The electron pairs shared between two different atoms are usually not shared equally. We can visualize two extreme cases in the degree to which electron pairs are shared. At one extreme we have bonding between two identical atoms, as in Cl2 or N2, where the electron pairs must be equally shared. At the other extreme, illustrated by NaCl, there will be essentially no sharing of electrons. We know that the compound in this case is best described as composed of Na+ and Cl– ions. The 3s electron of the Na atom is, in effect, transferred completely to chlorine. The bonds occurring in most covalent substances fall somewhere between these extremes.
The concept of bond polarity is useful in describing the sharing of electrons between atoms. A nonpolar covalent bond is one in which the electrons are shared equally between two atoms. In a polar covalent bond one of the atoms exerts a greater attraction for the bonding electrons than the other. If the difference in relative ability to attract electrons is large enough, an ionic bond is formed.
We use a quantity called electronegativity to estimate whether a given bond will be nonpolar covalent, polar covalent, or ionic. Electronegativity is defined as the ability of an atom in a molecule to attract electrons to itself. The greater an atom's electronegativity, the greater is its ability to attract electrons to itself. The electronegativity of an atom in a molecule is related to its ionization energy and electron affinity, which are properties of isolated atoms. The ionization energy measures how strongly an atom holds on to its electrons. Likewise, the electron affinity is a measure of how strongly an atom attracts additional electrons. An atom with a very negative electron affinity and high ionization energy will both attract electrons from other atoms and resist having its electrons attracted away; it will be highly electronegative.
Numerical estimates of electronegativity can be based on a variety of properties, not just ionization energy and electron affinity. The first and most widely used electronegativity scale was developed by the American chemist Linus Pauling (1901-1994), who based his electronegativity scale on thermochemical data. Figure 8.7 shows Pauling's electronegativity values for many of the elements. The values are unitless. Fluorine, the most electronegative element, an electronegativity of 4.0. The least electronegative element, cesium, has an electronegativity of 0.7. The values for all other elements lie between these two extremes.
Figure 8.7 Electronegativities of the elements.
Note that within each period there is generally a steady increase in electronegativity from left to right; that is, from the most metallic to the most nonmetallic elements. Notice also that, with some exceptions (especially within the transition metals), electronegativity decreases with increasing atomic number in any one group. This is what we might expect because we know that ionization energies tend to decrease with increasing atomic number in a group and electron affinities don't change very much. You do not need to memorize numerical values for electronegativity. However, you should know the periodic trends so that you can predict which of two elements is the more electronegative.
We can use the difference in electronegativity between two atoms to gauge the polarity of the bonding between them. Consider these three fluorine-containing compounds:
In F2 the electrons are shared equally between the fluorine atoms, and the covalent bond is nonpolar. A nonpolar covalent bond results when the electronegativities of the bonded atoms are equal.
In HF the fluorine atom has a greater electronegativity than does the hydrogen atom, so the sharing of electrons is unequal; the bond is polar. A polar covalent bonds results when the atoms differ in electronegativity. In HF the more electronegative fluorine atom attracts electron density away from the less electronegative hydrogen atom. Thus, some of the electron density around the hydrogen nucleus is pulled toward the fluorine nucleus, leaving a partial positive charge on the hydrogen atom and a partial negative charge on the fluorine atom. We can represent this situation as:
The + and - (read "delta plus" and "delta minus") symbolize the partial positive and negative charges, respectively.
In LiF the far greater electronegativity of fluorine compared with that of lithium leads to the complete transfer of the valence electron of Li to F. This transfer results in the formation of Li+ and F– ions; the resultant bond is therefore ionic.
As these examples illustrate, the greater the difference in electronegativity between two atoms, the more polar their bond. The nonpolar covalent bond lies at one end of a continuum of bond types, and the ionic bond lies at the other end. In between is a broad range of polar covalent bonds, differing in the extent to which there is unequal sharing of electrons.
Which bond is more polar: (a) BCl or CCl; (b) PF or PCl? Indicate in each case which atom has the partial negative charge.
SOLUTION (a) The difference in the electronegativities of chlorine and boron is 3.0 - 2.0 = 1.0; the difference between chlorine and carbon is 3.0 - 2.5 = 0.5. Consequently, the BCl bond is the more polar; the chlorine atom carries the partial negative charge because it has a higher electronegativity. We should be able to reach this same conclusion without using a table of electronegativities; we can rely on periodic trends. Because boron is to the left of carbon in the periodic table, we would predict that it has a lower attraction for electrons. Chlorine, being on the right side of the table, has a strong attraction for electrons. The more polar bond will be the one between the atoms having the lowest attraction for electrons (boron) and the highest attraction (chlorine).
(b) Because fluorine is above chlorine in the periodic table, we would predict it to be more electronegative. Consequently, the PF bond will be more polar than the PCl bond. You should compare the electronegativity differences for the two bonds to verify this prediction. The fluorine atom carries the partial negative charge.
Which of the following bonds is most polar: SCl, SBr, SeCl, or SeBr? Answer: SeCl
As we have just seen, there is a difference in electronegativity between H and F, which leads to a polar covalent bond in the HF molecule. As a consequence, there is a concentration of negative charge on the more electronegative F atom, leaving the less electronegative H atom at the positive end of the molecule. A molecule such as HF in which the centers of positive and negative charge do not coincide is said to be a polar molecule. Thus, we not only describe bonds as polar and nonpolar, but we also describe entire molecules this way as well.
