Short Answer

A molecule becomes polar when it has polar bonds and its shape does not allow the bond dipoles to cancel out. This means the molecule must have an uneven distribution of electrons, creating a positive end and a negative end. Electronegativity difference between atoms produces polar bonds.

However, for a molecule to be polar overall, its molecular geometry must be asymmetrical. Lone pairs, different surrounding atoms, or bent or pyramidal shapes often cause this asymmetry, resulting in a net dipole moment and making the molecule polar.

Detailed Explanation :

Conditions That Make a Molecule Polar

A molecule is considered polar when it has a net dipole moment, meaning there is an overall separation of positive and negative charges. Polarity arises from both the electronegativity difference between atoms and the shape of the molecule. Even if bonds are polar, the molecule may still be non-polar if the shape allows dipole moments to cancel. Therefore, polarity depends on two major conditions working together: polar bonds and asymmetrical geometry.

Understanding polarity is important because it determines solubility, reactivity, intermolecular forces, melting and boiling points, and many physical and chemical behaviours of molecules.

1. Condition 1: Presence of Polar Bonds

The first requirement for molecular polarity is that the molecule must contain at least one polar covalent bond. A polar covalent bond forms when the two atoms involved have different electronegativities.

• The atom with higher electronegativity pulls the electron pair closer and becomes slightly negative (δ–).
• The atom with lower electronegativity becomes slightly positive (δ+).

This creates a bond dipole.

Examples of polar bonds:

• H–Cl
• O–H
• N–H
• C–O

Without polar bonds, there can be no dipole moment, and therefore no molecular polarity.

2. Condition 2: Asymmetrical Molecular Shape

Even if a molecule has polar bonds, it will only be polar if the shape is asymmetrical. Asymmetry prevents the bond dipoles from canceling out.

Molecules become asymmetrical when:

• They contain lone pairs on the central atom
• They have different atoms attached to the central atom
• They have shapes like bent, trigonal pyramidal, seesaw, or T-shaped

These arrangements make the dipoles point in different directions and create a net dipole.

Examples of polar molecules due to asymmetry:

• H₂O (bent shape)
• NH₃ (trigonal pyramidal)
• SO₂ (bent due to a lone pair)
• CH₃Cl (different surrounding atoms)

These molecules all have nonzero dipole moments.

3. Condition 3: Dipole Moments Do Not Cancel Out

Dipole moments behave like vectors—they have direction and magnitude.
For a molecule to be polar:

• The dipole vectors must not cancel each other.

Dipole cancellation happens in symmetrical molecules, such as:

• CO₂ (linear)
• BF₃ (trigonal planar)
• CCl₄ (tetrahedral)
• XeF₂ (linear with lone pairs symmetric)

Even though these molecules may contain polar bonds, the geometry is perfectly symmetrical, so the dipoles cancel, resulting in non-polar molecules.

In contrast, polar molecules have dipole vectors that add up to give a net dipole moment because the geometry is not symmetrical.

4. Effect of Lone Pairs on Polarity

Lone pairs almost always increase polarity by distorting molecular shape.

Examples:

• In water, two lone pairs force the molecule into a bent shape → polar
• In ammonia, one lone pair creates asymmetry → polar
• In SO₂, lone pair creates a bent shape → polar

However, if lone pairs are arranged symmetrically (as in XeF₂ or square planar XeF₄), the molecule may still be non-polar.

Thus, lone pairs contribute to polarity but only if they cause asymmetry.

5. Condition 4: Different Surrounding Atoms

If a central atom is bonded to different types of atoms, symmetry is broken, causing polarity.

Example:

• CH₃Cl is polar because C–Cl bond is highly polarized, and the geometry is not symmetrical.
• CH₂Cl₂ is polar because the Cl atoms are arranged unevenly.

In contrast:

• CCl₄ is non-polar because all four substituents are identical, keeping symmetry intact.

Thus, having different surrounding atoms increases polarity.

6. Role of Electronegativity Difference

Greater electronegativity difference creates stronger dipoles, which increases molecular polarity.

Examples:

• H–F has a very large difference → strong dipole → HF is polar
• C–H has a small difference → weak dipole → hydrocarbons are mostly non-polar

However, dipole strength alone cannot make a molecule polar unless the shape also supports polarity.

7. Summary of Conditions for Polarity

To summarize, a molecule is polar if all these conditions are met:

1. It contains polar bonds (electronegativity difference).
2. The molecular shape is asymmetrical (dipoles do not cancel).
3. Dipole moments add together to produce a net dipole.
4. Lone pairs or different surrounding atoms create uneven electron distribution.

If any of these conditions are missing, the molecule is likely non-polar.

Conclusion

A molecule becomes polar when it has polar bonds and an asymmetrical shape that prevents dipole cancellation. Electronegativity differences create bond dipoles, while molecular geometry determines how these dipoles combine. Lone pairs, different surrounding atoms, and bent or pyramidal shapes often make molecules polar. Understanding these conditions helps predict physical properties like solubility, boiling point, and molecular interactions.