What Is Electronegativity?

Electronegativity is a measure of an atom's ability to attract bonding electrons toward itself when it is chemically bonded to another atom.

When two atoms form a covalent bond, they share a pair of electrons.

However, those electrons are not always shared equally.

If one atom has a greater electronegativity, it attracts the shared electrons more strongly.

For example, in an H–Cl bond, chlorine attracts the bonding electrons more strongly than hydrogen.

This produces an unequal distribution of charge.

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Electronegativity Is About Bonding Electrons

Electronegativity describes how strongly an atom attracts electrons in a chemical bond.

Imagine two atoms connected by a rope, with the bonding electrons represented by the middle of the rope.

If both atoms pull equally strongly, the electrons remain approximately evenly shared.

If one atom pulls more strongly, the electrons spend more time closer to that atom.

The greater the difference in electronegativity, the more unequal the sharing can become.

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The Electronegativity Scale

Electronegativity is commonly described using the Pauling scale.

Fluorine has the highest electronegativity:

F ≈ 3.98

Some commonly encountered approximate values are:

H = 2.20

C = 2.55

N = 3.04

O = 3.44

F = 3.98

Na = 0.93

Cl = 3.16

These values allow us to compare how strongly different atoms attract bonding electrons.

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Fluorine Is the Most Electronegative Element

Fluorine is located near the upper-right corner of the periodic table.

It has:

• a relatively small atomic radius
• relatively little shielding
• a strong effective nuclear attraction for bonding electrons

As a result, fluorine attracts bonding electrons extremely strongly.

This makes fluorine the most electronegative element.

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Electronegativity Across a Period

As we move from left to right across a period:

electronegativity generally increases.

For example, across Period 2:

Li → Be → B → C → N → O → F

the general electronegativity trend is upward.

Elements toward the right side of a period attract bonding electrons more strongly than elements toward the left.

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Why Electronegativity Increases Across a Period

Moving across a period:

• proton number increases
• electrons are added to the same main energy level
• shielding does not increase greatly
• effective nuclear attraction increases
• atomic radius generally decreases

Therefore, the nucleus can attract bonding electrons more strongly.

The result is:

electronegativity increases across a period.

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Effective Nuclear Attraction

The nucleus contains positively charged protons.

Electrons are negatively charged.

Therefore, the nucleus attracts electrons through electrostatic forces.

Across a period, the number of protons increases while electrons are added to the same main shell.

This generally increases the effective attraction experienced by electrons involved in bonding.

That stronger attraction contributes to higher electronegativity.

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Electronegativity Down a Group

As we move down a group:

electronegativity generally decreases.

For example:

F > Cl > Br > I

Fluorine attracts bonding electrons more strongly than chlorine.

Chlorine attracts them more strongly than bromine.

Bromine attracts them more strongly than iodine.

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Why Electronegativity Decreases Down a Group

Moving down a group:

• more occupied electron shells are added
• atomic radius increases
• shielding increases
• bonding electrons are generally farther from the nucleus
• effective nuclear attraction for bonding electrons decreases

Therefore:

electronegativity decreases down a group.

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Electron Shielding

Inner electrons produce an effect called shielding.

The inner electrons partially reduce the attraction experienced by electrons farther from the nucleus.

Down a group, additional electron shells create greater shielding.

This is one reason atoms farther down a group attract bonding electrons less strongly.

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Atomic Radius and Electronegativity

Atomic size is closely related to electronegativity.

Smaller atoms generally have bonding electrons closer to their nuclei.

This allows stronger attraction.

Larger atoms have bonding electrons farther from their nuclei and experience greater shielding.

Therefore, the broad relationship is:

smaller atomic radius → higher electronegativity

larger atomic radius → lower electronegativity

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The Overall Electronegativity Trend

Combining the two trends:

Across a period → electronegativity increases

Down a group ↓ electronegativity decreases

Therefore, electronegativity generally increases toward the:

upper-right corner of the periodic table

with fluorine having the highest value.

This pattern allows us to compare elements without memorizing every electronegativity value.

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Electronegativity and the Periodic Table

The general trend can be summarized as:

lower-left → lower electronegativity

upper-right → higher electronegativity

For example:

Potassium is toward the lower-left and has relatively low electronegativity.

Fluorine is toward the upper-right and has extremely high electronegativity.

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What About Noble Gases?

Noble gases are located in Group 18.

They already have very stable outer electron arrangements and do not commonly form the same kinds of bonds as many other main-group elements.

Because electronegativity specifically concerns attraction for bonding electrons, electronegativity values for noble gases are often omitted in introductory periodic tables.

Some noble gases can form compounds under particular conditions, so the full chemistry is more complex than simply saying that noble gases have "no electronegativity."