We can indicate the polarity of the HF molecule in two ways:
Recall from the proceeding section that the "+" and "-" indicate the partial positive and negative charges on the H and F atoms. In the notation on the right, the arrow denotes the shift in electron density toward the fluorine atom. The crossed end of the arrow can be thought of as a plus sign that designates the positive end of the molecule.
Polar molecules align themselves with respect to each other and with respect to ions. The negative end of one molecule and the positive end of another attract each other. Polar molecules are likewise attracted to ions. The negative end of a polar molecule is attracted to a positive ion, and the positive end is attracted to a negative ion. These interactions help to explain the properties of liquids, solids, and solutions, as you will see in Chapters 11, 12, and 13.
How can we quantify the polarity of a molecule such as HF? Whenever two electrical charges of equal magnitude but opposite sign are separated by a distance, a dipole is established. The quantitative measure of the magnitude of a dipole is called its dipole moment, denoted . If two equal and opposite charges, Q+ and Q-, are separated by a distance r, the magnitude of the dipole moment is the product of Q and r (Figure 8.8):
Figure 8.8 When charges Q + and Q - are separated by a distance r, a dipole is produced. The size of the dipole is given by the dipole moment, , which is the product of the charge separated and the distance of separation between their centers: = Qr.
We see that the dipole moment will increase in size as the quantity of charge that is separated increases, and as the distance between the charges increases.
Dipole moments of molecules are usually reported in debyes (D), a unit that equals 3.34 10–30 coulomb-meters (C-m). For molecules, we usually measure charge in units of the electronic charge e, 1.60 10–19 C, and distance in units of Angstroms, Å. Suppose that two charges, 1+ and 1- (in units of e), are separated by a distance of 1.00 Å. The dipole moment produced is:
Measurement of the dipole moments of molecules can provide us with valuable information about the charge distributions in molecules, as illustrated in the following Sample Exercise.
The distance between the centers of the H and Cl atoms in the HCl molecule (called its bond length) is 1.27 Å. (a) Calculate the dipole moment, in D, that would result if the charges on the H and Cl atoms were 1+ and 1-, respectively. (b) The experimentally measured dipole moment of HCl(g) is 1.08 D. What magnitude of charge, in units of e, on the H and Cl atoms would lead to this dipole moment?
SOLUTION (a) The charge on each atom is the electronic charge, e: 1.60 10–19 C. The separation is 1.27 Å. By analogy to the calculation in the text above, the dipole moment is
Notice that the calculated dipole moment is greater than in the earlier example because the distance between the charges has increased from 1.00 Å to 1.27 Å.
(b) In this instance we know the value of , 1.08 D, and the value of r, 1.27 Å, and we want to calculate the value of Q:
We can readily convert this charge to units of e:
Thus, the experimental dipole moment indicates the following charge separation in the HCl molecule:
Because the experimental dipole moment is less than that calculated in part (a), the charges on the atoms are less than a full electronic charge. We could have anticipated this because the HCl bond is polar covalent rather than ionic.
The dipole moment of chlorine monofluoride, ClF(g), is 0.88 D. The bond length of the molecule is 1.63 Å. (a) Which atom is expected to have a negative charge? (b) What is the charge on that atom, in e? Answers: (a) F; (b) 0.11-
Table 8.3 presents the bond lengths and dipole moments of the hydrogen halides. Notice that as we proceed from HF to HI, the electronegativity difference decreases and the bond length increases. The first effect decreases the amount of charge separated and causes the dipole moment to decrease from HF to HI, even though the bond length is increasing. For these molecules, the change in the difference in electronegativity is a more important factor on the dipole moment than is the bond length.
We saw in Section 2.7 that there are two general approaches to naming binary compounds (compounds composed of two elements): one used for ionic compounds and the other for molecular ones. It is important to recognize now that in both approaches, the name of the less electronegative element is given first. The name of the more electronegative element then follows, modified to have an -ide ending. You will recall that compounds that are ionic are given names based on their component ions, including the charge of the cation if it is variable. Those that are molecular are named using the prefixes listed in Table 2.6 to indicate the number of atoms of each kind in the substance:
The dividing line between the two approaches, however, is not always clear, and both approaches are often applied to the same substances. For example, TiO2, which is a commercially important white paint pigment, is sometimes referred to as titanium(IV) oxide but is more commonly called titanium dioxide. The Roman numeral in the first name is the oxidation number of the titanium.
One reason for the overlap in the two approaches to nomenclature is that compounds of metals with higher oxidation numbers tend to be molecular rather than ionic. For example, SnCl4 [tin tetrachloride or tin(IV) chloride] is a colorless liquid that freezes at -33°C and boils at 114°C; Mn2O7 [dimanganese heptoxide or manganese(VII) oxide] is a green liquid that freezes at 5.9°C. Recall that ionic compounds are solids at room temperature. When we see the formula of a compound containing a metal with a high oxidation number (above +3), we should not be surprised that it does not exhibit the general properties of ionic compounds.