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We now have an understanding of what it means for two atoms to share an electron pair and why this results in a bond. In addition to H2, this description works well in describing and understanding chemical bonds such as in F2 or Cl2. More work is required to understand multiple bonds such as the double bond in O2, but it turns out that the same principles apply. What about molecules where the atoms are not the same, e.g. HF? The Lewis model of these molecules still assumes a sharing of an electron pair. But we learned in our study of atomic structure that the properties of H atoms and F atoms are quite dierent. For example, the ionization energy of an F atom is larger than the ionization energy of an H atom. F atoms also have very strong electron anity. Do such dierent types of atoms share electrons? If so, does it matter that the properties of the two atoms are so dierent?

We need observations to answer these questions. First, we can do as we did with H₂ and examine the energies of the bonds between dierent atoms. The bond energy of HF is 568 kJ/mol, larger than the bond energy of H₂. By contrast, the bond energy of F₂ is 154 kJ/mol, quite a bit weaker than the bonds in either HF or H₂. The bond energy of HCl is 432 kJ/mol, weaker than the HF bond. The energy of one O-H bond in H2O is 463 kJ/mol. Clearly, the strength of bond depends on what types of atoms are bonded together. This suggests that the sharing of an electron pair depends on the properties of the atoms in the bond, including the sizes of the atoms and the charges on the nuclei. In fact, it would not be too surprising if dierent atoms did not share the electrons equally.

To nd out if this is the case, we observe a property of molecules called the dipole moment. An electric dipole is simply a separation of a positive and a negative charge. The dipole moment, usually labeled µ, measures how strong the dipole is by taking the product of the amount of the charge times the separation of the charge. Just having a positive charge, let's say a proton, and a negative charge, let's say an electron, does not mean that there is a dipole moment. The hydrogen atom does not have a dipole moment because the electron is in constant rapid motion around the positive nucleus. As such, viewed from outside the atom, there is no positive end or negative end to the hydrogen atom, so there is no dipole moment. This is true of all atoms.

It might seem that molecules could not have dipole moments for the same reason. The electrons are moving about the nuclei rapidly. In the H₂molecule, although there are positive and negative charges, there is no end of the molecule which looks more positive and no end which looks more negative. H2 does not have a dipole moment.

What about a molecule like HF? If an H atom has no dipole moment and an F atom has no dipole moment, does HF have a dipole moment? Experimentally, molecular dipoles can be observed in a number of ways. If you put a molecule with a dipole moment in an electric eld, it will line up with the eld, with

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92 CHAPTER 9. ENERGY AND POLARITY OF COVALENT CHEMICAL BONDS the positive end of the dipole pointed towards the negative end of the eld. Dipoles can interact with each other as well, so that the negative end of one dipole will point towards the positive end of another dipole.

It is also possible to measure the magnitude of the dipole moment. A number of these are given in Table 9.1: Dipole Moments of Specic Molecules for some simple molecules. We can see that, not only does HF have a dipole moment, but in fact it has the largest dipole moment of the molecules listed. What does it mean that HF has a dipole moment? An HF molecule must have a permanent negative end and a permanent positive end. Experimentation shows that the F end of the molecule has a net negative charge and the H end has a net positive charge.

Dipole Moments of Specic Molecules

Molecule µ (debye)

H2O 1.85

HF 1.91

HCl 1.08

HBr 0.80

HI 0.42

CO 0.12

CO₂ 0

NH₃ 1.47

PH3 0.58

AsH3 0.20

CH4 0

NaCl 9.00

Table 9.1

How can this be so? The total number of electrons and protons in the molecule are evenly matched, of course, and the electrons are moving rapidly about the two nuclei just as in H₂. Perhaps the electrons are not moving uniformly around the H and the F nuclei. To observe this, we look at the molecular orbital for the shared electrons in the HF bond, shown in Figure 9.3.

