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 and from bottom to top within a group.

• Nonmetals generally have higher electronegativity values, with fluorine being the most electronegative element.

• Metals have lower electronegativity values.

Bond Polarity

The polarity of a bond depends on the difference in electronegativity between the two atoms involved:

• Nonpolar Covalent Bonds: Electrons are shared equally (e.g., H2, Cl2), typically between identical nonmetals or atoms with very similar electronegativity values.

• Polar Covalent Bonds: Electrons are shared unequally, resulting in partial charges (δ+ and δ-) on the atoms (e.g., HCl, H2O). The greater the electronegativity difference, the more polar the bond.

Examples of Dipoles in Polar Covalent Bonds

Predicting Bond Type Using Electronegativity Difference

• Nonpolar Covalent: Electronegativity difference < 0.5

• Polar Covalent: Electronegativity difference between 0.5 and 1.7

• Ionic: Electronegativity difference > 1.7

Molecular Geometry: VSEPR Theory

Valence Shell Electron-Pair Repulsion (VSEPR) Theory

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

• Electron groups are arranged as far apart as possible to minimize repulsion.

• The number of electron groups determines the molecular geometry.

• Each of the following counts as one electron group: a lone pair, a single bond, a double bond, or a triple bond.

Linear Geometry Example: CO2

• Two electron groups around the central atom

• Bond angle: 180°

• Shape: Linear

Other Common Molecular Shapes

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

• Bent: Two bonds and one or more lone pairs (e.g., SO2)

• Tetrahedral: Four electron groups, 109° bond angles (e.g., CH4)

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, CCl4

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

Determining Molecular Polarity: Steps

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.

Example: PBr3

• P = 2.1, Br = 2.8 (polar covalent)

• Shape: Trigonal pyramidal (one lone pair on P)

• Result: Polar molecule

Example: CH4

• C = 2.5, H = 2.1 (nonpolar covalent)

• Shape: Tetrahedral

• Result: Nonpolar molecule (dipoles cancel)

Intermolecular Forces

Types of Attractive Forces

• Ionic Bonds: Strongest attractive forces, typically found in ionic compounds (e.g., NaCl). 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: Special type of dipole-dipole attraction; occurs when H is bonded to F, O, or N. Strongest force between molecules, 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. Compounds with weak forces (dispersion) have lower melting points, while those with strong forces (hydrogen bonds) have higher melting points.

Summary Table: Types of Intermolecular Forces

References

• Timberlake, K. (2018). Chemistry: An Introduction to General, Organic and Biological Chemistry (13th ed.). Pearson Education.