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Comparing Elements in the Same Period

Compare:

Na and Cl

Both are in Period 3.

Chlorine is farther to the right.

Electronegativity generally increases across a period.

Therefore:

Cl is more electronegative than Na.

Chlorine attracts bonding electrons much more strongly than sodium.

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Comparing Elements in the Same Group

Compare:

F and Br

Both are in Group 17.

Fluorine is higher in the group.

Electronegativity decreases down a group.

Therefore:

F is more electronegative than Br.

Fluorine has fewer electron shells, less shielding, and a smaller atomic radius.

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Comparing Elements in Different Groups and Periods

Suppose we compare:

Mg and O

Oxygen is:

• farther to the right
• higher in the periodic table

Both trends indicate that oxygen should have the greater electronegativity.

Therefore:

O is more electronegative than Mg.

This type of reasoning allows us to predict relative electronegativity without knowing exact numerical values.

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Worked Example 1

Which is more electronegative?

C or O

Both are in Period 2.

Oxygen is farther to the right.

Therefore:

O > C

in electronegativity.

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Worked Example 2

Which is more electronegative?

Cl or I

Both are in Group 17.

Chlorine is higher in the group.

Therefore:

Cl > I

in electronegativity.

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Worked Example 3

Which is more electronegative?

K or Br

Bromine is farther right and higher than potassium.

Therefore:

Br is much more electronegative than K.

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Ranking Electronegativity

Suppose we need to arrange:

Na, Mg, Si, Cl

from lowest to highest electronegativity.

All four are in Period 3.

Electronegativity generally increases from left to right.

Therefore:

Na < Mg < Si < Cl

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Electronegativity and Chemical Bonds

Electronegativity becomes especially useful when predicting how atoms will behave when bonded.

If two atoms have similar electronegativities, they tend to share electrons relatively evenly.

If their electronegativities are different, the electrons are shared unequally.

If the difference is very large, electron transfer and ionic bonding may become a useful model.

This allows electronegativity to help us understand:

• non-polar covalent bonds
• polar covalent bonds
• ionic bonding

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Electronegativity Difference

We can calculate the difference in electronegativity between two bonded atoms.

This is often written:

ΔEN = |EN₁ − EN₂|

The vertical bars mean that we take the positive difference.

For example:

Hydrogen ≈ 2.20

Chlorine ≈ 3.16

Therefore:

ΔEN = |3.16 − 2.20|

ΔEN = 0.96

This significant difference means the H–Cl bond is polar covalent.

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Non-Polar Covalent Bonds

When two atoms have the same or very similar electronegativities, the bonding electrons are shared relatively equally.

This produces a non-polar covalent bond.

For example:

H–H

Both atoms are hydrogen.

They have exactly the same electronegativity.

Therefore:

ΔEN = 0

Neither atom pulls the bonding electrons more strongly than the other.

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Polar Covalent Bonds

When two non-metal atoms have noticeably different electronegativities, the bonding electrons are shared unequally.

This produces a polar covalent bond.

Consider:

H–Cl

Chlorine is more electronegative.

Therefore, the bonding electrons spend more time closer to chlorine.

This creates partial charges:

Hδ⁺ — Clδ⁻

The symbol δ means "partial."

Hydrogen is slightly positive.

Chlorine is slightly negative.

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Bond Dipoles

A difference in electronegativity produces a bond dipole.

The more electronegative atom attracts the bonding electrons more strongly and develops a partial negative charge.

The less electronegative atom develops a partial positive charge.

For example:

Hδ⁺ — Clδ⁻

The unequal electron distribution is called bond polarity.

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Ionic Bonding

When the electronegativity difference between two elements is large, the bonding can have strong ionic character.

A common example is:

NaCl

Sodium has relatively low electronegativity.

Chlorine has much higher electronegativity.

Sodium loses an electron:

Na → Na⁺ + e⁻

Chlorine gains an electron:

Cl + e⁻ → Cl⁻

The oppositely charged ions then attract each other.

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Electronegativity Difference and Bond Type

A common introductory guideline using Pauling electronegativity differences is:

ΔEN about 0–0.4

Mostly non-polar covalent

ΔEN about 0.5–1.7

Mostly polar covalent

ΔEN greater than about 1.7

Often described as having strong ionic character

These values are useful guidelines, not absolute boundaries.

Chemical bonding exists along a continuum rather than falling into perfectly separated categories.

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Bonding Is a Continuum

It is useful to imagine bonding as a spectrum:

equal sharing → unequal sharing → strong ionic character

As the electronegativity difference increases:

electron sharing becomes increasingly unequal

and:

ionic character generally increases.

This is more accurate than imagining three completely separate types of bonds with sharp boundaries.

