electronegativity values tend to increase up a group. The elements at the top of a group have a lower number of core electrons and they tend to have a smaller atomic radius. The elements at the top of a group have relatively high electronegativity values because they are small and smaller atoms are highly effective at β€œpulling” electrons from chemical bonds.

Electronegativity values are related to first ionization energy values because both atomic parameters depend on atomic radii and effective nuclear charge. Both atomic parameters tend to be large when elements have small atoms and a high effective nuclear charge. This explains why electronegativity and first ionization energies similarly tend to increase from the bottom-left corner of the periodic table to the top-right corner of the periodic table. Similar comparisons can be made between electronegativity and electron affinity parameters. Both of these atomic parameters tend to increase toward the top-right corner of the periodic table because they both depend on atomic radii and effective nuclear charge. Electronegativity and electron affinity parameters tend to be high when elements have small atoms and a high effective nuclear charge.

Example 3: Identifying Which Element Has a Similar Electronegativity to Aluminum

Which group 2 element is likely to have a similar electronegativity to aluminum?

  1. Beryllium
  2. Calcium
  3. Strontium
  4. Barium
  5. Magnesium

Answer

Electronegativity values tend to increase with atomic number across a period. Electronegativity values also tend to increase up a group. This means that electronegativity values tend to increase as we move from the bottom left-hand side of the periodic table to the top right-hand side. The following image shows the periodic table of elements.

It is clear that aluminum is a group 13 element that can be found in the third row of the periodic table. Elements should have the same or a similar electronegativity as aluminum if they are close to it in the periodic table. Elements should have a much lower electronegativity value than aluminum if they are right next to or close to cesium or francium.

The list shows five elements from group 2 of the periodic table. Beryllium and magnesium are relatively close to aluminum, but the other listed options are not. The other listed options can be discounted because they are too close to cesium and francium.

Beryllium should have a slightly higher electronegativity value than magnesium because beryllium is situated directly above magnesium in the periodic table. Aluminum should have a slightly higher electronegativity value than magnesium because it is effectively situated just to the right of magnesium. Aluminum and beryllium must have the more similar electronegativity values because they both have an electronegativity value that is slightly higher than the electronegativity of magnesium. This line of reasoning can be used to determine that option A is the correct answer for this question.

The Pauling scale of electronegativity can be used to understand why some atoms form simple molecular compounds and other atoms form much larger and more elaborate giant ionic lattices. The bonding type is determined from the absolute difference of electronegativity values.

Compounds are usually covalently bonded when they contain elements that have the same or very similar electronegativity values. Compounds are almost always ionically bonded when they contain one metal element that has a very low electronegativity value and another nonmetal that has a much higher electronegativity value.

Definition: Absolute Difference of Electronegativity Values (Δ𝐸𝑁)

The absolute difference of electronegativity values can be defined as Δ𝐸𝑁=|πΈβˆ’πΈ|, where 𝐸 and 𝐸 are the electronegativity values of two chemically bonded elements.

Bond TypeElectronegativity DifferenceExample Molecules
Pure covalent0Hydrogen gas (H)2
Nonpolar covalentLess than 0.4Methane (CH)4
Polar CovalentBetween 0.4 and 1.7Hydrogen fluoride (HF)
IonicGreater than 1.7Sodium fluoride (NaF)

Compounds are usually covalently bonded when the difference of electronegativity values is less than 1.7 and they tend to be ionically bonded when the difference of electronegativity values is greater than 1.7. The absolute difference of electronegativity values can even be used to understand why simple molecular compounds can sometimes form nonpolar molecules and other times form polar molecules. Compounds usually form nonpolar covalent compounds if the difference of electronegativity values is less than 0.4. Compounds usually form polar covalent compounds if the difference of electronegativity values is greater than 0.4 but less than 1.7.

Example 4: Predicting the Type of Bond That Is Formed between Two Different Elements

The difference in electronegativity between two atoms of different elements is 0.2, using values from the Pauling scale. What type of bond is likely to form between these two atoms in a chemical reaction?

  1. A nonpolar covalent bond
  2. A polar bond
  3. An ionic bond

Answer

The absolute difference of electronegativity values can be defined with the equation Δ𝐸𝑁=|πΈβˆ’πΈ|, where 𝐸 and 𝐸 are the electronegativity values of two chemically bonded elements. Compounds tend to be covalently bonded when Δ𝐸𝑁 is smaller than 1.7 and ionically bonded when Δ𝐸𝑁 is greater than 1.7. The absolute difference of electronegativity values can also be used to determine if bonding pairs of electrons are shared equally or unequally between covalently bonded atoms. Compounds tend to make nonpolar covalent bonds when Δ𝐸𝑁 is smaller than 0.4 and polar covalent bonds when Δ𝐸𝑁 is greater than 0.4 but also smaller than 1.7.

The question discusses two atoms that have an absolute difference of electronegativity values that is less than 0.4. It discusses two atoms that have an absolute difference of electronegativity values that is equal to 0.2. The two atoms can be expected to form a nonpolar covalent bond because Δ𝐸𝑁<0.4. The previous statements can be used to determine that option A is the correct answer for this question.

