Understanding Lewis Structures, Resonance, and Hybridization

Posted on Dec 10, 2024 in Chemistry

Lewis Structures of Bonding Electrons

Lewis Structures represent bonding electrons, non-bonding electrons (lone pairs), and the dipole moment. The dipole moment is the product of the electric charge and the length of the bond. Elements with higher electronegativity attract electrons more strongly.

Writing Lewis Structures

  1. Write the basic structure of the compound, joining the atoms together. The less electronegative atom tends to go in the center.
  2. Count the total number of valence electrons present. The number of valence electrons equals the group number of the atom.
  3. Draw a single bond between the central atom and each of the surrounding atoms. Complete the octets (for Hydrogen, a duet). Non-bonding electrons are represented as lone pairs.
  4. If the central atom does not follow the octet rule, add double or triple bonds between this atom and the surrounding atoms, using the unused pairs of the latter.

Formal Charge in Lewis Structures

For each atom in a valid Lewis structure:

  • Count the number of valence electrons.
  • Subtract all nonbonding electrons.
  • Subtract half of its shared electrons.

Resonance Structures

Resonant structures are converted into each other due to the movement of electrons. The use of curved arrows helps us represent that movement.

Resonance: General Rules

  1. Must be a valid Lewis structure.
  2. Nuclei cannot move, and bond angles should remain the same. Only electrons move (lone pairs and pi electrons change more often).
  3. The core cannot change positions, and bond angles must be the same. All electrons must stay in pairs in all resonance structures.
  4. The number of unpaired electrons must remain the same.
  5. It must satisfy the octet rule. A small separation of charge is allowed. The negative charge will be on the more electronegative atom.
  6. Resonance stabilization is most important when it serves to relocate two or more charges on the atom.

Atomic Orbitals

An orbital is a region of space where the probability of finding an electron is large (the volume may contain an electron 90-95% of the time).

  • The 1s and 2s orbitals represent the core focus areas.
  • Each s orbital can accommodate two electrons.
  • Each 2p orbital has two lobes (near areas) that are very close.
  • There are three 2p orbitals that are perpendicular (orthogonal) to one another.
  • Each orbital can accommodate two electrons, for a total of six.
  • The three p orbitals are degenerate (have the same energy).

Properties of Carbon: Structure of Organic Compounds

Carbon’s small size enables it to form double and triple bonds.

Hybridization

Hybridization is the combination of orbitals belonging to the atomic valence shell to form new orbitals suitable for the qualitative description of bond properties. Hybrid orbitals are very useful for explaining the shape of the orbitals in molecules and, therefore, their geometry. Hybridization is an integral part of valence bond theory.

sp3 Hybridization

The four sp3 orbitals can be oriented at 109.5° angles with respect to each other. The sp3 orbitals give rise to the tetrahedral structure of methane.

Sigma and Pi Bonds

Sigma bonds are symmetric in the cross-section when viewed along the bond. All single bonds are sigma bonds. For example, in ethane, the hypothetical formation of bonding molecular orbitals involves two hybrid carbon atoms and six hydrogen atoms. All bonds are sigma bonds, and the rotation of the bond does not require much energy.

Ethene contains a carbon-carbon double bond and is among the so-called alkenes. Its geometry is trigonal, and the bond angle is 120°. Pi bonds involve parallel p orbitals that overlap above and below the plane of the sigma bond array. Lateral overlapping of p orbitals results in the formation of a pi bond. A pi bond has a nodal plane passing through two linked nuclei and between the lobes of the pi molecular orbital. Bonding and antibonding molecular orbitals: The lower energy pi orbital contains both pi electrons at the baseline. The antibonding orbital has higher energy and is not occupied by electrons at the baseline.

Rotation of double bonds: Maximum overlap between p orbitals of a pi bond occurs when the axes of the orbitals are parallel. A 90° rotation of the carbon-carbon pi bond breaks it. The pi bond strength is 264 kJ/mol, which is the rotational barrier. The rotational barrier of the single bond is 13-26 kJ/mol.

Ethyne (acetylene): sp hybridization. The sp orbital is a linear combination of s and p valence orbitals of the central atom. One s and one p orbital give two sp orbitals (linear geometry). The sp hybrid orbital has lobes set at an angle of 180° to each other. The carbon-carbon triple bond consists of one sigma and two pi bonds. The decrease in the C-H bond distance is associated with more s character in the orbital.

Molecular Geometry

The three-dimensional distribution of atoms in a molecule is responsible for its physical and chemical properties. The electrons in the valence shell repel each other and are responsible for the bond. The final geometry is one that signals a lower repulsion between electrons.

Functional Groups and Homologous Series

A functional group is an atom or group of atoms joined together in a characteristic way, preferably determining the properties of the compound in which they are present. A homologous series is a series of compounds whose members differ from the following one by a constant value.

Hydrocarbons

  • Open chain (alkanes, alkenes, alkynes)
  • Closed chain (cyclic, aromatic)