Covalent Bonding
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Covalent bonding results from the sharing of electrons between multiple nuclei. Molecular compounds (i.e., compounds that feature only covalent bonding) consist of individual molecules. Molecular compounds are essentially always composed only of nonmetals (or nonmetals and metalloids, which come from the right side of the periodic table as well). Hydrogen chloride (HCl), water, acetic acid (CH$$_3$$COOH, the acid found in vinegar), and ammonia (NH$$_3$$) are all molecular compounds, and they all consist exclusively of nonmetals. In contrast to ionic compounds, molecular compounds tend to have lower melting points.
In contrast to ionic bonds that tend to form when one atom gives one or more electrons to another atom, covalent bonding occurs when neither atom is “willing” to completely give up its electrons. Instead, the two atoms, red and blue in Figure 1 above, share the electrons, allowing both atoms to have access to eight electrons (completing their valence shell, or octet). (Note: Hydrogen only needs two electrons to fill its valence shell. When bonded to another atom, hydrogen has a "duet" of electrons.) Electrons are shared equally or unequally, depending on the atom (we will examine this later).
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Figure 2 illustrates covalent bonding in an H$$_2$$ molecule. Two hydrogen atoms, with their lone electrons, come together to share their electrons with each other (as described in the reaction equation, Figure 1). This allows each atom to have access to two electrons, completing the duet of each hydrogen. The bond forms because it releases energy. This is because when bonded, the electrons are able to interact with both nuclei simultaneously, in effect exposing each electron to more positive charge than it would experience in a lone H atom. There is, of course, some repulsion between the two electrons and between the two nuclei, but overall, the attractive interactions outweigh the repulsive interactions (i.e., there more attractive interactions in H$$_2$$ than there in the two isolated atoms). Since opposite charges are, on average, being placed closer together, the bonded molecule is lower in potential energy than the separated atoms.
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Figure 3 on the right shows the region of space an electron could inhabit between two nuclei in which it would experience greater electrostatic interactions than it would in an isolated atom.
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The curve shown in Figure 4 describes energy as a function of the distance between two bonded nuclei (i.e., internuclear distance). When the nuclei are far apart (on the far right), there is no interaction. As the nuclei come closer together, the nuclei begin to share electrons and form a bond. As that happens, the energy goes down since electrons and protons are brought closer together on average. As the distance between the nuclei continues to decrease, energy reaches some minimum before beginning to increase again. The distance at which energy is minimized is the bond distance. This describes the physical distance between two bonded nuclei (i.e., the length of the bond). The reason energy increases as distance continues to decrease is more complicated than it may seem, and is beyond the scope of this course.
The bond dissociation energy (often abbreviated BDE) describes the energy needed to break a bond. As shown on the graph, it is the difference in energy between the two bonded atoms and the two lone, non-bonded atoms.
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