5.3.3 Types of Covalent Bond on the Basis of Polarity

This topic explains how covalent bonds are classified as non-polar or polar, depending on whether the two atoms sharing electrons are the same or different.

Definition

Non-Polar Covalent Bond

A covalent bond may form between two similar atoms, such as in H₂, N₂, O₂, Cl₂, etc.

When two identical atoms share electron pairs, both atoms exert the same force on the shared electron pairs. Such a covalent bond is called a non-polar covalent bond.

Simple explanation: When the two atoms are the same, they pull the shared electrons equally — no side becomes charged.

Polar Covalent Bond

When two different atoms share an electron pair, both atoms exert different forces on the shared electron pair.

The more electronegative atom pulls the shared electron pairs towards itself with a greater force than the other atom. Thus, the more electronegative atom attracts some of the electron density towards itself. This makes it partially negatively charged and the other atom partially positively charged.

Such a covalent bond is called a polar covalent bond.

Simple explanation: When the two atoms are different, the more electronegative atom hogs the shared electrons, creating partial charges.

Main Idea

The type of covalent bond depends on electronegativity difference between the two bonded atoms:

• Same atoms → equal pull → non-polar covalent bond
• Different atoms → unequal pull → polar covalent bond

Simple explanation: Identical twins share equally; different siblings share unequally.

Core Concept

Detailed Explanation

Non-Polar Covalent Bond

Formation: Between identical atoms. Electron sharing: Equal — both atoms exert the same force on the shared pair. Result: No partial charges.

• H–H (in H₂)
• **O=O** (in O₂)
• N≡N (in N₂)
• Cl–Cl (in Cl₂)

Simple explanation: Because the atoms are identical, the shared electrons stay exactly in the middle.

Polar Covalent Bond

Formation: Between different atoms. Electron sharing: Unequal — the more electronegative atom pulls the shared pair toward itself. Result: Partial positive charge (δ+) on one atom, partial negative charge (δ−) on the other.

• H–Cl (in HCl) → Cl is more electronegative, so it gets δ−; H gets δ+
• H₂O (O–H bonds) → O is more electronegative, so it gets δ−; H gets δ+
• NH₃ (N–H bonds) → N is more electronegative, so it gets δ−; H gets δ+

Simple explanation: The stronger atom grabs the shared electrons; the weaker one is left slightly positive.

Intermolecular Forces

The book mentions that the forces of attraction between molecules are called intermolecular forces.

• In polar molecules, dipole–dipole interactions occur between the δ+ end of one molecule and the δ− end of another.
• In non-polar molecules, only weak dispersion forces exist.

Intermediate Level

Why Unequal Sharing Happens

Electronegativity is the ability of an atom to attract electrons in a bond.

• Higher electronegativity → stronger pull on shared electrons.
• Lower electronegativity → weaker pull on shared electrons.

When the two atoms have different electronegativities, the shared electron pair is pulled toward the more electronegative atom, creating partial charges.

Partial Charges (δ+ and δ−)

The symbol δ (delta) means "partial" or "slightly."

• δ− = partially negative (the more electronegative atom)
• δ+ = partially positive (the less electronegative atom)

• Cl is more electronegative → δ−
• H is less electronegative → δ+

Electronegativity Difference and Bond Type

Simple explanation: The bigger the difference, the more polar the bond. If the difference is very large, the bond becomes ionic.

Examples of Polar and Non-Polar Bonds

Important Points

• - Non-polar covalent bond: formed between identical atoms; shared electrons are pulled equally; no partial charges.
• - Polar covalent bond: formed between different atoms; shared electrons are pulled unequally; partial charges (δ+ and δ−) appear.
• - More electronegative atom → δ− (partially negative).
• - Less electronegative atom → δ+ (partially positive).
• - Examples of non-polar bonds: H–H, O=O, N≡N, Cl–Cl.
• - Examples of polar bonds: H–Cl, H₂O, NH₃.
• - Intermolecular forces: forces of attraction between molecules (dipole–dipole for polar molecules).
• - Electronegativity difference determines bond polarity.

Remember

• - Non-polar = same atoms = equal sharing = no charges.
• - Polar = different atoms = unequal sharing = δ+ and δ−.
• - More electronegative atom → δ−.
• - Less electronegative atom → δ+.
• - Examples:
• - Non-polar: H–H, O=O, N≡N, Cl–Cl
• - Polar: H–Cl, H₂O, NH₃
• - Electronegativity difference determines bond polarity.
• - Intermolecular forces are forces between molecules.

Common Mistakes

• 1. Thinking all covalent bonds are non-polar. Bonds between different atoms are polar.
• 2. Thinking polar bonds have full charges. They have partial charges (δ+ and δ−), not full charges.
• 3. Confusing polar covalent with ionic. Polar covalent = unequal sharing; ionic = complete transfer.
• 4. Thinking the bigger atom gets the δ− charge. It's the more electronegative atom that gets δ−.
• 5. Forgetting that H–O and N–H bonds are polar. They are polar because O and N are more electronegative than H.
• 6. Thinking non-polar bonds have zero electronegativity. They have equal electronegativity (difference = 0).

Quick Revision

• Non-polar covalent bond: formed between two identical atoms (e.g., H₂, O₂, N₂, Cl₂); both atoms exert the same force on shared electron pairs.
• Polar covalent bond: formed between two different atoms (e.g., HCl, H₂O, NH₃); the more electronegative atom pulls the shared electron pair toward itself, becoming partially negative (δ−), leaving the other partially positive (δ+).
• Intermolecular forces: forces of attraction between molecules.
• Electronegativity difference determines whether a bond is non-polar, polar, or ionic.

60-Second Revision

• Non-polar = same atoms = equal sharing.
• Polar = different atoms = unequal sharing.
• More electronegative atom → δ−.
• Less electronegative atom → δ+.
• Examples: H–H (non-polar); H–Cl (polar).
• Warning: polar bonds have partial charges (δ+, δ−), not full charges.