four atomic orbitals combine to generate

#Four atomic orbitals combine to generate full

Combining waves can lead to constructive interference, in which peaks line up with peaks, or destructive interference, in which peaks line up with troughs ( Figure 5.28). Molecular orbitals are combinations of atomic orbital wave functions. The wave function describes the wavelike properties of an electron. The mathematical process of combining atomic orbitals to generate molecular orbitals is called the linear combination of atomic orbitals (LCAO). In these diatomic molecules, several types of molecular orbitals occur. Such molecules are called homonuclear diatomic molecules. We will consider the molecular orbitals in molecules composed of two identical atoms (H 2 or Cl 2, for example).

#Four atomic orbitals combine to generate full

Like an atomic orbital, a molecular orbital is full when it contains two electrons with opposite spin. The region of space in which a valence electron in a molecule is likely to be found is called a molecular orbital ( Ψ 2). Just like electrons around isolated atoms, electrons around atoms in molecules are limited to discrete (quantized) energies. Using quantum mechanics, the behavior of an electron in a molecule is still described by a wave function, Ψ, analogous to the behavior in an atom. Molecular orbital theory describes the distribution of electrons in molecules in much the same way that the distribution of electrons in atoms is described using atomic orbitals. Both theories provide different, useful ways of describing molecular structure. Table 5.1 summarizes the main points of the two complementary bonding theories. MO theory also helps us understand why some substances are electrical conductors, others are semiconductors, and still others are insulators. Unlike valence bond theory, which uses hybrid orbitals that are assigned to one specific atom, MO theory uses the combination of atomic orbitals to yield molecular orbitals that are delocalized over the entire molecule rather than being localized on its constituent atoms. Additionally, it provides a model for describing the energies of electrons in a molecule and the probable location of these electrons. It also explains the bonding in a number of other molecules, such as violations of the octet rule and more molecules with more complicated bonding (beyond the scope of this text) that are difficult to describe with Lewis structures. Molecular orbital theory (MO theory) provides an explanation of chemical bonding that accounts for the paramagnetism of the oxygen molecule. You can see videos of diamagnetic floating frogs, strawberries, and more. If you place a frog near a sufficiently large magnet, it will levitate. Living things contain a large percentage of water, so they demonstrate diamagnetic behavior. Water, like most molecules, contains all paired electrons. We can calculate the number of unpaired electrons based on the increase in weight. When we compare the weight of a sample to the weight measured in a magnetic field ( Figure 5.27), paramagnetic samples that are attracted to the magnet will appear heavier because of the force exerted by the magnetic field. Magnetic susceptibility measures the force experienced by a substance in a magnetic field. And yet, the Lewis structure of O 2 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 5.1. By itself, O 2 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 O 2: However, one of the most important molecules we know, the oxygen molecule O 2, presents a problem with respect to its Lewis structure.

Relate these electron configurations to the molecules’ 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.By the end of this section, you will be able to: