S2.2.5 Bond polarity

Electronegativity Differences and Dipole Moments in Covalent Bonds

Imagine pouring water into a glass and then adding oil. Despite vigorous shaking, the two liquids stubbornly refuse to mix. Have you ever wondered why?

The answer lies in bond polarity—a concept that explains how electrons are shared between atoms in a molecule.

Electronegativity: The Driving Force Behind Bond Polarity

Definition

Electronegativity

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

Analogy

Think of it as a "tug-of-war" for electrons—the stronger the atom's pull, the higher its electronegativity.

Example

Fluorine (F), the most electronegative element, has a value of 4.0.
Hydrogen (H), with a lower electronegativity of 2.2, exerts a weaker pull.

Electronegativity Differences and Bond Types

The difference in electronegativity ( $Δ χ$ ) between two bonded atoms determines the type of bond they form:

Nonpolar Covalent Bond ( $Δ χ = 0$ ): Electrons are shared equally (e.g., $H_{2}$ ).
Polar Covalent Bond ( $0 < Δ χ < 1.8$ ): Electrons are shared unequally, creating partial charges (e.g., $H C l$ ).
Ionic Bond ( $Δ χ > 1.8$ ): Electrons are transferred completely, forming ions (e.g., $N a C l$ ).

Example

Classifying Bond Types

Carbon–Hydrogen (C–H):
- Electronegativity of C = 2.6, H = 2.2.
- $Δ χ = | 2.6 - 2.2 | = 0.4$ , so the bond is nonpolar covalent.
Hydrogen–Fluorine (H–F):
- Electronegativity of H = 2.2, F = 4.0.
- $Δ χ = | 4.0 - 2.2 | = 1.8$ , so the bond is highly polar covalent.

Tip

Electronegativity differences provide a guideline, but remember that bonding exists on a continuum rather than in strict categories.

Hint

Electronegativity values can be found in Section 9 of the Data Booklet.

Bond Polarity and Partial Charges (δ+ and δ−)

Definition

Bond polarity

Bond polarity refers to the unequal sharing of electrons between two atoms in a covalent bond, resulting in a partial positive charge on one atom and a partial negative charge on the other.

When two atoms in a bond have different electronegativities, the shared electrons spend more time near the more electronegative atom. This creates:

A partial negative charge ( $δ^{-}$ ) on the more electronegative atom.
A partial positive charge ( $δ^{+}$ ) on the less electronegative atom.

Example

In $H C l$ :

Chlorine ( $χ = 3.2$ ) attracts the electrons more strongly, becoming $δ^{-}$ .
Hydrogen ( $χ = 2.2$ ) becomes $δ^{+}$ .

Representing Bond Polarity

Bond polarity can be visualized in two ways:

Partial Charges ( $δ^{+}$ and $δ^{-}$ ):
- Example: $H^{δ +} - C l^{δ -}$ .
Dipole Moment Vector ( $\vec{μ}$ ):

Definition

Dipole moment vector

A dipole moment vector is a vector pointing from the less electronegative atom to the more electronegative atom.

The arrow starts with a “+” sign at the positive end and points toward the negative end. Example: $H \to C l$ .

Tip

When drawing dipole moment vectors, always use molecular geometry to determine the direction and magnitude of the dipoles.

Illustration of dipole moment for hydrochloric acid.

Dipole Moments: Measuring Bond Polarity

Definition

Dipole moment

A dipole moment describes the separation of electrical charge in a bond or molecule due to differences in electronegativity.

It results in partial positive and negative regions, making the molecule polar.

Example

In water $(H_{2} O)$ , oxygen attracts electrons more strongly than hydrogen, creating a dipole. A larger dipole moment indicates a greater charge separation, influencing properties like solubility and boiling point.

Molecular Dipole Moment vs. Bond Dipoles

While individual bonds may have dipoles, the overall molecular dipole depends on the geometry of the molecule:

Polar Molecules: Bond dipoles do not cancel out, resulting in a net dipole moment (e.g., $H_{2} O$ ).
Nonpolar Molecules: Bond dipoles cancel out due to symmetry, resulting in no net dipole moment (e.g., $C O_{2}$ ).

Example

Water vs. Carbon Dioxide

Water ( $H_{2} O$ ):Geometry: Bent. The bond dipoles of the two $O - H$ bonds do not cancel, resulting in a net dipole moment. -> Polar molecule.
Carbon Dioxide ( $C O_{2}$ ):Geometry: Linear. The bond dipoles of the two $C = O$ bonds cancel each other out. -> Nonpolar molecule.

Hint

Use VSEPR theory to predict molecular geometry and determine whether bond dipoles will cancel.

Applications of Bond Polarity and Dipole Moments

Explaining Physical Properties

Boiling and Melting Points: Polar molecules experience stronger intermolecular forces (e.g., dipole-dipole interactions), leading to higher boiling points.

Example

Water ( $H_{2} O$ ) has a higher boiling point than methane ( $C H_{4}$ ) due to hydrogen bonding.

Solubility: Polar molecules dissolve in polar solvents (e.g., water), while nonpolar molecules dissolve in nonpolar solvents (e.g., oil). Like dissolves like.

Common Mistake

Mistake: Assuming all molecules with polar bonds are polar.

Solution: Always consider molecular geometry. Symmetrical molecules (e.g., $C O_{2}$ ) can have polar bonds but still be nonpolar overall.

Common Mistake

Mistake: Forgetting to calculate electronegativity differences.

Solution: Always use $Δ χ$ to determine bond polarity and compare it to the bonding continuum.

Practice and Reflection

Self review

Calculate the electronegativity difference for the bonds in $H F$ , $C H_{4}$ , and $N H_{3}$ . Classify each bond as nonpolar, polar, or ionic.
Draw the dipole moment vector for $H C l$ and $C O_{2}$ .
Explain why $C H C l_{3}$ is polar but $C C l_{4}$ is nonpolar.

Theory of Knowledge

How does the concept of electronegativity reflect the balance between empirical data and theoretical models in science?
Can you think of other fields where abstract models explain real-world phenomena?

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S2.2.5 Bond polarity

Electronegativity Differences and Dipole Moments in Covalent Bonds

Electronegativity: The Driving Force Behind Bond Polarity

Electronegativity Differences and Bond Types

Classifying Bond Types

Bond Polarity and Partial Charges (δ+ and δ−)

Representing Bond Polarity

Dipole Moments: Measuring Bond Polarity

Molecular Dipole Moment vs. Bond Dipoles

Water vs. Carbon Dioxide

Applications of Bond Polarity and Dipole Moments

Explaining Physical Properties

Practice and Reflection

You've read 2/2 free chapters this week.

Introduction to Bond Polarity