Electrostatic Interpretation of Different Bond Types
Now that we have the concept of electronegativity, we are ready to explain the different types of bonding within the electrostatic interpretation of the atomic bond. As discussed in the answer to the last question of Exercise 3.5.1, all atomic bonds are the result of increased electron density between two nuclei. The difference between different bond types is due to how the increased electron density is distributed. As we will see, the different types of bonds are actually on a spectrum.
Covalent to Ionic Bonds
In a covalent bond, the bonding atoms have nearly equal electronegativity. This means that they attract electrons equally. So, when electron density shifts to be between the two nuclei, it is equally spaced between them. This is the sense in which covalent bonds "share" electrons. However, what causes the atoms to bond is the electrostatic forces between the nuclei and the increased electron density between them. The fact that it is equally shared is simply a result of the electrons' equal attraction to both nuclei. The left side of Figure 3.7.1 below illustrates this with a bond between two oxygen atoms.
In an ionic bond, the two atoms have very unequal electronegativity. So, when electron density shifts to be between the two nuclei, it is pulled more towards the side of the more electronegative atom. The left side of Figure 3.7.1 below illustrates this with a bond between a sodium and a chlorine atom.
In the case of a pure covalent bond, the difference in electronegativity between the atoms is zero. As the electronegativity difference increases, the bond will become more ionic. However, there will never be a complete transfer of electrons to the more electronegative atoms, because then there wouldn't be increased electron density between the two nuclei and there would be no bond. The only time that the electron completely transfers is when the two atoms disassociate and the more electronegative atom keeps the electron (e.g., when NaCl dissolves in water and becomes $\text{Na}^+$ and $\text{Cl}^+$). It is still a useful classification, but it is important to realize that the amount of electron "sharing" vs "transfer" is actually on a spectrum.
In your chemistry course, you probably learned about the "octet rule" for explaining bonding. How the octet rule relates to the electrostatic interpretation of bonding will be explored on the next page (Section 3.8).
Figure 3.7.1 Cartoon illustration of electron density shift in a pure covalent bond between two oxygen atoms and a highly ionic bond between a sodium atom and a chlorine atom. The pairs of little red circles on the outsides of the electron clouds represent the number of electron pairs in the valence shells of the atoms, both before and after bonding.
In the context of materials science, we mostly discuss materials with extended bonding networks/structures, not discrete molecules. Figure 3.7.2 and Figure 3.7.3 depict ionic and covalent bonding in solids in terms of shifted electron density. Although these figures are cartoons, modern imaging techniques can allow us to "see" electron density distribution in real bonds, such as the image of pentacene constructed with the help of an atomic force microscope. You can see that example here.
Figure 3.7.2 Cartoon depiction of electron density in an ionic solid of NaCl.
Figure 3.7.3 Cartoon depiction of electron density in graphene, a 2D covalent solid of carbon atoms.
Covalent to Metallic Bonds
Metallic bonds are similar to covalent bonds in that electron density is shared approximately equally between the bonding atoms. The difference is that electron density in metallic bonds is much more delocalized, i.e., spread out. Figure 3.7.4 illustrates the difference between covalent and metallic bonds on a spectrum of how delocalized the electrons are. In general, metallic bonding leads to solids as depicted in Figure 3.7.5. In a solid metal, valence electrons are very delocalized which leads to various properties we will discuss later in the course.
What would cause electrons to be more or less delocalized? The answer again is related electronegativity. Oxygen atoms have an electronegativity of 3.44 while aluminum has an electronegativity of 1.61. This means that oxygen atoms exert a much stronger attractive force on valence electrons, causing them to localize between the two nuclei. In contrast, aluminum atoms exert a much weaker force on valence electrons. There is still an increase in electron density between nuclei, but it is less localized.
A second related factor in electron delocalization is whether all the valence orbitals are filled. Solid metals typically do not obey the octet "rule" and only have partially filled valence orbitals. When this happens, the electrons fractionally occupy all the valence orbitals—instead of filling some and leaving others empty—allowing them to delocalize and exist between multiple nuclei at once (you could also think of them as hopping around between different pairs of nuclei). These fractional bonds are discussed further in Chapter 13 on electronic properties. The reason this is related to electronegativity is that, in general, the fewer valence electrons an atom has, the less electronegative it will be (as seen in the electronegativity trends of the periodic table discussed in the solution to Question 3.6.1.1). So, the fact that aluminum has only three valence electrons is related to both its relatively low electronegativity and the fact that it doesn't fill all its valence orbitals in a solid. This is true for other metals as well.
