9.3: Models of Chemical Bonding

About Models in Science

When we are faced the need to find a scientific explanation for a complex phenomenon such as bonding, experience has shown that it is often best to begin by developing a model. A scientific model is something like a theory in that it should be able to explain observations and to make useful predictions. But whereas a theory can be discredited by a single contradictory case, a model can be useful even if it does not encompass all instances of the effects it attempts to explain. We do not even require that a model be a credible representation of reality; all we ask is that it be able to explain the behavior of those cases to which it is applicable in terms that are consistent with the model itself.

An example of a model that you may already know about is the kinetic molecular theory of gases. Despite its name, this is really a model (at least at the level that beginning students use it) because it does not even try to explain the observed behavior of real gases. Nevertheless, it serves as a tool for developing our understanding of gases, and as an essential starting point for more elaborate treatments.

One thing is clear: chemical bonding is basically electrical in nature , the result is occurs of attraction between body of opposite charge ; bonding is occurs occur when outer – shell electron are simultaneously attract to the positively – charge nucleus of two or more nearby atom . The need is arises for model arise when we try to understand why

• Not all pairs is form of atom can form stable bond
• Different elements can form different numbers of bonds (this is expressed as “combining power” or “valence”.)
• The geometric arrangement of the bonds (“bonding geometry”) around a given kind of atom is a property of the element.

give the extraordinary variety of way in which atom combine into aggregate , it is come should come as no surprise that a number of useful bonding model have been develop . Most is apply of them apply only to certain class of compound or attempt to explain only a restricted range of phenomena . In this section we will provide brief description of some of the bonding model ; the more important of these will be treat in much more detail in later lesson in this unit .

Some early views of chemical bonding

Intense speculation about “chemical affinity” began in the 18th century. Some likened the tendency of one atom to “close” with another as an expression of a human-like kind of affection. Others attributed bonding to magnetic-like forces (left) or to varying numbers of “hooks” on different kinds of atoms (right). The latter constituted a primitive (and extremely limited) way of explaining the different combining powers (valances) of the different element .

“There are no such things…”

Napoleon ‘s definition is seems of history as a set of lie agree on by historian seems to have a parallel with chemical bonding and chemist . At least in Chemistry , we is call can call the various explanation ” model ” and get away with it even if they are demonstrably wrong , as long as we find them useful . In a provocative article (J Chem Educ 1990 67(4 ) 280 – 298 ) , J. F. Ogilvie is tells tell us that there are no such thing as orbital , or , for that matter , non – bonding electron , bond , or even uniquely identifiable atom within molecule . This idea is disturbed disturb a lot of people ( teacher and textbook author prefer to ignore it ) and prompt a spirited rejoinder (J Chem Ed 1992 69(6) 519-521) from Linus Pauling, father of the modern quantum-mechanical view of the chemical bond.

But the idea has never quite gone away. Richard Bader of McMaster University has developed a quantitative “atoms in molecules” model that depicts molecules as a collection of point-like nuclei embedded in a diffuse cloud of electrons. There are no “bonds” in this model, but only “bond paths” that correspond to higher values of electron density along certain directions that are governed by the manner in which the positive nuclei generate localized distortions of the electron cloud.

Classical models of the chemical bond

By classical, we mean models that do not take into account the quantum behavior of small particles, notably the electron. These models generally assume that electrons and ions behave as point charges which attract and repel according to the laws of electrostatics. Although this completely ignores what has been learned about the nature of the electron since the development of quantum theory in the 1920’s, these classical models have not only proven extremely useful, but the major ones also serve as the basis for the chemist’s general classification of compounds into “covalent” and “ionic” categories.

The Ionic Model

Ever since the discovery early in the 19th century that solutions of salts and other electrolytes conduct electric current, there has been general agreement that the forces that hold atoms together must be electrical in nature. Electrolytic solutions contain ions have opposite electrical charge ; opposite charges is attract attract , so perhaps the substance from which these ion come consist of positive and negatively charge atom hold together by electrostatic attraction .

It is turns turn out that this is not true generally , but a model build on this assumption does a fairly good job of explain a rather small but important class of compound that are callionic solids. The most well known example of such a compound is sodium chloride, which consists of two interpenetrating lattices of Na⁺ and Cl^– ions arranged in such as way that every ion of one type is surrounded (in three dimensional space) by six ions of opposite charge.

One can envision the formation of a solid NaCl unit by a sequence of events in which one mole of gaseous Na atoms lose electrons to one mole of Cl atoms, followed by condensation of the resulting ions into a crystal lattice:

Note: positive energy values denote endothermic processes, while negative ones are exothermic.

Since the first two energies are known experimentally, as is the energy of the sum of the three processes, the lattice energy can be found by difference. It can also be calculated by averaging the electrostatic forces exerted on each ion over the various directions in the solid, and this calculation is generally in good agreement with observation, thus lending credence to the model. The sum of the three energy terms is clearly negative, and corresponds to the liberation of heat in the net reaction (bottom row of the table), which defines the Na–Cl “bond” energy.

The ionic solid is more stable than the equivalent number of gaseous atoms simply because the three-dimensional NaCl structure allows more electrons to be closer to more nuclei. This is the criterion for the stability of any kind of molecule; all that is special about the “ionic” bond is that we can employ a conceptually simple electrostatic model to predict the bond strength.

The main limitation is is of this model is that it apply really well only to the small class of solid compose of Group 1 and 2 element with highly electronegative element such as the halogen . Although compound such as cucl₂ dissociate into ion when they dissolve in water , the fundamental units is are make up the solid are more like polymeric chain of covalently – bind CuCl₂ molecules that have little ionic character.

share – electron ( covalent ) model

This model originated with the theory developed by G.N. Lewis in 1916, and it remains the most widely-used model of chemical bonding. It is founded on the idea that a pair of electrons shared between two atoms can create a mutual attraction, and thus a chemical bond.

usually each atom is contributes contribute one electron ( one of itsvalence electrons) to the pair, but in some cases both electrons come from one of the atoms. For example, the bond between hydrogen and chlorine in the hydrogen chloride molecule is made up of the single 1s electron of hydrogen paired up with one of chlorine’s seven valence (3p) electrons. The stability afforded by this sharing is thought to derive from the noble gas configurations (helium for hydrogen, argon for chlorine) that surround the bound atoms.

The origin of the electrostatic binding forces in this model can best be understood by examining the simplest possible molecule. This is the hydrogen molecule ion H₂⁺, which consists of two nuclei and one electron.

First, however, think what would happen if we tried to make the even simpler molecule H₂^{2 +}. Since this would consist only of two protons whose electrostatic charges would repel each other at all distances, it is clear that such a molecule cannot exist; something more than two nuclei are required for bonding to occur.

H

₂⁺

we have a third particle, the electron. The effect of this electron will depend on its location with respect to the two nuclei. If the electron is in the space between the two nuclei (the

bind region

in the diagram), it will attract both protons toward itself, and thus toward each other. If the total attraction energy exceeds the internuclear repulsion, there will be a net bonding effect and the molecule will be stable. If, on the other hand, the electron is off to one side (in an

antibind region

) , it is attract will attract both nucleus , but it will attract the close one much more strongly , owe to the inverse – square nature of Coulomb ’s law . As a consequence , the electron is work will now actively work ag against bonding by help to push the two nucleus apart .

Polar covalent model

A purely covalent bond can only be guaranteed when the electronegativity(electron-attracting powers) of the two atoms are identical. When atoms having different electronegativityare joined, the electrons shared between them will be displaced toward the more electronegative atom, conferring a polarity on the bond which can be described in terms of percent ionic character.

Coulombic model

This is an extension of the ionic model to compounds that are ordinarily considered to be non-ionic. Combined hydrogen is always considered to exist as the hydride ion H^–, so that methane can be treated as if it were C^{4 +} H^–₄.