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Example: H₂

Hydrogen bonded to hydrogen:

H–H

Both atoms have the same electronegativity.

Therefore:

ΔEN = 0

The bond is:

non-polar covalent

The electrons are shared equally.

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Example: HCl

Approximate electronegativities:

H = 2.20

Cl = 3.16

Difference:

ΔEN = 0.96

Therefore, the bond is:

polar covalent

The bonding electrons are attracted more strongly toward chlorine.

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Example: NaCl

Approximate electronegativities:

Na = 0.93

Cl = 3.16

Difference:

ΔEN = 2.23

This large difference is consistent with the strong ionic character of sodium chloride.

Sodium forms:

Na⁺

and chlorine forms:

Cl⁻

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Example: C–H

Approximate electronegativities:

C = 2.55

H = 2.20

Difference:

ΔEN = 0.35

The electronegativity difference is relatively small.

The C–H bond is therefore commonly treated as essentially non-polar in introductory chemistry.

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Electronegativity and Water

Water provides an important example of electronegativity affecting molecular properties.

Oxygen is more electronegative than hydrogen.

Therefore, each O–H bond is polar.

Oxygen attracts the bonding electrons more strongly.

This gives oxygen a partial negative charge:

Oδ⁻

and the hydrogen atoms partial positive charges:

Hδ⁺

Because of the bent shape of the water molecule, these bond dipoles do not cancel.

Water is therefore a polar molecule.

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Bond Polarity Is Not the Same as Molecular Polarity

This distinction is important.

Electronegativity differences tell us about the polarity of individual bonds.

But the overall polarity of a molecule also depends on its shape.

For example, carbon dioxide contains polar C=O bonds.

However, CO₂ is linear:

O=C=O

The two bond dipoles point in opposite directions and cancel.

Therefore, CO₂ is overall non-polar, even though its individual C=O bonds are polar.

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Electronegativity and Metallic Character

Electronegativity and metallic character show broadly opposite trends.

Metallic character generally increases toward the:

bottom-left

Electronegativity generally increases toward the:

top-right

Strongly metallic elements tend to lose electrons relatively easily.

Strongly electronegative non-metals tend to attract bonding electrons strongly.

This helps explain why metals and non-metals often react together to form ionic compounds.

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Electronegativity and Reactivity

Electronegativity also helps explain some reactivity trends.

For Group 17:

F > Cl > Br > I

in electronegativity.

The same group also shows decreasing halogen reactivity down the group.

Fluorine's small atomic radius and strong attraction for electrons contribute to its very high reactivity.

However, electronegativity and reactivity are not the same property, and reactivity depends on the type of reaction being considered.

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Predicting an Unfamiliar Element

Suppose element X is located immediately to the left of fluorine in the same period.

Even without knowing its exact electronegativity, we can predict that X should have:

• relatively high electronegativity
• lower electronegativity than fluorine
• strong non-metallic character
• a strong tendency to attract bonding electrons

This prediction comes directly from periodic position.

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Worked Example: Predicting Bond Polarity

Consider a bond between carbon and oxygen.

Oxygen is farther right in the periodic table.

Therefore:

oxygen is more electronegative than carbon.

We predict that the bonding electrons will be pulled toward oxygen.

Therefore:

Cδ⁺ — Oδ⁻

The C–O bond is polar.

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Worked Example: Comparing Three Elements

Arrange:

K, Br, F

from lowest to highest electronegativity.

Potassium is far toward the lower-left.

Bromine is toward the upper-right.

Fluorine is above bromine and is the most electronegative element.

Therefore:

K < Br < F

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Worked Example: Across Period 3

Arrange:

Na, Al, S, Cl

from lowest to highest electronegativity.

All are in Period 3.

Moving left to right:

Na → Al → S → Cl

Electronegativity generally increases.

Therefore:

Na < Al < S < Cl

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Using Periodic Position Instead of Memorization

You do not need to memorize every electronegativity value.

Instead, remember the broad trend:

Electronegativity increases →

and:

Electronegativity increases ↑

toward fluorine.

Then compare the locations of the elements.

For most introductory problems, periodic position provides enough information to predict relative electronegativity.

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A Strategy for Predicting Bond Type

When given two elements:

Step 1: Locate both elements on the periodic table.

Step 2: Determine which is more electronegative.

Step 3: If numerical values are provided, calculate:

ΔEN = |EN₁ − EN₂|

Step 4: Consider the difference.

Very small difference:

mostly non-polar covalent

Moderate difference:

polar covalent

Large difference:

strong ionic character

Step 5: Remember that these categories form a continuum rather than having perfectly sharp boundaries.

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A Strategy for Comparing Electronegativity

When comparing two elements:

Step 1: Look at their positions.

