7.3 Covalent Bonding
Covalent bonding is made from the combination of two or more nonmetals. Nonmetals in the covalent bonding do not transfer electron (valence electrons) instead they share them to form single, double or even triple bonds.
An example is given below depicting the covalent bonding formation of PCl3 between phosphorous and chlorine atoms: The covalent single bonds are designated with shared valence electrons 
Figure 7.23 Covalent Bond Formation
Or:
Figure 7.24 Sharing of Electrons in Covalent Bond Formation
There are three single bonds and one set of lone pair in PCl3.
Another example: Methane CH4. Hydrogen is ingroup 1 A, and it has one valence electron and carbon is in group 4 A, and therefore it has four valence electrons. Methane has four single covalent bonds.
Figure 7.25 Sharing of Electrons in methane
Covalent compounds
The shared-electron pair model introduced by G.N. Lewis showed how chemical bonds could form in the absence of electrostatic attraction between oppositely-charged ions. As such, it has become the most popular and generally useful model of bonding in all substances other than metals. A chemical bond occurs when electrons are simultaneously attracted to two nuclei, thus acting to bind them together in an energetically-stable arrangement. The covalent bond is formed when two atoms are able to share a pair of electrons:
Figure 7.29 Formation of covalent Bond
In general, however, different kinds of atoms exert different degrees of attraction on their electrons, so in most cases the sharing will not be equal. One can even imagine an extreme case in which the sharing is so unequal that the resulting “molecule” is simply a pair of ions:
Figure 7.30 Formation of Ion Pair
The resulting substance is sometimes said to contain an ionic bond. Indeed, the properties of a number of compounds can be adequately explained using the ionic model. But does this mean that there are really two kinds of chemical bonds, ionic and covalent?
Comparison Study: Bonding in ionic solids
According to the ionic electrostatic model, solids such as NaCl consist of positive and negative ions arranged in a crystal lattice. Each ion is attracted to neighboring ions of opposite charge, and is repelled by ions of like charge; this combination of attractions and repulsions, acting in all directions, causes the ion to be tightly fixed in its own location in the crystal lattice.
Figure 7.31 Formation of Ionic Crystal
Since electrostatic forces are non-directional, the structure of an ionic solid is determined purely by geometry: two kinds of ions, each with its own radius, will fall into whatever repeating pattern will achieve the lowest possible potential energy. Surprisingly, there are only a small number of possible structures; one of the most common of these, the simple cubic lattice of NaCl, is shown here.
Is there really such a thing as an ionic bond?
When two elements form an ionic compound, is an electron really lost by one atom and transferred to the other one? In order to deal with this question, consider the data on the ionic solid LiF. The average radius of the neutral Li atom is about 2.52Å. Now if this Li atom reacts with an atom of F to form LiF, what is the average distance between the Li nucleus and the electron it has “lost” to the fluorine atom? The answer is 1.56Å; the electron is now closer to the lithium nucleus than it was in neutral lithium!
Figure 7.32 Formation of Ionic bond
So the answer to the above question is both yes and no: yes, the electron that was now in the 2s orbital of Li is now within the grasp of a fluorine 2p orbital, but no, the electron is now even closer to the Li nucleus than before, so how can it be “lost”? The one thing that is inarguably true about LiF is that there are more electrons closer to positive nuclei than there are in the separated Li and F atoms. But this is just the rule we stated at the beginning of this unit: chemical bonds form when electrons can be simultaneously near two or more nuclei.
It is obvious that the electron-pair bond brings about this situation, and this is the reason for the stability of the covalent bond. What is not so obvious (until you look at the numbers such as are quoted for LiF above) is that the “ionic” bond results in the same condition; even in the most highly ionic compounds, both electrons are close to both nuclei, and the resulting mutual attractions bind the nuclei together. This being the case, is there really any fundamental difference between the ionic and covalent bond?
The answer, according to modern chemical thinking is probably “no”; in fact, there is some question as to whether it is realistic to consider that these solids consist of “ions” in the usual sense. The preferred picture that seems to be emerging is one in which the electron orbitals of adjacent atom pairs are simply skewed so as to place more electron density around the “negative” element than around the “positive” one.
This being said, it must be reiterated that the ionic model of bonding is a useful one for many purposes, and there is nothing wrong with using the term “ionic bond” to describe the interactions between the atoms in the very small class of “ionic solids” such as LiF and NaCl.
More on polar covalence
“Covalent, ionic or metallic” is an oversimplification!
