0.9 Energy and polarity of covalent chemical bonds  (Page 8/8)

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 H ₂ O and H ₂ S. H ₂ O has a dipole moment with the O being the negative end. This means that O is more electronegative than H. H ₂ S also has a dipole moment, but it is much smaller than that of H ₂ O, so S is less electronegative than O. Let’s compare the dipole moments of NH ₃ and PH ₃ : we see the same trend. It is generally true that electronegativity is larger for smaller atoms.

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 NH ₃ <H ₂ O<HF, and PH ₃ <H ₂ S<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 defined physical property. In fact, there are several possible definitions of how to measure electronegativity each of which leads to slightly different values for each atom Although the exact values vary, the overall trends we observed are similar. Using one popular definition, [link] shows electronegativities 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 find 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 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 affinity 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 find out, we should examine other elements with high electronegativities. The highest electronegativities 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 affinity whereas O has a strong electron affinity, so electron affinity must also be important in understanding electronegativity. 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 affinity 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 affinities have higher electronegativity. The reasons for high ionization energy and high electron affinity 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 affinity, and larger electronegativity. One way to define electronegativity is simply as the average of the ionization energy and the electron affinity. In fact, the values in [link] are just this average multiplied by a constant to give a simplified 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.

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 affect be measured experimentally?
2. The bond in an H ₂ molecule is almost twice as strong as the bond in the H ₂ ⁺ ion. Explain why the H ₂ bond is so much stronger. Why isn’t the H ₂ bond exactly twice as strong as the H ₂ ⁺ bond?
3. The ionization energy of H ₂ is slightly less than the ionization energy of H ₂ ⁺ . But the bond energy of H ₂ is much larger than the bond energy of H ₂ ⁺ . Explain how these two facts are consistent with each other.
4. In this study, we referred to H ₂ as a “stable” molecule. But H ₂ gas can be explosively reactive, as a viewing of the Hindenberg disaster clearly reveals. In what sense is H ₂ “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 affinity 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 affinity, meaning that N atoms do not attract electrons. Explain how and why these facts are not inconsistent.