Polar and Non-polar Covalent Bonds — Revision Notes

⚡ 30-Second Revision

Electronegativity ($Delta EN$) — Determines bond polarity.

- $Delta EN approx 0$: Non-polar covalent (e.g., $H_2, O_2$). - $0.4 le Delta EN < 1.7$: Polar covalent (e.g., $HCl, H_2O$ bonds).

Dipole Moment ($mu$) — Quantitative measure of polarity. Vector quantity. $\mu = q \times r$.
Molecular Polarity — Determined by bond polarity + molecular geometry.

- Non-polar molecule: Net $\mu = 0$. Either all bonds are non-polar, or polar bonds are arranged symmetrically and cancel (e.g., $CO_2$ (linear), $CCl_4$ (tetrahedral), $BF_3$ (trigonal planar), $XeF_4$ (square planar)). - Polar molecule: Net $\mu \ne 0$. Polar bonds with asymmetrical geometry (e.g., $H_2O$ (bent), $NH_3$ (pyramidal), $CHCl_3$ (asymmetrical tetrahedral)).

Properties — Polar molecules have stronger IMFs (dipole-dipole, H-bonding), higher BP/MP, soluble in polar solvents ('like dissolves like').

2-Minute Revision

Polar and non-polar covalent bonds are distinguished by the equality of electron sharing, which is dictated by the **electronegativity difference ($Delta EN$)** between bonded atoms. A non-polar bond occurs when $Delta EN$ is zero or very small (e.

g., $H_2$, C-H), leading to equal sharing. A polar bond forms when there's a significant $Delta EN$ (e.g., O-H, H-Cl), causing unequal sharing and the development of partial charges ($delta^+, delta^-$).

This charge separation creates a bond dipole moment. The overall molecular polarity is determined by the vector sum of all bond dipoles and the molecular geometry (VSEPR theory). Symmetrical molecules like $CO_2$ (linear) or $CCl_4$ (tetrahedral) can have polar bonds but be non-polar overall because their bond dipoles cancel out.

Asymmetrical molecules like $H_2O$ (bent) or $NH_3$ (pyramidal) have a net dipole moment and are thus polar. Molecular polarity profoundly affects physical properties like solubility ('like dissolves like') and boiling/melting points, with polar molecules generally having stronger intermolecular forces.

5-Minute Revision

Understanding polar and non-polar covalent bonds is crucial for NEET. Start by recalling that a covalent bond involves electron sharing. The key concept is electronegativity (EN), an atom's ability to attract shared electrons. The **difference in electronegativity ($Delta EN$)** between two bonded atoms determines bond polarity.

1
Non-polar Covalent Bonds — Occur when $Delta EN$ is zero or very small (typically $<0.4$). Electrons are shared equally. Examples: $H_2, O_2, N_2, Cl_2$. C-H bonds are often considered non-polar.
2
Polar Covalent Bonds — Occur when there's a significant $Delta EN$ (typically $0.4 le Delta EN < 1.7$). Electrons are shared unequally, creating partial positive ($delta^+$) and partial negative ($delta^-$) charges. Examples: H-Cl, O-H, N-H. This creates a bond dipole.

**Dipole Moment ($mu$)** quantifies polarity ($mu = q \times r$). It's a vector pointing from $delta^+$ to $delta^-$. The unit is Debye (D).

Molecular Polarity is the overall polarity of a molecule, determined by the vector sum of all individual bond dipoles and the molecular geometry (predicted by VSEPR theory).

Non-polar Molecules (Net $mu = 0$)

* Contain only non-polar bonds (e.g., $CH_4$, if C-H is considered non-polar). * Contain polar bonds, but the molecule's symmetrical geometry causes the bond dipoles to cancel out. Examples: $CO_2$ (linear), $CCl_4$ (tetrahedral), $BF_3$ (trigonal planar), $XeF_4$ (square planar).

**Polar Molecules (Net $mu

e 0$)**: * Contain polar bonds and have an asymmetrical geometry, so bond dipoles do not cancel. Examples:$H_2O$(bent),$NH_3$(trigonal pyramidal),$CHCl_3$(asymmetrical tetrahedral),$SO_2$ (bent).

Impact on Physical Properties: Polarity dictates intermolecular forces (IMFs). Polar molecules exhibit stronger IMFs (dipole-dipole interactions, hydrogen bonding) in addition to London dispersion forces (LDF).

Non-polar molecules primarily rely on weaker LDF. Stronger IMFs lead to higher boiling points, melting points, and surface tension. The 'like dissolves like' rule applies: polar solvents dissolve polar solutes, and non-polar solvents dissolve non-polar solutes.

For instance, water (polar) dissolves sugar (polar) but not oil (non-polar). Practice identifying geometries and vector sums for various molecules to master this concept.

Prelims Revision Notes

Polar and Non-polar Covalent Bonds: NEET Revision Notes

1. Covalent Bond Basics:

Formed by sharing of electrons between atoms.
Electron sharing can be equal or unequal.

2. Electronegativity (EN):

Atom's ability to attract shared electrons in a bond.
Pauling scale: F (3.98) is highest, Cs (0.79) is lowest.
Trends: Increases across a period, decreases down a group.

3. Bond Polarity (Based on $Delta EN$):

Non-polar Covalent Bond:

* $Delta EN approx 0$ (typically $< 0.4$). * Equal sharing of electrons. * No partial charges ($delta^+, delta^-$). * Examples: $H_2, O_2, N_2, Cl_2$, C-C, C-H (often considered non-polar).

Polar Covalent Bond:

* Significant $Delta EN$ (typically $0.4 le Delta EN < 1.7$). * Unequal sharing of electrons. * More electronegative atom gets $delta^-$, less electronegative atom gets $delta^+$. * Examples: H-Cl, O-H, N-H, C-O.

Ionic Bond:

* Very large $Delta EN$ (typically $ge 1.7$). * Complete transfer of electrons.

4. Dipole Moment ($mu$):

Quantitative measure of polarity.
Formula: $\mu = q \times r$ (charge magnitude $imes$ distance).
Unit: Debye (D).
Vector quantity: Direction from $delta^+$ to $delta^-$ (less EN to more EN atom).
Higher $mu$ indicates greater polarity.

5. Molecular Polarity (Crucial Distinction):

Determined by BOTH bond polarity AND molecular geometry (VSEPR theory).
It's the vector sum of all individual bond dipoles in the molecule.

* **Non-polar Molecules (Net $mu = 0$):** * All bonds are non-polar (e.g., $CH_4$ if C-H is non-polar). * OR, molecule contains polar bonds, but its symmetrical geometry causes the bond dipoles to cancel out. * Key Symmetrical Geometries (leading to non-polar molecules if terminal atoms are identical): * Linear ($CO_2$) * Trigonal Planar ($BF_3$) * Tetrahedral ($CCl_4, SiH_4$) * Trigonal Bipyramidal ($PCl_5$) * Octahedral ($SF_6$) * Square Planar ($XeF_4$)

* **Polar Molecules (Net $mu e 0$):** * Molecule contains polar bonds and has an asymmetrical geometry, so bond dipoles do not cancel. * Key Asymmetrical Geometries (leading to polar molecules): * Bent ($H_2O, SO_2, H_2S$) * Trigonal Pyramidal ($NH_3, PCl_3$) * Asymmetrical Tetrahedral ($CHCl_3, CH_2Cl_2$) * See-saw ($SF_4$) * T-shaped ($ClF_3$)

6. Impact on Physical Properties:

Intermolecular Forces (IMFs):

* Non-polar molecules: Only London Dispersion Forces (LDF). * Polar molecules: LDF + Dipole-Dipole interactions (and Hydrogen Bonding if H is bonded to F, O, or N).

Boiling/Melting Points: — Stronger IMFs $ightarrow$ Higher BP/MP. Polar molecules generally have higher BP/MP than non-polar molecules of comparable size.
Solubility: — 'Like dissolves like'. Polar solutes dissolve in polar solvents (e.g., sugar in water). Non-polar solutes dissolve in non-polar solvents (e.g., oil in hexane).

7. Common Traps:

Confusing bond polarity with molecular polarity.
Ignoring lone pairs when determining VSEPR geometry.
Not considering the vector nature of dipole moments.

Vyyuha Quick Recall

To remember when a molecule is POLAR or NON-POLAR, think of 'S.A.N.D.':

Symmetry: If the molecule is Symmetrical, it's likely Non-polar (dipoles cancel). Asymmetry: If the molecule is Asymmetrical, it's likely Polar (dipoles don't cancel). No Lone Pairs + Identical Atoms: If the central atom has No Lone Pairs and is bonded to Identical Atoms, it's usually symmetrical and Non-polar (e.

g., $CO_2, CCl_4, BF_3$). Different Atoms / Lone Pairs: If the central atom has Different Atoms bonded to it OR has Lone Pairs, it's usually asymmetrical and Polar (e.g., $CHCl_3, H_2O, NH_3$).

(Remember, this is a general guide; always confirm with VSEPR and vector analysis!)