If there is no such thing as a “completely ionic” bond, can we have one that is completely covalent? The answer is yes, if the two nuclei have equal electron attracting powers. This situation is guaranteed to be the case with homonuclear diatomic molecules– molecules consisting of two identical atoms. Thus in Cl2, O2, and H2, electron sharing between the two identical atoms must be exactly even; in such molecules, the center of positive charge corresponds exactly to the center of negative charge: halfway between the two nuclei.
Figure 7.33 Metallic, covalent and ionic Compounds of different elements
Categorizing all chemical bonds as either ionic, covalent, or metallic is a gross oversimplification; as this diagram shows, there are examples of substances that exhibit varying degrees of all three bonding characteristics.
- A bond’s percent ionic character (discussed more in detail later under “Electronegativity section) is the amount of electron sharing between two atoms; limited electron sharing corresponds with a high percent ionic character.
- To determine a bond’s percent ionic character, the atoms’ electronegativities are used to predict the electron sharing between the atoms.
Electron donor-acceptor bonds
In most covalent bonds, we think of the electron pair as having a dual parentage, one electron being contributed by each atom. There are, however, many cases in which both electrons come from only one atom. This can happen if the donor atom has a non-bonding pair of electrons and the acceptor atom has a completely empty orbital that can accommodate them.
Figure 7.34 Electronic arrangement in donor-acceptor bond
This is the case, for example, with boron trifluoride and ammonia. In BF3, one the 2p orbitals is unoccupied and can accommodate the lone pair on the nitrogen atom of ammonia. The electron acceptor, BF3, acts as a Lewis acid here, and NH3 is the Lewis base.
Bonds of this type (sometimes known as coordinate covalent or dative bonds) tend to be rather weak (usually 50-200kJ/mol); in many cases the two joined units retain sufficient individuality to justify writing the formula as a molecular complex or adduct.
Single, Double and Triple Covalent Bonding
Covalent bonding requires a very discrete and specific orientation to establish the overlap between the bonding orbitals. The overlap can form sigma (single) bonds or pi (double and triple) bonds. the sigma bonds have very strong interaction due to the fact shared electrons are moving freely along the axis and the overlapping orbital are lying exactly on the orbital axis, while the pi bonds have weak interaction because the overlapping orbitals are laying above or below the orbital axis.
Single bonds are made when two electrons are shared, they form one sigma bond between the two atoms. Double bonds are made when four electrons are shared between the two atoms. As a result, sigma bond and one pi bond is formed. Triple bonds occur when six electrons are shared between the two atoms and as result one sigma bond and two pi bonds are formed.
More information on this type of bonding is discussed in next module.(Module 8) Advance Theories of Bonding.
The videos below illustrate the concept of the covalent bonds.
Reference: https://www.youtube.com/watch?v=ZxWmyZmwXtA
Reference: https://www.youtube.com/watch?v=S_k0kr2eZSQ
Polar and Nonpolar Covalent Bonds
Two important types of covalent bonds are nonpolar or pure covalent bonds and polar covalent bonds. Nonpolar bonds occur when atoms equally share electron pairs. Since only identical atoms (having the same electronegativity) truly engage in equal sharing, the definition is expanded to include covalent bonding between any atoms with an electronegativity difference less than 0.4. Examples of molecules with nonpolar bonds are H2, N2, and CH4.
As the electronegativity difference increases, the electron pair in a bond is more closely associated with one nucleus than the other. If the electronegativity difference is between 0.4 and 1.7, the bond is polar. If the electronegativity difference is greater than 1.7, the bond is ionic.
Covalent Bond Examples
There is a covalent bond between the oxygen and each hydrogen in a water molecule (H2O). Each of the covalent bonds contains two electrons, one from a hydrogen atom and one from the oxygen atom. Both atoms share the electrons.
A hydrogen molecule, H2, consists of two hydrogen atoms joined by a covalent bond. Each hydrogen atom needs two electrons to achieve a stable outer electron shell. The pair of electrons is attracted to the positive charge of both atomic nuclei, holding the molecule together.
Phosphorus can form either PCl3 or PCl5. In both cases, the phosphorus and chlorine atoms are connected by covalent bonds. PCl3 assumes the expected noble gas structure, in which the atoms achieve complete outer electron shells. Yet PCl5 is also stable, so it’s important to remember covalent bonds in chemistry don’t always abide by the octet rule.
Reference: https://www.youtube.com/watch?v=RK07Bwwxkjg