Kaplan MCAT General Chemistry Review

When two atoms of different electronegativities bond covalently by sharing one or more pairs of electrons, the resulting bond is polar, with the more electronegative atom possessing the greater share of the electron density. However, the mere presence of bond dipoles does not necessarily result in a molecular dipole; that is, an overall separation of charge across the molecule. We must first consider the molecular geometry and the vector addition of the bond dipoles based upon that molecular geometry. A compound with nonpolar bonds is always nonpolar; a compound with polar bonds may be polar or nonpolar, depending upon the spatial orientation of the polar bonds within the given molecular geometry.

A compound composed of two atoms bound by a polar bond must have a net dipole moment and is therefore polar. The two equal and opposite partial charges are localized on the two atoms at the ends of the compound. HCl (hydrogen chloride) is a good example of this, because the bond between the hydrogen and chlorine atom is polar (with the hydrogen atom assuming a partial positive charge and the chlorine atom assuming a partial negative charge); the compound must also be polar, with a molecular dipole moment in the same direction and same magnitude as the bond dipole moment. A compound consisting of more than two atoms bound with polar bonds may be either polar or nonpolar, because the overall dipole moment of a molecule is the vector sum of the individual bond dipole moments. If the compound has a particular molecular geometry such that the bond dipole moments cancel each other (i.e., if the vector sum is zero), then the result is a nonpolar compound. For instance, CCl₄ has four polar C–Cl bonds, but because the molecular geometry of carbon tetrachloride is tetrahedral, the four bond dipoles point to the vertices of the tetrahedron and, therefore, cancel each other out, resulting in a nonpolar compound (see Figure 3.6).

MCAT Expertise

Back to that tug-of-war from earlier, sometimes we can see the winner before the final flag. A polar covalent bond will have one atom that carries more electron density than the other does (and therefore a partial negative charge) but hasn’t won the match yet.

Figure 3.6

However, when the molecular geometry is such that the bond dipoles do not cancel each other, the molecule will have a net dipole moment and, therefore, will be polar. For instance, the O–H bonds in H₂O are polar, with each hydrogen molecule assuming a partial positive charge and the oxygen assuming a partial negative charge. Because water’s molecular geometry is angular (bent), the vector summation of the bond dipoles results in a molecular dipole moment from the partially positive hydrogen end to the partially negative oxygen end, as illustrated in Figure 3.7.

Key Concept

A molecule with polar bonds need not be polar: The bond dipole moments may cancel each other out, resulting in a nonpolar molecule. Although a molecule with polar bonds need not be polar overall, a polar molecule must have polar bonds.

Figure 3.7

ATOMIC AND MOLECULAR ORBITALS

To finish our discussion of covalent bonds, we need to address the issue of atomic and molecular orbitals. If you remember back to the first chapter, we described the modern understanding of the atom as a dense, positively charged nucleus surrounded by clouds of electrons organized into orbitals (regions in space surrounding the nucleus within which there are certain probabilities of finding an electron). The four quantum numbers completely describe the energy and position of any electron of an atom. While the principal quantum number indicates the average energy level of the orbitals, the azimuthal quantum number, l, describes the subshells within each principal energy level, n. When l = 0, this indicates the s-subshell, which has one orbital that is spherical in shape. The 1s-orbital (n = 1, l = 0) is plotted in Figure 3.8.

Bridge

Quantum Numbers (Chapter 1) revisited:

• For any value of n, there are n values of l(0n- 1).

•

l= 0s

l= 1 p

l= 2d

• For any value of l, there are 2l + 1 values of m_l (number of orbitals); values themselves will range from -l to l.

Figure 3.8

When l = 1, this indicates the p-subshell, which has three orbitals shaped like barbells along the x, y, and z axes at right angles to each other. The p-orbitals are plotted in Figure 3.9.

Figure 3.9