Figure 9.3: Distribution of Electron Probability in HF

93 We observe that the electrons in the molecule move with greater probability near the F atom. There is thus more electron charge around the F nucleus than the positive charge on the F nucleus. And the opposite is true for the H end of the molecule. Apparently, when an H atom and an F atom share electrons, they do not share them equally!

This is an extremely important result, one of the most useful in all of Chemistry. Dierent atoms have dierent tendencies to draw electrons to themselves when sharing electrons in a covalent bond. The relative strength with which an atom draws electrons to itself in a bond is called electronegativity. A very electronegative atom will attract the shared electrons more strongly than a less electronegative atom. This will produce a negative charge near the more electronegative atom and a positive charge near the atom other.

This will, in most cases, create a permanent dipole moment in the molecule. At this point, we don't know how much negative charge is near the more electronegative atom, but it probably is not the full charge on a single electron since the electrons are still mostly shared by the two atoms. As such, we label the negative charge by δ-, where δ is some number between 0 and 1. In most cases, we don't even need to know how large δis. The positive end of the molecule is labeled by δ+, since whatever the negative charge on the negative end must equal the positive charge on the positive end. HF is a good starting example. The F atom is more electronegative than the H atom; hence, the HF molecule has a dipole moment, which makes HF what we call a polar molecule.

Figure 9.4

Of course, we next want to know why the F atom is more electronegative. We need more data to develop a model for electronegativity. Table 9.1: Dipole Moments of Specic Molecules listing the dipole moments of several molecules provides good data for comparison.

Let's rst study the set of molecules HF, HCl, HBr, and HI. It is easy to see that the dipole moments of these molecules are in the order of HF > HCl > HBr > HI. To analyze this comparison, we need to remember that the dipole moment measures the product of the amount of charge separated times the distance by which the charges are separated. Since the dipole moment of HF is larger than that of HCl, let's compare these two. How do the properties of F atoms compare to those of Cl atoms? One thing we can say for sure: Cl is a larger atom than F. This means that the HCl bond (127 pm) is longer than the HF bond (92 pm). Since the dipole moment depends on the distance between the positive and negative charges, the comparison of bond lengths would perhaps lead us to predict that HCl would have a larger dipole moment than HF, But that's not what the data tell us. There is only one way to explain this. The dierence between HF and HCl must be in the amount of charge that is separated. It must be true that the negative charge on the F atom is greater than the negative charge on the Cl atom. This means that F attracts the shared electrons in the H-F bond more than the Cl atom attracts the shared electrons in the H-Cl bond. This means that F is more electronegative than Cl.

Looking at the dipole moments of these four molecules and remembering that the atoms get larger in the order F < Cl < Br < I, it must be true that the electronegativities go in the order F > Cl > Br > I. In this group, the electronegativity is larger for smaller atoms. To see if this turns out to be generally true, we can examine other families in the periodic table.

First, let's look at the dipole moments for H2O and H2S. H2O has a dipole moment with the O being the negative end. This means that O is more electronegative than H. H2S also has a dipole moment, but it is much smaller than that of H2O, so S is less electronegative than O. Let's compare the dipole moments of NH3 and PH3: we see the same trend. It is generally true that electronegativity is larger for smaller atoms.

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94 CHAPTER 9. ENERGY AND POLARITY OF COVALENT CHEMICAL BONDS We should also compare the electronegativities of elements in the same row of the Periodic Table. For example, the dipole moments increase in the order of NH3<H2O < HF, and PH3<H2S < HCl. From this, we can conclude that, in general, electronegativity increases with increasing atomic number within a single row of the Periodic Table.

Electronegativity is an extremely useful concept in chemistry, but it is not a precisely dened physical property. In fact, there are several possible denitions of how to measure electronegativity each of which leads to slightly dierent values for each atom Although the exact values vary, the overall trends we observed are similar. Using one popular denition, Table 9.2: Electronegativity of Selected Atoms shows electronega-tivities for many atoms. Looking at these numbers, you should be able to see the trends we have developed from analyzing dipole moments of simple molecules.

