Kaplan MCAT General Chemistry Review

Covalent compounds contain discrete molecular units with relatively weak intermolecular interactions. Consequently, these compounds tend to have lower melting and boiling points, and because they do not break down into constituent ions, they are poor conductors of electricity in the liquid state or in aqueous solutions.

PROPERTIES OF COVALENT BONDS

The formation of one covalent bond may not be sufficient to fill the valence shell for a given atom. Thus, many atoms can form bonds with more than one other atom, and each atom can form multiple bonds between itself and another atom, with few exceptions. Two atoms sharing one, two, or three pairs of electrons are said to be joined by a single, double, or triple covalent bond, respectively. The number of shared electron pairs between two atoms is called the bond order; hence, a single bond has a bond order of one, a double bond has a bond order of two, and a triple bond has a bond order of three. Three features characterize a covalent bond: bond length, bond energy, and polarity.

Bond Length

If the arrangement of atoms in covalent bonds can be likened to two wooden spools connected by a dowel in a very sophisticated children’s construction set, then the length of the dowel is analogous to the length of the bond between the atoms sharing the electron pair(s). Bond length is the average distance between the two nuclei of the atoms involved in the bond. As the number of shared electron pairs increases, the two atoms are pulled closer together, leading to a decrease in bond lengths. Thus, for a given pair of atoms, a triple bond is shorter than a double bond, which is shorter than a single bond.

Bridge

We will see a great example of covalent bonds in organic chemistry, and we can see here the inverse proportionality between bond length and strength.

Bond Length Bond Strength C–ClongestweakestC=CmediummediumCCshorteststrongest

Bond Energy

Bond energy is the energy required to break a bond by separating its components into their isolated, gaseous atomic states. Just as it becomes increasingly difficult to snap a tree branch of increasing thickness, the greater the number of pairs of electrons shared between the atomic nuclei, the more energy is required to “break” the bond(s) holding the atoms together. Thus, triple bonds have the greatest bond energy, and single bonds have the lowest bond energy. We will discuss bond energy and calculations involving actual bond energy values (called bond enthalpy) in Chapter 6, Thermochemistry. By convention, the greater the bond energy is, the “stronger” the bond.

Polarity

As in the case of two people who decide to share some commodity for the purpose of achieving a larger goal and who must then decide the degree to which the common commodity will be shared equally or unequally between them (that commodity might be money, land, pizza, or an apartment), atoms that come together in covalent bonds also “negotiate” the degree to which their sharing of electron pair(s) will be equal or unequal. The nature and degree of sharing between the nuclei of two atoms in a covalent bond is determined by the relative difference in their respective electronegativities, with the atom of higher electronegativity getting the larger “share” of the electron pair(s). A polar bond is a dipole, with the positive end of the dipole at the less electronegative atom and the negative end at the more electronegative atom.

Nonpolar Covalent Bond

When atoms that have identical or nearly identical electronegativities share electron pair(s), they do so with equal distribution of the electrons. This is called a nonpolar covalent bond, and there is no separation of charge across the bond. Of course, only bonds between atoms of the same element will have exactly the same electronegativity and, therefore, share with perfectly equal distribution the pair(s) of electrons in the covalent bond. Examples of diatomic molecules include H₂, Cl₂, O₂, and N₂. At the same time, many bonds can be said to be approximately nonpolar. For example, the electronegativity difference between carbon and hydrogen is so sufficiently small that we can usually consider the C–H bond to be effectively nonpolar.

Polar Covalent Bond