Electronegativity, Bond Polarity, Molecular Shape, and Intermolecular Forces

Electronegativity and Bond Polarity

Electronegativity

Electronegativity is a measure of an atom's ability to attract shared electrons in a chemical bond. It is a fundamental property that influences the type and polarity of chemical bonds formed between atoms.

• Trends in the Periodic Table:

  • Electronegativity increases from left to right across a period.

  • Electronegativity increases from bottom to top within a group.

  • Nonmetals, especially fluorine, have the highest electronegativity values.

  • Metals have relatively low electronegativity values.

• Application: Electronegativity values are used to predict bond polarity and molecular properties.

Bond Polarity

The polarity of a bond depends on the difference in electronegativity between the two atoms involved. This difference determines how electrons are distributed in the bond.

• Nonpolar Covalent Bonds:

  • Electrons are shared equally or almost equally.

  • Occurs between atoms with very small or zero electronegativity difference (ΔEN = 0.0–0.4).

  • Example: H2, Cl2, N2.

• Polar Covalent Bonds:

  • Electrons are shared unequally.

  • Occurs between atoms with a moderate electronegativity difference (ΔEN = 0.5–1.8).

  • Example: HCl, H2O, NH3.

• Ionic Bonds:

  • Electrons are transferred from one atom to another.

  • Occurs between atoms with a large electronegativity difference (ΔEN = 1.9–3.3).

  • Example: NaCl, KBr.

Dipoles and Bond Polarity

A dipole is created in a polar covalent bond due to the unequal sharing of electrons, resulting in partial positive (δ+) and partial negative (δ−) charges at opposite ends of the bond. The greater the electronegativity difference, the more polar the bond.

• Dipoles are represented by an arrow pointing toward the more electronegative atom, with a plus sign at the tail.

• Examples of dipoles: C=O, N=O, Cl–F.

Molecular Geometry and VSEPR Theory

Valence Shell Electron-Pair Repulsion (VSEPR) Theory

VSEPR theory is used to predict the three-dimensional shape of molecules based on the repulsion between electron groups (bonding and lone pairs) around a central atom.

• Electron groups arrange themselves as far apart as possible to minimize repulsion.

• Each of the following counts as one electron group:

  • Lone pair

  • Single bond

  • Double bond

  • Triple bond

Common Molecular Shapes

• Linear: 2 electron groups, 180° bond angle (e.g., CO2).

• Trigonal Planar: 3 electron groups, 120° bond angle (e.g., H2CO).

• Bent: 3 electron groups (2 bonds, 1 lone pair), 120° bond angle (e.g., SO2).

• Tetrahedral: 4 electron groups, 109° bond angle (e.g., CH4).

Summary Table: Electron Groups and Molecular Shapes

Electron Groups	Geometry	Bonded Atoms	Lone Pairs	Bond Angle	Molecular Shape	Example
2	Linear	2	0	180°	Linear	CO2
3	Trigonal planar	3	0	120°	Trigonal planar	H2CO
3	Trigonal planar	2	1	120°	Bent	SO2
4	Tetrahedral	4	0	109°	Tetrahedral	CH4
4	Tetrahedral	3	1	109°	Trigonal pyramidal	NH3
4	Tetrahedral	2	2	109°	Bent	H2O

Polarity of Molecules

Nonpolar Molecules

A molecule is nonpolar if it contains only nonpolar bonds or if the dipoles in polar bonds cancel each other out due to the molecule's symmetry.

• Examples: CO2, CH4

Polar Molecules

A molecule is polar if it contains polar bonds and the dipoles do not cancel out, resulting in a molecule with a positive end and a negative end.

• Examples: HCl, H2O, NH3

Steps to Determining Molecular Polarity

1. Determine if the bonds are polar covalent or nonpolar covalent using electronegativity values.

2. If the bonds are polar covalent, draw the Lewis structure and determine if the dipoles cancel.

Intermolecular Forces

Types of Attractive Forces

• Ionic Bonds: Strongest attractive forces, occur between ions, solids at room temperature.

• Dipole-Dipole Attractions: Present in polar covalent compounds; positive end of one molecule is attracted to the negative end of another.

• Hydrogen Bonds: Strongest force between molecules, present when H is bonded to F, O, or N; important in biological molecules like DNA.

• Dispersion Forces: Weak attractions between nonpolar molecules due to temporary dipoles.

Note: Melting point is related to the strength of attractive forces: lower melting points for weak forces (dispersion), higher for strong forces (hydrogen bonds).

Key Definitions and Concepts

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

• Bond Polarity: The unequal sharing of electrons in a bond, leading to dipoles.

• VSEPR Theory: Predicts molecular shape based on electron group repulsion.

• Dipole: A separation of charge in a bond or molecule due to differences in electronegativity.

• Intermolecular Forces: Forces of attraction between molecules, influencing physical properties.

Sample Problems and Applications

• Predicting Bond Type: Use electronegativity difference to classify bonds as nonpolar covalent, polar covalent, or ionic.

• Determining Molecular Shape: Use VSEPR theory to predict the geometry of molecules based on the number of electron groups and lone pairs.

• Classifying Molecular Polarity: Analyze the shape and bond dipoles to determine if a molecule is polar or nonpolar.

• Identifying Intermolecular Forces: Recognize the main attractive forces present in different compounds (ionic, dipole-dipole, hydrogen bonding, dispersion).