Polar Covalent Bonds, Acids, and Bases: Structure, Polarity, and Noncovalent Interactions

Polar Covalent Bonds and Electronegativity

Bond Types and Ionic Character

Chemical bonds can be classified based on the distribution of electrons between the bonded atoms. The degree of electron sharing or transfer determines whether a bond is covalent, polar covalent, or ionic.

• Covalent Bonds: Electrons are shared equally between atoms with similar electronegativities.

• Polar Covalent Bonds: Electrons are shared unequally due to a difference in electronegativity, resulting in partial charges (δ+ and δ-).

• Ionic Bonds: Electrons are transferred from one atom to another, creating ions with full charges.

Electronegativity

Electronegativity (EN) is the intrinsic ability of an atom to attract shared electrons in a covalent bond. Differences in EN between atoms lead to bond polarity. The EN scale is arbitrary, with fluorine (F) being the most electronegative (EN = 4.0) and cesium (Cs) the least (EN = 0.7).

• Metals (left side of the periodic table) have low EN.

• Nonmetals (right side, especially halogens) have high EN.

• EN difference < 2: polar covalent bond; EN difference > 2: ionic bond.

Bond Polarity and Inductive Effect

Bonding electrons are drawn toward the more electronegative atom, resulting in partial charges. The inductive effect describes the shifting of electrons in response to the EN of nearby atoms.

Dipole Moments and Molecular Polarity

Dipole Moment

The dipole moment (μ) quantifies molecular polarity and is defined as the product of the magnitude of the charge (Q) and the distance (r) between charges:

$\mu = Q \times r$

Units: debye (D), where 1 D = $3.336 \times 10^{-30}$ C·m. Molecules with large dipole moments are strongly polar and tend to be soluble in polar solvents like water.

Dipole Moments of Common Compounds

The table below summarizes dipole moments for several compounds, illustrating the relationship between molecular structure and polarity.

Compound	Dipole moment (D)	Compound	Dipole moment (D)
NaCl	9.00	NH3	1.47
CH2O	2.33	CH3NH2	1.31
CH3Cl	1.87	CO2	0
H2O	1.85	CH4	0
CH3OH	1.70	CH3CH3	0
CH3CO2H	1.70	Benzene	0
CH3SH	1.52

Formal Charges

Definition and Calculation

Formal charge is used to keep track of electron distribution in molecules. It is calculated by comparing the number of valence electrons in the free atom to the number assigned in the molecule:

$\text{Formal charge} = \text{Valence electrons} - (\text{Number of bonds} + \text{Nonbonding electrons})$

• If an atom has one more electron than its neutral state, it is assigned a negative charge.

• If it has one less, it is assigned a positive charge.

Common Formal Charges Table

Atom	C	N	O	S	P
Valence electrons	4	5	6	6	5
Number of bonds	4	4	3	3	4
Number of nonbonding electrons	0	0	2	2	0
Formal charge	0	+1	+1	+1	+1

Resonance and Resonance Hybrids

Resonance Forms

Some molecules cannot be accurately represented by a single Lewis structure. Instead, they are described by resonance forms, which differ only in the placement of π bonds or lone pairs. The actual structure is a resonance hybrid, more stable than any individual form.

• Resonance forms are connected by a double-headed arrow.

• Electrons move in pairs, indicated by curved arrows.

• Resonance forms must obey the octet rule and be valid Lewis structures.

Acids and Bases

Brønsted–Lowry Definition

The Brønsted–Lowry theory defines acids and bases by their ability to donate or accept protons (H+):

• Acid: Proton donor

• Base: Proton acceptor

When an acid reacts with a base, conjugate acid-base pairs are formed.

Acid and Base Strength: Ka and pKa

The acid dissociation constant (Ka) measures the strength of an acid in water. The pKa is the negative logarithm of Ka:

$\mathrm{p}K_a = -\log K_a$

• Lower pKa = stronger acid

• Higher pKa = weaker acid

Predicting Acid–Base Reactions

The direction of acid-base reactions can be predicted using pKa values. The acid with the lower pKa will donate a proton to the base with the higher pKa conjugate acid.

Organic Acids and Bases

Organic Acids

Organic acids are characterized by the presence of a positively polarized hydrogen atom. They typically lose a proton from an O–H or C–H group adjacent to a carbonyl (C=O).

Organic Bases

Organic bases have atoms with lone pairs that can bond to H+. Nitrogen-containing compounds are the most common, but oxygen-containing compounds can also act as bases.

Lewis Acids and Bases

Lewis Definition

The Lewis definition broadens the concept of acids and bases:

• Lewis acid: Electron pair acceptor

• Lewis base: Electron pair donor

This definition includes metal cations and compounds with vacant orbitals as Lewis acids.

Molecular Models and Noncovalent Interactions

Molecular Models

Molecular models are useful for visualizing the three-dimensional structure of molecules, which is critical for understanding their chemical behavior.

Noncovalent Interactions

Noncovalent interactions are forces between molecules that do not involve sharing or transfer of electrons. They include:

• Dipole-dipole forces: Occur between polar molecules due to electrostatic interactions among dipoles.

• Dispersion forces: Present in all molecules, arising from temporary fluctuations in electron distribution.

• Hydrogen bonds: Strongest noncovalent interaction, occurring between a hydrogen atom bonded to an electronegative atom (O or N) and another electronegative atom with a lone pair.

Summary

• Polar covalent bonds arise from differences in electronegativity, leading to molecular polarity measured by dipole moment (μ).

• Formal charges help track electron distribution in molecules.

• Resonance hybrids represent delocalized electron structures, increasing stability.

• Brønsted–Lowry acids donate protons; bases accept protons. Acid strength is measured by pKa.

• Lewis acids accept electron pairs; Lewis bases donate electron pairs.

• Noncovalent interactions (dipole-dipole, dispersion, hydrogen bonds) are crucial for molecular properties and biological function.