(credit: modification of work by Jefferson Lab) The detailed explanation of bonding described in this chapter allows us to understand this phenomenon. Other diatomic molecules (like N2) flow past the magnet. However, when we pour liquid oxygen through a magnet, the molecules line up with the magnetic field, and the attraction allows them to stay suspended between the poles of the magnet where the magnetic field is strongest. Oxygen molecules orient randomly most of the time, as shown in the top magnified view. And yet, the Lewis structure of O2 indicates that all electrons are paired. Such attraction to a magnetic field is called paramagnetism, and it arises in molecules that have unpaired electrons. Thus, when we pour liquid oxygen past a strong magnet, it collects between the poles of the magnet and defies gravity, as in Figure 22.1. By itself, O2 is not magnetic, but it is attracted to magnetic fields. However, this picture is at odds with the magnetic behavior of oxygen. There is an O=O double bond, and each oxygen atom has eight electrons around it. This electronic structure adheres to all the rules governing Lewis theory. We would write the following Lewis structure for O2: However, one of the most important molecules we know, the oxygen molecule O2, presents a problem with respect to its Lewis structure. Relate these electron configurations to the moleculesâ stabilities and magnetic propertiesįor almost every covalent molecule that exists, we can now draw the Lewis structure, predict the electron-pair geometry, predict the molecular geometry, and come close to predicting bond angles.Write molecular electron configurations for first- and second-row diatomic molecules.Calculate bond orders based on molecular electron configurations.Describe traits of bonding and antibonding molecular orbitals.Outline the basic quantum-mechanical approach to deriving molecular orbitals from atomic orbitals.Correctly predicting the magnetic properties of molecules is in advantage of molecular orbital theory over Lewis structures and valence bond theory.īy the end of this section, you will be able to: Materials with unpaired electrons are paramagnetic and attracted to a magnetic field, while those with all-paired electrons are diamagnetic and repelled by a magnetic field. Electrons fill molecular orbitals following the same rules that apply to filling atomic orbitals Hund’s rule and the Aufbau principle tell us that lower-energy orbitals will fill first, electrons will spread out before they pair up, and each orbital can hold a maximum of two electrons with opposite spins. We can describe the electronic structure of diatomic molecules by applying molecular orbital theory to the valence electrons of the atoms. Molecular orbitals formed from p orbitals oriented in a side-by-side fashion have electron density on opposite sides of the internuclear axis and are called π orbitals. They can be formed from s orbitals or from p orbitals oriented in an end-to-end fashion. Molecular orbitals located along an internuclear axis are called σ MOs. Antibonding molecular orbitals result from out-of-phase combinations of atomic wave functions and electrons in these orbitals make a molecule less stable. Bonding molecular orbitals are formed by in-phase combinations of atomic wave functions, and electrons in these orbitals stabilize a molecule. The resulting molecular orbitals may extend over all the atoms in the molecule. Molecular orbital (MO) theory describes the behavior of electrons in a molecule in terms of combinations of the atomic wave functions.

Li2 molecular orbital diagram