Step 2: Remember:

right → higher electronegativity

up → higher electronegativity

Step 3: Consider atomic radius.

Smaller atoms generally attract bonding electrons more strongly.

Step 4: Consider shielding.

More shielding generally reduces attraction for bonding electrons.

Step 5: Use these ideas to justify your comparison.

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Common Mistakes

Mistake 1: Saying electronegativity measures electron loss

Electronegativity measures an atom's ability to attract bonding electrons.

Ionization energy is related to the energy required to remove an electron.

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Mistake 2: Saying electronegativity decreases across a period

The general trend is:

electronegativity increases from left to right.

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Mistake 3: Saying electronegativity increases down a group

Generally:

electronegativity decreases down a group.

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Mistake 4: Forgetting shielding

Atoms lower in a group have more occupied electron shells.

This increases shielding and reduces the effective attraction for bonding electrons.

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Mistake 5: Assuming all covalent bonds share electrons equally

Many covalent bonds involve unequal electron sharing because the bonded atoms have different electronegativities.

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Mistake 6: Assuming a polar bond automatically makes a polar molecule

Molecular shape must also be considered.

Polar bonds can sometimes have dipoles that cancel.

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Mistake 7: Treating bond-type boundaries as absolute

Electronegativity difference is a useful guide, but bonding exists on a continuum between covalent and ionic character.

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Real-World Importance of Electronegativity

Electronegativity helps explain many properties of substances.

It helps chemists understand and predict:

• bond polarity
• molecular polarity
• solubility
• intermolecular attractions
• chemical reactivity
• molecular structure
• properties of materials
• behaviour of biological molecules

For example, the polarity of water strongly influences its ability to dissolve many ionic and polar substances.

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Did You Know?

The electronegativity scale most commonly introduced in chemistry is associated with chemist Linus Pauling.

Pauling developed a way of comparing how strongly atoms attract electrons in chemical bonds.

Unlike measurements such as mass or length, electronegativity is a relative scale used to compare atoms.

Fluorine is assigned the highest value on the commonly used Pauling scale.

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Key Terms

Electronegativity: An atom's tendency to attract bonding electrons.

Bonding electrons: Electrons involved in forming a chemical bond.

Period: A horizontal row of the periodic table.

Group: A vertical column of the periodic table.

Atomic radius: A measure of atomic size.

Shielding: Reduction in effective nuclear attraction caused by inner electrons.

Effective nuclear attraction: The net attraction experienced by electrons due to the nucleus after considering shielding.

Non-polar covalent bond: A covalent bond in which electrons are shared approximately equally.

Polar covalent bond: A covalent bond in which electrons are shared unequally.

Bond polarity: Unequal distribution of electron density across a bond.

Dipole: A separation of partial positive and negative charge.

Ionic character: The degree to which bonding resembles attraction between oppositely charged ions.

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Key Trends

Across a period:

electronegativity generally increases →

because effective nuclear attraction increases and atomic radius generally decreases.

Down a group:

electronegativity generally decreases ↓

because atomic radius and shielding increase.

Overall:

Electronegativity increases toward the upper-right of the periodic table.

Fluorine is the most electronegative element.

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Key Bonding Relationship

As the difference in electronegativity increases:

similar electronegativity

↓

approximately equal electron sharing

↓

non-polar covalent character

Then:

moderate electronegativity difference

↓

unequal electron sharing

↓

polar covalent character

And with a very large difference:

strong electron attraction difference

↓

greater ionic character

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Key Takeaways

• Electronegativity describes an atom's ability to attract bonding electrons.
• Electronegativity generally increases from left to right across a period.
• Electronegativity generally decreases down a group.
• The highest electronegativities are found toward the upper-right of the periodic table.
• Fluorine is the most electronegative element.
• Across a period, increasing effective nuclear attraction and decreasing atomic radius generally increase electronegativity.
• Down a group, increasing atomic radius and shielding generally decrease electronegativity.
• Periodic position can be used to compare elements without memorizing every electronegativity value.
• A difference in electronegativity can cause bonding electrons to be shared unequally.
• Similar electronegativities generally produce greater non-polar covalent character.
• Moderate electronegativity differences can produce polar covalent bonds.
• Large electronegativity differences are often associated with strong ionic character.
• Bonding exists on a continuum rather than fitting into perfectly separated categories.
• The more electronegative atom in a polar bond develops a partial negative charge, while the less electronegative atom develops a partial positive charge.
• Bond polarity does not automatically determine molecular polarity; molecular shape must also be considered.
• Electronegativity connects atomic structure, periodic trends, bonding, polarity, and chemical properties, making it one of the most useful ideas for understanding patterns in chemistry.