# Electronic Structure and Bonding

## Overview

### Description

Understanding the electronic structure and bonding properties of elements and molecules is critical to predicting how chemical reactions will occur. One of the ways elements are organized in the periodic table is according to the number of electrons in their outermost shell. Electrons are held in orbitals, which may overlap to allow elements to share electrons and form covalent bonds. Carbon uses hybridized orbitals to form four covalent bonds. The electronegativity difference between the two elements bonding together determines their bond strength. Intermolecular forces govern how molecules interact with each other based on permanent dipoles within individual molecules.

### At A Glance

• Lewis structures are used to show the connectivity of organic compounds.
• The octet rule states that atoms prefer to have eight valence electrons when forming compounds. Oxygen, nitrogen, and the halogens have lone pairs that are part of their valence shell. Boron is an exception because it forms compounds with only six valence electrons; hydrogen and helium are also exceptions as they follow the duet rule and only have two valence electrons.
• Formal charge determines the charge of each atom in a Lewis structure.
• When compounds form, the connectivity of the atoms will determine the structures that they form. In forming compounds, atoms without a formal charge will have a typical number of bonds.
• Generally, covalent bonds between atoms with identical (or almost identical) electronegativity will be nonpolar covalent bonds. Covalent bonds between atoms with different electronegativity will be polar covalent bonds. If the difference is big enough, the compound will have an ionic bond. Molecules with polar covalent bonds that do not cancel will have a net dipole moment.
• Quantum mechanics is the study of the wavelike properties of electrons. An atomic orbital represents a region of space occupied by electrons.
• The Aufbau principle, Pauli's exclusion principle, and Hund's rule determine the electron configuration of an atom.
• The valence bond theory and molecular orbital theory help explain how atomic orbitals overlap to form covalent bonds.
• Carbon atoms cannot form four bonds in their unhybridized state. Carbon must hybridize to form four covalent bonds. The hybridization states of carbon are sp3, sp2, and sp.
• Bond length and bond strength of carbon are determined by the hybridization state of carbon. Bond length order is $sp^3>sp^2>sp$. Bond strength order is $sp>sp^2>sp^3$. VSEPR theory can help predict the molecular geometry of organic compounds. Most organic compounds have tetrahedral, trigonal planar, or linear geometry.
• All organic molecules have London (dispersion) forces. Some organic molecules have dipole-dipole interactions, and even fewer have hydrogen bonding. Ionic interactions are seen between cations and anions, and ion-dipole forces can be seen between an ion and a polar molecule. Intermolecular forces are responsible for the melting point, boiling point, and water solubility or insolubility of organic molecules.