Electronegativity, Bond Polarity, and Covalent Bond Energies: Study Notes

Electronegativity and Bond Polarity

Introduction to Electronegativity and Bond Polarity

Atoms in molecules do not always share electrons equally, leading to differences in charge distribution and molecular properties. This unequal sharing is fundamental to understanding chemical bonding and molecular behavior.

• Electronegativity (EN): The tendency of an atom to attract shared electrons in a covalent bond.

• Bond Polarity: A measure of the unequal sharing of electrons in a bond, resulting in partial charges.

• Polar Covalent Bond: A covalent bond in which electrons are shared unequally, creating a dipole moment.

• Nonpolar Covalent Bond: Electrons are shared equally between atoms.

• Ionic Bond: Electrons are transferred from one atom to another, resulting in full charges.

Visualizing Bond Polarity

Polar molecules, such as HF, align with an external electric field due to their charge separation. This alignment demonstrates the presence of a molecular dipole.

• Partial Charges: Denoted by δ+ (partial positive) and δ− (partial negative).

• Example: In HF, F is more electronegative and gains δ−, while H gains δ+.

Electronegativity Trends

Electronegativity values are determined from thermochemical data and range from 0 (low) to 4 (high). The most electronegative element is fluorine (F).

• Periodic Trend: Electronegativity increases across a period (left to right) and decreases down a group (top to bottom).

• Comparison with Electron Affinity: Electronegativity refers to an atom's ability to attract electrons in a bond, while electron affinity is the energy change when an atom gains an electron.

Bond Polarity and Electronegativity Difference

The greater the difference in electronegativity (ΔEN) between two atoms, the more polar the bond.

• ΔEN ≤ 0.4: Nonpolar covalent bond

• ΔEN = 0.5–2.0: Polar covalent bond

• ΔEN > 2.0: Ionic bond

• Charge Distribution: The atom with higher EN gains a partial negative charge (δ−), and the atom with lower EN gains a partial positive charge (δ+).

Examples and Applications

• Order of Increasing Electronegativity: C < N < O < F

• Most Polar Bond: The bond with the greatest electronegativity difference, e.g., P-F is more polar than P-Cl.

Covalent Bond Energies, Lengths, and Vibrations

Bond Energy and Bond Length

Bond energy is the energy required to break a bond. The strength and length of a bond are inversely related: shorter bonds are generally stronger.

• Bond Energy: Measured in kJ/mol, it quantifies the stability of a bond.

• Bond Length: The distance between the nuclei of two bonded atoms, measured in picometers (pm).

• Multiple Bonds: Double and triple bonds are stronger and shorter than single bonds.

Bond	Bond Length (pm)	Bond Strength (kJ/mol)
N–F	139	272
N–Cl	191	200
N–Br	214	243
N–I	222	159

Bond	Bond Length (pm)	Bond Strength (kJ/mol)
C≡C	120	837
C=C	134	611
C–C	154	347

Bond Energy Calculation Example

Bond energy can be related to photon energy using the equation:

• $E = h u$

• $u = rac{c}{ ext{wavelength}}$

Example: A 256 nm photon has enough energy to break the C–Cl bond in CCl4, generating a Cl radical.

Bond Vibrations and Infrared Absorption

Visible and UV light cause electronic excitation, while infrared (IR) radiation causes vibrational excitation in molecules.

• IR Absorption: Molecules absorb IR light, causing bond vibrations to increase in amplitude.

• IR Activity: For a vibrational mode to be IR active, it must cause a change in the dipole moment of the molecule.

• Characteristic Frequencies: Different bonds vibrate at specific frequencies, allowing IR spectroscopy to identify bond types.

Additional info:

• Electron affinity is a property of isolated atoms, while electronegativity is a property of atoms in molecules.

• Bond strength exceptions may occur due to specific atomic interactions or resonance effects.