Polar Covalent Bonds and Electronegativity

Introduction to Covalent Bonding

Chemical bonds can be classified based on how electrons are shared between atoms. The nature of electron sharing determines whether a bond is covalent, polar covalent, or ionic.

• Covalent bond: Electrons are shared equally between two atoms.

• Polar covalent bond: Electrons are shared unequally, resulting in partial charges on the atoms.

• Ionic bond: Electrons are transferred from one atom to another, creating ions.

The continuum from covalent to ionic bonding is determined by the difference in electronegativity between the bonded atoms.

Electronegativity

Electronegativity is a measure of an atom's ability to attract electrons in a chemical bond. The greater the difference in electronegativity between two atoms, the more polar the bond.

• If the electronegativity difference () is less than 0.5, the bond is nonpolar covalent.

• If is between 0.5 and 2.0, the bond is polar covalent.

• If is greater than 2.0, the bond is ionic.

For example, the electronegativities of carbon (2.5) and hydrogen (2.1) are close, so C–H bonds are considered nonpolar.

Electronegativity Values and Trends

Electronegativity increases from left to right across a period and decreases from top to bottom in a group.

Additional info: The table above is a simplified version of the periodic trend for electronegativity.

Polar Covalent Bonds and Dipole Moments

Bond Polarity and Dipole Moments

Polar covalent bonds result in charge separation, which can be visualized using electrostatic potential maps. The direction of bond polarity is indicated by a crossed arrow, with the tail (plus sign) at the positive end and the head at the negative end.

• Dipole moment (): A measure of the separation of positive and negative charges in a molecule.

The dipole moment is calculated as:

where is the magnitude of the charge and is the distance between charges.

Examples of Dipole Moments

Additional info: Molecules with symmetrical charge distribution (e.g., CO2, CH4) have zero dipole moment.

Formal Charges

Definition and Calculation

Formal charge is a bookkeeping tool used to keep track of electron distribution in molecules. It may not correspond to actual charges but helps in understanding reactivity and resonance.

The formal charge (FC) is calculated as:

• For carbon: 4 valence electrons

• For nitrogen: 5 valence electrons

• For oxygen: 6 valence electrons

Example: In methane (CH4), carbon has no formal charge.

Resonance Forms

Introduction to Resonance

Some molecules cannot be adequately represented by a single Lewis structure. Resonance forms are multiple Lewis structures that differ only in the arrangement of electrons. The true structure is a resonance hybrid, a weighted average of all valid forms.

• Resonance stabilizes molecules by delocalizing charge or electrons.

• Curved arrows are used to show electron movement between resonance forms.

Rules for Resonance Forms

1. Resonance forms differ only in the position of electrons, not atoms.

2. Valid resonance forms obey normal rules of valency.

3. The resonance hybrid is more stable than any individual resonance form.

Examples of Resonance

• Acetate ion: Two resonance forms with negative charge delocalized over two oxygens.

• Benzene: Six-membered ring with alternating double bonds; resonance forms show delocalization.

Acids and Bases: The Brønsted–Lowry Definition

Brønsted–Lowry Acids and Bases

The Brønsted–Lowry definition classifies acids as proton donors and bases as proton acceptors.

• Acid: Donates a proton (H+).

• Base: Accepts a proton.

• When an acid donates a proton, it forms its conjugate base.

• When a base accepts a proton, it forms its conjugate acid.

General reaction:

Acid Strength and pKa

Acid strength is measured by the acid dissociation constant () and its logarithmic counterpart, .

• Lower means a stronger acid.

• Higher means a weaker acid.

Table: pKa Values of Common Acids and Their Conjugate Bases

Acid-Base Reactions from pKa Values

Predicting Reaction Direction

Acid-base reactions favor the formation of the weaker acid and base (higher values). The equilibrium lies toward the side with the weaker acid.

Example: Reaction of acetic acid () with hydroxide ion () favors formation of water and acetate ion.

Acids and Organic Bases

Organic Acids and Their Stabilization

Organic acids such as methanol, acetic acid, and acetone differ in acidity due to the stabilization of their conjugate bases.

• Negative charge on a highly electronegative atom (e.g., oxygen) stabilizes the anion.

• Resonance stabilization further increases acidity.

Example: Acetic acid is more acidic than methanol because its conjugate base is stabilized by resonance.

Acids and Bases: The Lewis Definition

Lewis Acids and Bases

The Lewis definition broadens the concept of acids and bases:

• Lewis acid: Electron pair acceptor (vacant orbital).

• Lewis base: Electron pair donor (filled orbital).

Lewis acid-base reactions involve the sharing of an electron pair to form a covalent bond.

Examples of Lewis Acids and Bases

• Lewis acids: BF3, AlCl3, H+, metal cations

• Lewis bases: NH3, H2O, ethers, amines

Noncovalent Interactions between Molecules

Types of Intermolecular Forces

Noncovalent interactions, also called intermolecular forces, influence physical properties such as melting point, boiling point, and solubility.

• Dipole-dipole forces: Attraction between polar molecules.

• Dispersion forces (London forces): Temporary dipoles in nonpolar molecules.

• Hydrogen bonding: Strong dipole-dipole attraction involving N–H or O–H groups.

Hydrogen Bonding in Biological Molecules

Hydrogen bonds are crucial in stabilizing structures such as DNA and proteins. For example, hydrogen bonds between base pairs hold the two strands of DNA together.

Solubility and Polarity

Polar molecules (hydrophilic) dissolve in water due to hydrogen bonding, while nonpolar molecules (hydrophobic) do not. For instance, vitamin C (ascorbic acid) is water-soluble due to multiple –OH groups, whereas vitamin A (retinol) is fat-soluble due to its nonpolar structure.