3.1.2 Trend of Electronegativity in the Periodic Table

This topic explains how electronegativity changes as we move across a period and down a group in the periodic table. It also provides examples to illustrate these trends.

Definition

Electronegativity: The power of a covalently bonded atom to attract the shared pair of electrons towards itself in a molecule.

Main Idea

• Down a group (top to bottom): Electronegativity decreases.
• Across a period (left to right): Electronegativity increases.

These trends can be explained by changes in nuclear charge, atomic size, and shielding effect.

Core Concept

• Nuclear charge — more protons pull electrons more strongly.
• Atomic size — smaller atoms hold electrons more tightly.
• Shielding effect — inner shells reduce the pull on outer electrons.

Detailed Explanation

Trend Down a Group (Top to Bottom)

• Electronegativity decreases from top to bottom in groups.
• There is an increase in the atomic number as we move down the group in the modern periodic table.
• The nuclear charge also increases, but the effect of the increase in nuclear charge is overcome by the addition of an extra shell (increase in shielding effect).
• The attraction of shared electrons with the nucleus decreases.
• Hence, the value of electronegativity decreases.

Example: - In the halogen group, as we move down the group from fluorine to astatine, the electronegativity value decreases.

• Fluorine is at the top of the group with the smallest size and highest nuclear charge relative to its size.
• As we go down, the atomic size increases, and the shielding effect increases.
• The outer electrons are farther from the nucleus, so the attraction for shared electrons decreases.
• Therefore, electronegativity decreases down the group.

Trend Across a Period (Left to Right)

• Electronegativity increases from left to right across a period.
• When we move from left to right in periods, the effective nuclear charge increases.
• Shielding effect remains constant.
• Hence, the size of atom decreases.
• The power of atoms to attract the shared pair of electrons increases.

Example: - In the third period, from sodium to chlorine, the electronegativity value increases.

• Sodium has a low electronegativity because it has a large atomic size and low effective nuclear charge.
• As we move across the period, the nuclear charge increases, and the atomic size decreases.
• The outer electrons are held more tightly by the nucleus.
• Therefore, electronegativity increases across the period.

Summary of Trends

Why This Matters

• Predicting bond types: Electronegativity differences determine whether a bond is ionic, polar covalent, or nonpolar covalent.
• Understanding reactivity: Elements with high electronegativity tend to gain electrons (non-metals); elements with low electronegativity tend to lose electrons (metals).
• Explaining periodic trends: Electronegativity is one of the key periodic properties that helps explain the behavior of elements.

Intermediate Level

Trend Down a Group

• Electronegativity decreases from top to bottom.
• Reason: Atomic size increases; shielding effect increases.
• The nucleus's attraction for shared electrons decreases.
• Example: F > Cl > Br > I > At

Trend Across a Period

• Electronegativity increases from left to right.
• Reason: Nuclear charge increases; atomic size decreases.
• The nucleus's attraction for shared electrons increases.
• Example: Na < Mg < Al < Si < P < S < Cl

Comparison Table

Important Points

• 1. Electronegativity decreases down a group.
• 2. Electronegativity increases across a period.
• 3. Down a group: Atomic size increases; shielding effect increases.
• 4. Across a period: Nuclear charge increases; atomic size decreases.
• 5. Example: Halogens → F > Cl > Br > I > At.
• 6. Example: Period 3 → Na < Mg < Al < Si < P < S < Cl.

Remember

• - Down a group: Electronegativity decreases
• - Across a period: Electronegativity increases
• - Reason (down): Size increases, shielding increases
• - Reason (across): Nuclear charge increases, size decreases
• - Example: F > Cl > Br > I > At
• - Example: Na < Mg < Al < Si < P < S < Cl

Common Mistakes

• 1. Thinking electronegativity increases down a group: It decreases.
• 2. Forgetting the shielding effect: It reduces the nuclear pull on outer electrons.
• 3. Confusing electronegativity with electron affinity: Electronegativity is about attracting shared electrons; electron affinity is about gaining an electron.
• 4. Mixing up the trends: Across a period = increases; Down a group = decreases.

Quick Revision

• Electronegativity decreases down a group.
• Electronegativity increases across a period.
• Down a group: Atomic size increases; shielding effect increases.
• Across a period: Nuclear charge increases; atomic size decreases.
• Example: Halogens → F > Cl > Br > I > At.
• Example: Period 3 → Na < Mg < Al < Si < P < S < Cl.

60-Second Revision

• Down a group: Electronegativity decreases
• Across a period: Electronegativity increases
• Reason (down): Size increases, shielding increases
• Reason (across): Nuclear charge increases, size decreases
• Example: F > Cl > Br > I > At
• Example: Na < Mg < Al < Si < P < S < Cl

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