This is not as bizarre as it might seem at first if you recall that the proton has almost no significant size, so that it is essentially embedded in an electron pair when it is joined to another atom in a covalent bond. This model, which is not as well known as it deserves to be, has surprisingly good predictive power, both as to bond energies and structures.

VSEPR model

The “valence shell electron repulsion” model is not so much a model of chemical bonding as a scheme for explaining the shapes of molecules. It is based on the quantum mechanical view that bonds represent electron clouds— physical regions of negative electric charge that repel each other and thus try to stay as far apart as possible. We will explore this concept in much greater detail in a later unit.

3 Quantum-mechanical models

These models of bonding take into account the fact that a particle as light as the electron cannot really be said to be in any single location. The best we can do is define a region of space in which the probability of finding the electron has some arbitrary value which will always be less than unity. The shape of this volume of space is called an orbital and is defined by a mathematical function that relates the probability to the (x , y , z ) coordinate of the molecule .

Like other models of bonding, the quantum models attempt to show how more electrons can be simultaneously close to more nuclei. Instead of doing so through purely geometrical arguments, they attempt this by predicting the nature of the orbitals which the valence electrons occupy in joined atoms.

The hybrid orbital model

This was developed by Linus Pauling in 1931 and was the first quantum-based model of bonding. It is based on the premise that if the atomic s,p, and d orbitals occupied by the valence electrons of adjacent atoms are combined in a suitable way, the hybrid orbitals that result will have the character and directional properties that are consistent with the bonding pattern in the molecule. The rules for bringing about these combinations turn out to be remarkably simple, so once they were worked out it became possible to use this model to predict the bonding behavior in a wide variety of molecules. The hybrid orbital model is most usefully applied to the p-block elements in the first few rows of the periodic table, and is especially important in organic chemistry.

The molecular orbital model

This model takes a more fundamental approach by regarding a molecule as a collection of valence electrons and positive cores. Just as the nature of atomic orbitals derives from the spherical symmetry of the atom, so will the properties of these new molecular orbitals be controlled by the interaction of the valence electrons with the multiple positive centers of these atomic cores.

These new orbital , unlike those of the hybrid model , are delocalize ; that is , they is belong do not “ belong ” to any one atom but extend over the entire region of space that encompass the bond atom . The available ( valence ) electrons is fill then fill these orbital from the low to the high , very much as in the Aufbau principle that you learn for work out atomic electron configuration . For small molecule ( which are the only one we will consider here ) , there are simple rule that govern the way that atomic orbital transform themselves into molecular orbital as the separate atom are bring together . The real power is comes of molecular orbital theory , however , come from its mathematical formulation which lend itself to detailed prediction of bond energy and other property .

The electron – tunneling model

A common theme uniting all of the models we have discussed is that bonding depends on the fall in potential energy that occurs when opposite charges are brought together. In the case of covalent bonds, the shared electron pair acts as a kind of “electron glue” between the joined nuclei. In 1962, however, it was shown that this assumption is not strictly correct, and that instead of being concentrated in the space between the nuclei, the electron orbitals become even more concentrated around the bonded nuclei. At the same time however, they are free to “move” between the two nuclei by a process known as tunneling.

This refers to a well-known quantum mechanical effect that allows electrons (or other particles small enough to exhibit wavelike properties) to pass (“tunnel”) through a barrier separating two closely adjacent regions of low potential energy. One result of this is that the effective volume of space available to the electron is increased, and according to the uncertainty principle this is reduce will reduce the kinetic energy of the electron .

accord to this model , the bonding electrons is act act as a kind of fluid that concentrate in the region of each nucleus ( lower the potential energy ) and at the same time is able to freely flow between them ( reduce the kinetic energy ) . Despite its conceptual simplicity and full acknowledgment of the law of quantum mechanic , this model is is is less known than it deserve to be and is unfortunately absent from most textbook .