Explain with the help of a suitable example a polar covalent bond.

Exercises · 4.18

Q. Explain with the help of a suitable example a polar covalent bond.

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Concept understanding — Dipole Moment Applications

Dipole Moment Applications: From Intuition to Precision

Imagine you have a magnet. One end pulls, the other pushes. Now imagine a molecule that behaves like a tiny magnet — not because of iron, but because of how its electrons are distributed. That's the idea behind a dipole moment.

A dipole moment arises when there's a separation of charge inside a molecule. One end becomes slightly negative (δ−), the other slightly positive (δ+). This imbalance creates a tiny electric "arrow" pointing from positive to negative. The arrow has both a size (how much charge is separated) and a direction (which way the molecule is polar).

Note

The dipole moment is a vector quantity. Its magnitude is given by $\mu = q \times d$ , where $q$ is the magnitude of separated charge and $d$ is the distance between the charge centers. The unit is the Debye (D).

Now, why does this matter? Because this tiny electric arrow determines how a molecule behaves around other molecules, around electric fields, and even how it interacts with light. Let's see the key applications.

1. Predicting Molecular Shape (Symmetry Check)

This is the most common exam application. If a molecule has polar bonds (like C–Cl or O–H), does it have a net dipole moment? The answer depends on symmetry.

Consider carbon dioxide, CO₂. Each C=O bond is polar (oxygen pulls electrons). But the molecule is linear: O=C=O. The two dipole arrows point in opposite directions and cancel out. Net dipole moment = zero. The molecule is nonpolar.

Now consider water, H₂O. Each O–H bond is polar. But water is bent (104.5°). The two arrows do not cancel — they add up to a net dipole pointing upward through the oxygen. Net dipole moment = 1.85 D. Water is polar.

Important

Symmetry kills polarity. If a molecule has a center of symmetry or identical polar bonds arranged symmetrically, the net dipole moment is zero. This is how you distinguish between linear CO₂ (nonpolar) and bent SO₂ (polar, 1.63 D).

Exam tip: For molecules like CH₄ (tetrahedral, zero dipole) vs. CH₃Cl (tetrahedral but one C–Cl bond, dipole = 1.87 D), the key is whether the polar bonds are arranged so their vectors cancel.

2. Determining Bond Character (Ionic vs. Covalent)

The dipole moment tells you how "unequal" the sharing of electrons is in a bond. A pure covalent bond (like H–H) has zero dipole. A pure ionic bond (like Na⁺Cl⁻) would have a huge dipole — but in reality, ions are separate.

For a bond like H–Cl, the measured dipole moment is 1.08 D. If the bond were 100% ionic (one full electron transferred), the dipole would be much larger (about 6.1 D for the same bond length). The ratio gives you the percent ionic character:

\% \text{ ionic character} = \frac{\mu_{\text{observed}}}{\mu_{\text{calculated for 100\% ionic}}} \times 100

For HCl: $\frac{1.08}{6.1} \times 100 \approx 17.7\%$ . So the H–Cl bond is about 18% ionic, 82% covalent.

Tip

This is a standard numerical problem. Remember: $\mu_{\text{ionic}} = e \times d$ , where $e = 4.8 \times 10^{-10}$ esu (or $1.6 \times 10^{-19}$ C in SI). Convert bond length to cm or m accordingly.

3. Intermolecular Forces and Physical Properties

Polar molecules (with a nonzero dipole) experience dipole-dipole interactions — the positive end of one molecule attracts the negative end of another. This is stronger than the London dispersion forces in nonpolar molecules of similar size.

Consequences:

Boiling points: Polar molecules have higher boiling points than nonpolar ones of similar molar mass. Example: HCl (polar, bp −85°C) vs. F₂ (nonpolar, bp −188°C). Both have about 38 g/mol, but HCl's dipole adds extra attraction.
Solubility: "Like dissolves like." Polar solutes dissolve in polar solvents (water, ethanol). Nonpolar solutes dissolve in nonpolar solvents (hexane, CCl₄). The dipole moment explains why NaCl dissolves in water but not in oil.
Dielectric constant: Polar liquids have high dielectric constants (water = 80), meaning they can weaken the electric field between charges. This is why water is such a good solvent for ionic compounds.

Watch out

Don't confuse dipole moment with boiling point directly. A molecule can have a large dipole but low boiling point if it's very small (like HF, bp 19.5°C, dipole 1.91 D). Hydrogen bonding (a special case of dipole interaction) is even stronger.

4. Reactivity and Orientation in Electric Fields

In an external electric field, polar molecules align themselves with the field. This is the principle behind microwave heating — water molecules in food rotate to align with the alternating microwave field, generating heat through friction.

In organic chemistry, the dipole moment helps predict reaction sites. The negative end of a dipole (where electrons are concentrated) is where electrophiles attack. The positive end (electron-deficient) is where nucleophiles attack. …

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A polar covalent bond forms when two atoms of different electronegativities share electrons unequally, creating partial charges. The classic example is the H–Cl bond, where chlorine pulls the shared electrons closer, becoming partially negative while hydrogen becomes partially positive.

What makes a bond polar?

When two atoms form a covalent bond, they share electrons. But not all atoms pull on those shared electrons with equal strength. Electronegativity measures an atom's ability to attract bonding electrons toward itself. When two atoms with different electronegativities bond, the electron cloud shifts toward the more electronegative atom. This creates a separation of charge: one end becomes slightly negative (δ⁻), the other slightly positive (δ⁺). That's a polar covalent bond.

The key difference from a pure covalent bond (like H–H, where both atoms are identical) is this asymmetry. The electrons aren't shared equally, so the bond has a dipole moment — a vector pointing from the positive end to the negative end.

The hydrogen chloride (HCl) example

Let's walk through hydrogen chloride, one of the clearest demonstrations of polarity.

Identify the atoms and their electronegativities

Hydrogen has an electronegativity of about 2.1 on the Pauling scale, while chlorine sits at 3.0. Chlorine is significantly more electronegative.
Electron sharing is unequal

When H and Cl form a single covalent bond, they share one pair of electrons. But chlorine's stronger pull means the shared electron pair spends more time near the Cl nucleus than near the H nucleus.
Partial charges develop

Because the electron density is greater around chlorine, it acquires a partial negative charge (δ⁻). Hydrogen, now electron-deficient in that region, carries a partial positive charge (δ⁺). We write this as:

\ce{H^{\delta+} - Cl^{\delta-}}

The bond has a dipole moment The dipole moment $\mu$ is defined as:

\mu = q \times d

where $q$ is the magnitude of the partial charge and $d$ is the distance between the charges (the bond length). For HCl, $\mu \approx 1.08 \, \text{D}$ (Debye units). The arrow representing the dipole points from H (positive) toward Cl (negative), or sometimes a crossed arrow ( $\rightarrow$ ) is drawn with the cross at the positive end. …

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Q.Explain with the help of a suitable example a polar covalent bond.

What makes a bond polar?

The hydrogen chloride (HCl) example

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