Short Answer

Periodic trends in electronegativity describe how the ability of an atom to attract shared electrons changes across periods and down groups in the periodic table. Electronegativity increases from left to right across a period because nuclear charge increases, pulling electrons more strongly. It decreases from top to bottom in a group because the atomic size increases, and the nucleus has less control over distant electrons.

These trends help predict the polarity of bonds, chemical reactivity, and the nature of compounds. Understanding electronegativity patterns is essential in studying bonding and periodic behaviour.

Detailed Explanation :

Periodic Trends in Electronegativity

Electronegativity is a fundamental concept in chemistry that explains how strongly an atom attracts electrons when it forms a chemical bond. When arranged in the periodic table, elements show a clear and predictable pattern in their electronegativity values. These patterns, known as periodic trends, help chemists understand why some elements attract electrons strongly while others do not.

Electronegativity plays a major role in determining the type of bond formed—ionic, polar covalent, or non-polar covalent. It also influences chemical reactions, molecular shapes, and properties of compounds. By studying periodic trends in electronegativity, we can predict the behaviour of elements in various chemical situations.

Electronegativity Across Periods and Groups

Electronegativity Across a Period

When moving from left to right across a period, electronegativity increases.

Reasons:

1. Increasing nuclear charge
   As the number of protons in the nucleus increases, the attraction for shared electrons also increases.
2. Decreasing atomic size
   Smaller atoms hold electrons more tightly.
   As atomic radius decreases, the nucleus pulls electrons more strongly.
3. Greater effective nuclear charge
   Electrons experience a stronger pull due to less shielding within the same shell.

Example Trend (Period 2):

Li < Be < B < C < N < O < F
Fluorine is the most electronegative element in the periodic table.

Electronegativity Down a Group

When moving from top to bottom in a group, electronegativity decreases.

Reasons:

1. Increasing atomic size
   The outer electrons are farther from the nucleus.
   This reduces the attractive force on shared electrons.
2. Increased shielding effect
   Inner electrons block the nuclear charge, weakening the pull on electrons.
3. Lower effective nuclear charge on valence electrons
   The nucleus has less influence on distant electrons, reducing electronegativity.

Example Trend (Group 17):

F > Cl > Br > I
Fluorine has the highest, iodine the lowest electronegativity in this group.

Why These Trends Occur

1. Role of Nuclear Charge

More protons mean stronger attraction for electrons.
This increases electronegativity across a period.

2. Role of Atomic Radius

Larger atoms have electrons farther from the nucleus, lowering their electronegativity.

3. Shielding Effect

Inner electrons reduce the pull of the nucleus on outer electrons.
Higher shielding decreases electronegativity.

4. Effective Nuclear Charge

The effective force felt by outer electrons affects electronegativity directly.

Special Points in Electronegativity Trends

1. Noble Gases

Most noble gases have zero electronegativity because they do not normally form bonds.

2. Transition Elements

Transition metals show irregular trends because d-electrons affect shielding and bonding.

3. Most Electronegative Element

Fluorine (F) is the most electronegative element.
It strongly attracts electrons in any bond.

Effects of Electronegativity on Chemical Properties

Electronegativity trends help explain several chemical behaviours:

1. Bond Type

• Large differences in electronegativity → ionic bond
• Moderate differences → polar covalent bond
• Small or no difference → non-polar covalent bond

2. Bond Polarity

Higher electronegativity difference creates stronger bond polarity.

Example:
H–F bond is very polar because fluorine attracts electrons strongly.

3. Reactivity

• Non-metals with high electronegativity (like halogens) are very reactive because they readily gain electrons.
• Metals with low electronegativity lose electrons easily, increasing their reactivity.

4. Acid–Base Behaviour

Electronegativity influences acidic and basic strength of compounds.

5. Molecular Properties

Polarity, solubility, boiling point, and melting point depend on electronegativity differences.

Real Examples Showing Electronegativity Trends

Water (H₂O)

Oxygen’s high electronegativity makes water a polar molecule.

Hydrogen chloride (HCl)

Chlorine pulls electrons strongly, making the bond polar.

Sodium chloride (NaCl)

Large electronegativity difference results in an ionic bond.

Oxygen vs Sulphur

Oxygen is more electronegative than sulphur because it is smaller and higher in the group.

Importance of Studying Electronegativity Trends

Understanding these trends helps in:

• Predicting how elements will react
• Understanding the polarity of molecules
• Determining bond types
• Explaining physical properties of compounds
• Analysing organic and inorganic reactions
• Studying biological molecules

Electronegativity is a key tool for predicting and interpreting chemical behaviour.

Conclusion

Periodic trends in electronegativity describe how the ability of atoms to attract shared electrons changes across the periodic table. Electronegativity increases from left to right across a period because of higher nuclear charge and smaller atomic size, and decreases from top to bottom in a group due to increased atomic size and shielding. These predictable trends help explain bond formation, molecular polarity, and chemical reactivity, making electronegativity an essential concept in chemistry.