Figure 3.5
In describing the shape of a molecule, only the arrangement of atoms (not electrons) is considered. Even though the electron pairs are arranged tetrahedrally, the shape of NH3 is pyramidal. It is not trigonal planar because the lone pair repels the three bonding electron pairs, causing them to move as far away as possible.
Example: Predict the geometry of CO2.
Solution: The Lewis structure of CO2 is
The double bond behaves just like a single bond for purposes of predicting molecular shape. This compound has two groups of electrons around the carbon. According to the VSEPR theory, the two sets of electrons will orient themselves 180° apart, on opposite sides of the carbon atom, minimizing electron repulsion. Therefore, the molecular structure of CO2 is linear:
Key Concept
The shapes from Table 3.2 refer to electronic geometry, which is different from molecular geometry. In Figure 3.5, we see an ammonia molecule that has a tetrahedral electronic structure but is considered to have a molecular structure that is trigonal pyramidal.
One subtlety that students sometimes miss is the difference between electronic geometry and molecular geometry. Electronic geometry describes the spatial arrangement of all pairs of electrons around the central atom, including the bonding and the lone pairs. Molecular geometry describes the spatial arrangement of only the bonding pairs of electrons, in much the same way that the geometry of an object made from our sophisticated construction set would be determined by the position of the wooden spools attached to the dowels. For example, consider the fact that CH4 (methane), NH3 (ammonia), and H2O all have the same electronic geometry: In each compound, four pairs of electrons surround the central atom. This is tetrahedral electronic geometry. However, because each has a different number of lone pairs, each has a different molecular geometry. Methane has tetrahedral geometry, ammonia has trigonal pyramidal, and water has angular or bent geometry. The distinction is important, and the MCAT will primarily focus on molecular geometry, but there is one important implication of electronic geometry: the determination of the ideal bond angle. Tetrahedral electronic geometry, for example, is associated with an ideal bond angle of 109.5°. Thus, molecular geometries that deviate from tetrahedral electronic geometry, such as those of ammonia and water, have bond angles that are deviations from 109.5°. You may have been tempted to say, for example, that water’s bond angle is a deviation from linear geometry with its ideal bond angle of 180°, but this is not the case. The actual bond angle in water is around 104.5°, a deviation from 109.5°.
Polarity of Molecules
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(0
•
l= 0
l= 1
l= 2