Chemists can look at the electronegativity values of two reacting elements to predict what type of chemical products can be formed. For example, sodium has an electronegativity value of 0.93 and chlorine has an electronegativity value of 3.16. The absolute difference of electronegativity values must be 2.23 for the compound that is formed between sodium and chlorine because Δ𝐸𝑁=|πΈβˆ’πΈ|=2.23 when 𝐸=3.16 and 𝐸=0.93. The value of 2.23 is much higher than the 1.70 cutoff point for covalent compounds. This suggests that sodium and chlorine atoms must react together and form an ionic salt that would have to be called sodium chloride (NaCl).

Example 5: Determining Which System Has the Greatest Difference of Electronegativity Values

Which of the following diatomic molecules is likely to have the greatest difference in electronegativity between its atoms?

  1. HBr
  2. O2
  3. CO
  4. CsF
  5. HF

Answer

The difference of electronegativity values is determined with the equation Δ𝐸𝑁=|πΈβˆ’πΈ|, where 𝐸 and 𝐸 are the electronegativity values of two chemically bonded elements. Option B is the diatomic oxygen (O)2 molecule. Option B cannot be the correct answer for this question because oxygen molecules contain two atoms that have identical electronegativity values. Option C is the carbon monoxide molecule (CO). Option C cannot be the correct answer for this question because carbon monoxide (CO) molecules contain carbon and oxygen. Carbon and oxygen are both nonmetal elements and they are both period 2 elements. They should have very similar electronegativity values. Option A is the hydrogen bromide molecule (HBr) and option E is the hydrogen fluoride molecule (HF). Option A can be discounted through comparison with option E. Hydrogen fluoride and hydrogen bromide are both hydrogen halide compounds, but hydrogen fluoride contains the highly electronegative fluorine element and hydrogen bromide contains the much less electronegative bromine element. Hydrogen fluoride can be expected to have the larger Δ𝐸𝑁 value because it contains the more highly electronegative halogen element. Option D is the cesium fluoride molecule (CsF). Option E can be discounted through comparison with option D. Hydrogen fluoride and cesium fluoride both contain the highly electronegative fluorine atom, but fluorine is bonded to cesium in cesium fluoride and it is bonded to hydrogen in hydrogen fluoride. Cesium has a much lower electronegativity value than hydrogen because it is much closer to the bottom left-hand side corner of the periodic table. Cesium fluoride must have the larger Δ𝐸𝑁 value because cesium has a much lower electronegativity value than hydrogen. This line of reasoning indicates that cesium fluoride must be the correct answer for this question. We can conclude that option D is the correct answer.

The following section shows how we can determine the answer in a different way. It shows how the answer can be determined via the direct calculation of electronegativity difference numbers with the Δ𝐸𝑁=|πΈβˆ’πΈ| formula. The following table shows the Pauling electronegativity values that are needed to determine the Δ𝐸𝑁 value for all of the listed compounds. This data can be found in the first figure of this explainer. The absolute difference of electronegativity values is determined with the Δ𝐸𝑁=|πΈβˆ’πΈ| formula.

ElementPauling Electronegativity
F3.98
O3.44
Br2.96
C2.55
H2.20
Cs0.79

The first Δ𝐸𝑁 calculation is for hydrogen bromide (HBr). Hydrogen bromide has an Δ𝐸𝑁 value of 0.76 because Δ𝐸𝑁=0.76 when 𝐸=2.96 and 𝐸=2.20. The second Δ𝐸𝑁 calculation is for diatomic oxygen (O)2. Diatomic oxygen has an Δ𝐸𝑁 value of 0.00 because Δ𝐸𝑁=0.00 when 𝐸=3.44 and 𝐸=3.44. The third Δ𝐸𝑁 calculation is for carbon monoxide (CO). Carbon monoxide has an Δ𝐸𝑁 value of 0.89 because Δ𝐸𝑁=0.89 when 𝐸=3.44 and 𝐸=2.55. The fourth Δ𝐸𝑁 calculation is cesium fluoride (CsF). Cesium fluoride has an Δ𝐸𝑁 value of 3.19 because Δ𝐸𝑁=3.19 when 𝐸=3.98 and 𝐸=0.79. The fifth Δ𝐸𝑁 calculation is for hydrogen fluoride (HF). Hydrogen fluoride has an Δ𝐸𝑁 value of 1.78 because Δ𝐸𝑁=1.78 when 𝐸=3.98 and 𝐸=2.20. These calculations can be compared with each other to determine that cesium fluoride has the highest absolute difference of electronegativity value. Two different lines of reasoning have been applied here to show that option D is the correct answer for this question.

Let us summarize what has been learned in this explainer regarding electronegativity.

Key Points

  • Electronegativity measures the tendency of an atom to attract a bonding pair of electrons from a chemical bond.
  • Electronegativity values tend to increase with atomic number across a period because the rightmost elements have smaller atoms than elements on their left.
  • Electronegativity values tend to increase up a group because elements near the top of the periodic table are small and elements near the bottom are large.
  • The noble gases have some of the most stable electron configurations and some of them do not have electronegativity values.
  • The bonding type is inextricably linked to the absolute difference of electronegativity values.
  • Two elements will tend to form ionic compounds when the absolute difference of their electronegativity values is greater than 1.7.
  • Two elements will tend to make nonpolar covalent compounds when the absolute difference of their electronegativity values is less than 0.4.
  • Two elements will tend to make polar covalent compounds when the absolute difference of their electronegativity values is between 0.4 and 1.7.

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