Figure 3.7.4 Covalent bonds, such as in the $\text{O}_2$ molecule, are highly localized electron density between nuclei resulting from the high electronegativity of the nuclei. Metallic bonds, such as in solid aluminum, are delocalized with electron density much more spread out. Of course, there still has to be in increase in electron density between the nuclei compared to non-bonded atoms. The pairs of little red circles on the outsides of the electron clouds represent the number of electron pairs in the valence shells of the atoms, both before and after bonding.
Figure 3.7.5 Cartoon depiction of a layer of aluminum atoms bonded together. The bonding electrons are delocalized and overlapping with neighboring atoms such that you can't say that an electron really "belongs" to one pair of atoms rather than another.
Secondary Bonds
So far, we have discussed the three primary types of bonds. There are also so-called "secondary" bonds (or simply secondary interactions) such as hydrogen bonds and those resulting from Van der Waals interactions. What causes these bonds? As with all other bonds, secondary bonds are the result of increased electron density between nuclei. The difference with secondary bonds is simply that the valence shells of the atoms involved are already filled. This means that any nearly electrons feel a much weaker force from the nucleus of these atoms because they are highly shielded.
Figure 3.7.6 depicts the weak bonds that can form between noble gases from shifting electron density. Since they have full valence shells, the effective nuclear charge that outside electrons feel is close to zero. So, the electron clouds of two adjacent atoms with filled valances only feel a very weak attraction - or shift in the electron cloud. This is why these bonds are so weak and why noble gases only form liquids and solids at extremely low temperatures.
Figure 3.7.7 depicts hydrogen bonds between water molecules. In a single water molecule, electron density is concentrated more around the oxygen atom than the hydrogen atoms because the oxygen atom is more electronegative. When multiple water molecules interact, the negatively charged part of the molecule around the oxygen nucleus attracts the positively charged parts of other water molecules around the hydrogen nuclei. These bonds are much stronger than those in noble gases, but still generally weaker than covalent, ionic, or metallic bonds.
Figure 3.7.6 Even with noble gasses that have full valence shells, the electron density can shift slightly, resulting in weak bonds between atoms. Since these bonds are so weak, noble gases only form liquids and solids at extremely low temperatures.
Figure 3.7.7 Hydrogen bonds form between the hydrogen nuclei in a molecule and the electron density concentrated around a more electronegative atom in another molecule. The example depicted here is in water molecules.
Predicting Bond Types
We now have everything we need to know to be able to approximately predict bond types. If one of the atoms already has a full valence shell, then we will only have secondary bonding. If both atoms have incomplete valence shells, then primary bonds can form. The type will depend on the electronegativities of both of the two atom types. It isn't enough to know the difference in their electronegativities ($\Delta \chi$), we need to know both numbers. Or, alternatively, we need to know ($\Delta \chi$) and the average electronegativity, $\chi_{\text{avg}}$. Either way, we need two pieces of information, not one. Here is how we can predict the three primary bond types:
- Covalent bonds will form when both electronegativities are high. Put in other terms, when $\Delta \chi$ is small and $\chi_{\text{avg}}$ is high.
- Ionic bonds will form when one electronegativity is high and one is low. Put in other terms, when $\Delta \chi$ is high.
- Metallic bonds will form when both electronegativities are low. Put in other terms, when $\Delta \chi$ is small and $\chi_{\text{avg}}$ is low.
As we have already discussed, these bonds are on a spectrum. For example, if both electronegativities are high, but not exactly the same, the bond will have a little bit of ionic character. The more different the electronegativities become, the more the ionic character.
Figure 3.7.8 shows two different graphical "maps" of bond type based on plotting binary compounds according to two electronegativity numbers.
Figure 3.7.8 The van Arkel-Ketelaar triangle of bonding which plots compounds on two axes of electronegativity difference vs average electronegativity. The vertices of the triangle represent the three ideal bond types and different regions of the triangle represent where each type of bond is predominant. Even though the bond types are on a continuum, it is often still useful to identify the predominant type.