With these observations in mind, we need to develop a model to understand why the electronegativity is larger for a larger atomic number in a single row, but is smaller for a larger atomic number in a single group. These trends seem to contradict each other. There must be a good physical explanation of these observations.

To nd one, let's note that both trends point to F being the most electronegative element. F is as far to the right on the table as we can go and as far to the top of the table as we can go. (We might consider He or Ne to be more electronegative, but since there are no known molecules containing bonds with He or

95 Ne, electronegativity has no meaning for these atoms.) What else do we know about F atoms? Thinking back on our understanding of atomic energy levels, we recall that F atoms have the highest ionization energy (except for He and Ne) and the highest electron anity of any of the elements. Electrons are clearly most strongly attracted to F atoms. We developed a model to explain this based on Coulomb's law. The F atom uniquely combines the largest core charge and the smallest distance of the valence electrons to the nucleus (measured as the average orbital distance). Perhaps these two factors are also responsible for F having the largest electronegativity.

To nd out, we should examine other elements with high electronegativities. The highest electronegativ-ities are all for elements with high ionization energies. Interesting examples include N, O, and Cl. Of these, O has the highest electronegativity, and this makes sense: it has a large core charge and a small shell radius.

But if we compare N and O, N has the higher ionization energy. There must be more to electronegativity than just ionization energy. Another interesting comparison of N and O is that N has no electron anity whereas O has a strong electron anity, so electron anity must also be important in understanding elec-tronegativity. This makes sense: an atom with a higher ionization energy is less likely to have its electrons drawn to another atom in a chemical bond, and an atom with a higher electron anity is more likely to draw electrons from another atom in a chemical bond.

This produces a simple model to understand electronegativity. Atoms with higher ionization energies and higher electron anities have higher electronegativity. The reasons for high ionization energy and high electron anity are the same as the reasons for high electronegativity. On the basis of Coulomb's law, a larger core charge and a smaller shell radius generally give larger ionization energy, larger electron anity, and larger electronegativity. One way to dene electronegativity is simply as the average of the ionization energy and the electron anity. In fact, the values in Table 9.2: Electronegativity of Selected Atoms are just this average multiplied by a constant to give a simplied scale.

Understanding these trends is extremely useful. Electronegativity is one of the most powerful concepts in chemistry for predicting chemical reactivity. For example, positive ends of molecules are often attracted to the negative ends of other molecules. Understanding where there may be a more negative charge in a molecule can then help us predict the location in a molecule where a reaction may take place or even predict whether a reaction is expected to occur or not. We will have many occasions to apply the concept of electronegativity, including in the next concept study.

9.6 Review and Discussion Questions

1. Why does an electron shared by two nuclei have a lower potential energy than an electron on a single atom? Why does an electron shared by two nuclei have a lower kinetic energy than an electron on a single atom? How does this sharing result in a stable molecule? How can this aect be measured experimentally?

2. The bond in an H2 molecule is almost twice as strong as the bond in the H2+ ion. Explain why the H2bond is so much stronger. Why isn't the H2bond exactly twice as strong as the H2+ bond?

3. The ionization energy of H2 is slightly less than the ionization energy of H2+. But the bond energy of H₂is much larger than the bond energy of H₂⁺. Explain how these two facts are consistent with each other.

4. In this study, we referred to H₂ as a stable molecule. But H₂ gas can be explosively reactive, as a viewing of the Hindenberg disaster clearly reveals. In what sense is H₂ stable? How can a stable molecule be a highly reactive molecule?

5. Explain why an atom with a high ionization energy is expected to have a high electronegativity.

6. Explain why an atom with a high electron anity is expected to have a high electronegativity.

7. Explain why S has a greater electronegativity than P but a smaller electronegativity than O.

8. N atoms have a high electronegativity. However, N atoms have no electron anity, meaning that N atoms do not attract electrons. Explain how and why these facts are not inconsistent.

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96 CHAPTER 9. ENERGY AND POLARITY OF COVALENT CHEMICAL BONDS

Chapter 10

Bonding in Metals and Metal-Non-Metal