Covalent Bonding, Polarity, and Molecular Structure: Study Guide

Covalent Bonding and Molecular Structure

What is a Covalent Bond?

Covalent bonds are a fundamental type of chemical bond where two atoms share one or more pairs of electrons. This sharing allows each atom to achieve a stable electron configuration, often following the octet rule.

• Source of Attraction: The mutual attraction between the positively charged nuclei and the shared electrons.

• Types of Covalent Bonds: Single (one pair shared), double (two pairs shared), and triple (three pairs shared) bonds.

• Lone Pairs: Non-bonding pairs of electrons that can influence molecular shape and reactivity.

Example: In a water molecule (H2O), each hydrogen shares one electron with oxygen, forming two single covalent bonds.

Polarity

Polarity in chemistry refers to the distribution of electrical charge over the atoms joined by the bond. It is determined by the difference in electronegativity (END) between the bonded atoms.

• Bond Polarity: Based on electronegativity difference (END) between atoms.

• Classification by END:

  • END > 2.0: Ionic bond

  • 1.9 ≥ END ≥ 0.5: Polar covalent bond

  • 0.4 ≥ END: Nonpolar covalent bond

• Molecular Polarity: Determined by both bond polarity and molecular geometry (shape).

Example: In HCl, the END between H and Cl is large, making the bond polar covalent.

Determining Molecular Polarity

• Polar Molecules: Asymmetrical shape, presence of lone pairs on the central atom, or different atoms attached to the central atom.

• Nonpolar Molecules: Symmetrical shape, same atoms on all sides of the central atom, and no lone pairs on the central atom (or only one element present).

Example: Carbon dioxide (CO2) is nonpolar due to its linear, symmetrical shape, even though the bonds are polar.

Intermolecular Forces (IMFs)

Intermolecular forces are attractions between molecules that influence physical properties like boiling and melting points.

• Types of IMFs:

  1. London Dispersion Forces: Present in all molecules, especially nonpolar ones; caused by temporary dipoles.

  2. Dipole-Dipole Forces: Occur between polar molecules due to permanent dipoles.

  3. Hydrogen Bonding: A strong type of dipole-dipole force, occurs when H is bonded to N, O, or F.

• Source of Attraction: Differences in charge distribution within or between molecules.

Example: Water exhibits hydrogen bonding, leading to its high boiling point.

Naming Covalent Compounds

Covalent compounds are named using prefixes to indicate the number of each atom present. Special rules apply to avoid double vowels and unnecessary prefixes.

• Prefixes: mono-, di-, tri-, tetra-, penta-, hexa-, etc.

• Special Rules: Do not use 'mono-' for the first element; avoid double vowels (e.g., monoxide, not monooxide).

Example: CO2 is named carbon dioxide, not monocarbon dioxide.

Lewis Structures

Lewis structures are diagrams that show the arrangement of electrons in a molecule. They help predict molecular shape, reactivity, and formal charge.

• Drawing Lewis Structures: Show all valence electrons as dots or lines for bonds.

• Resonance Structures: Multiple valid Lewis structures for a molecule; recognize but do not draw all forms.

• Formal Charge: Used to determine the most stable Lewis structure. Calculated as:

$\text{Formal Charge} = (\text{Valence electrons}) - (\text{Non-bonding electrons}) - \frac{1}{2}(\text{Bonding electrons})$

• Central Atom: Usually the least electronegative element (except hydrogen).

Example: In nitrate ion (NO3-), resonance structures distribute the negative charge over the oxygen atoms.

VSEPR Theory and Molecular Shape

Valence Shell Electron Pair Repulsion (VSEPR) theory predicts the 3D shape of molecules based on electron pair repulsion around the central atom.

• Basis: Electron pairs (bonding and lone pairs) repel each other and arrange themselves as far apart as possible.

• Common Shapes: Linear, trigonal planar, tetrahedral, trigonal bipyramidal, octahedral.

• Matching Lewis Structures to Shapes: Use the number of bonding and lone pairs to determine geometry.

Example: Methane (CH4) is tetrahedral; water (H2O) is bent due to two lone pairs on oxygen.

Mass Spectrometry and Structure

Mass spectrometry (Mass Spec) is used to determine the structure of molecules by analyzing the mass-to-charge ratio of fragments.

• Connecting Peaks to Structure: Peaks correspond to fragments; helps identify what is central and what is connected in the molecule.

Example: A peak at a certain mass may indicate the presence of a methyl group (CH3).

Summary Table: Bond Types by Electronegativity Difference

Electronegativity Difference (END)	Bond Type
> 2.0	Ionic
1.9 – 0.5	Polar Covalent
< 0.4	Nonpolar Covalent

Practice and Application

• Draw Lewis structures for given molecules and ions.

• Determine molecular shape using VSEPR theory.

• Assign formal charges and identify the most stable structure.

• Name covalent compounds and write their formulas using prefixes.

• Classify bonds and molecules as polar or nonpolar using electronegativity and geometry.

• Identify intermolecular forces present in a compound.