Assumptions: Part 2 — Fact versus theory: electronegativity

Assumptions: Part 2 — Fact versus theory: electronegativity

​In the first article of the series,¹ I discussed the assumptions behind the preparation of solutions; i.e., why are volumes not always additive? In the second article of the series, I would like to discuss the confusion that arises in the minds of some of our students between fact and theory, with particular attention to the concept of electronegativity.

Many beginning chemistry students believe that nearly everything they read in their textbook is completely factual. How many instructors, when collecting papers at the end of an exam, have found periodic tables with a collection of vertical and horizontal arrows drawn by their students, depicting the trends for atomic size, ionization energy, electron affinity and electronegativity? The students have “memorized” these trends and have quickly drawn the arrows on their copy of the periodic table at the beginning of the exam. What many of them do not realize is that these are merely “trends”, and they hope that these trends will be sufficient to arrive at the correct answer to a given multiple choice question.

For many years, I started off the first class of a general chemistry course by writing the words FACT and THEORY on opposite ends of the blackboard/whiteboard in the classroom. I impressed upon my students that one must have some facts before one can hope to formulate a theory to be able to predict the behaviour of a different substance in a similar situation or the same substance in a different situation. Far too often, the textbook coverage gives students a simple neat “picture” for a given periodic trend. For this portion of the course, I would constantly refer to the two words on the blackboard/whiteboard and insisted that the students make the distinction between the two, and realize that one must have experimental facts first.

One particular area of confusion involves the concept of electronegativity. Far too many students “believe” it to be determined experimentally, rather than being a theoretical construct. Most textbooks will include a graph of electronegativity values as a function of atomic number, along with statements such as²:

Electronegativity generally increases across a period in the periodic table.

Electronegativity generally decreases down a column in the periodic table

The key word in these statements is “generally”, as the electronegativity of an atom can vary from compound to compound. For example,³ the value for carbon in CH₃Cl is 2.4, while in CCl₄ it is 2.6. A consideration of electronegativity values is also often used to classify bonds as being purely covalent, polar covalent or ionic.

Students need to be reminded that this classification scheme is qualitative not quantitative, as electronegativity values cannot be measured directly. They are calculated in a variety of ways (see reference 6). One physical property that can be measured directly is the dipole moment of a compound (fact). One may also calculate the theoretical value of the dipole moment for a diatomic substance (theory). The percent ionic character is the ratio of the measured dipole moment of a compound compared to the calculated value, assuming that the electron was completely transferred, i.e., 100% ionic. The percent ionic character for a series of diatomic molecules in the gas phase is often plotted against the difference in electronegativity values of the two atoms. The results of such an analysis² gives 20% ionic character for HCl, 45% for HF, 60% for LiBr and 75% for NaCl. It is important for students to appreciate that by using this analysis there are no truly 100% ionic compounds. Most textbooks propose that ionic compounds are those that have greater than 50% ionic character; i.e., ΔEN greater than 2.0).

At this point, students are likely quite concerned as to where this property known as electronegativity actually came from. To help ease the confusion, I offer the following exercise for the students to complete.

The bond energies of the pure diatomic elements, in the gas phase, can be determined experimentally. The energy required is for the homolytic cleavage of the “covalent bond” in X₂ to form two free radicals (X.).

The experimentally determined values for this endothermic process are⁴:

You can now ask your students to calculate the expected bond energy for the series of HX compounds. The expected bond energy should be equal to one-half of the sum of the bond energies of H₂ and X₂.

The results of these calculations, as well as the actual experimentally determined values, are shown below:

The students will likely be puzzled as to why there is a difference between the theoretical and actual bond energy values. You can now ask them to consider what assumption they must have unconsciously made in calculating the theoretical values. For the pure elements, since each atom is the same, one must assume that the covalent bond between them is a result of equal sharing of the bonding electrons, i.e., ΔEN = 0. This means that all five diatomic molecules must be nonpolar. For the series of the HX compounds, since the bond energies are “larger” than expected, the covalent bonds must be “stronger” than expected. This implies that there must be unequal sharing of the bonding electrons, and that the bonding electrons must be closer to one atom than the other. This unequal sharing of the bonding electrons can be related to the dipole moment of each of the four compounds (fact).⁵

Since the magnitude of the difference between the predicted bond energy (theory), and the actual bond energy (fact) is HI < HBr < HCl < HF, the F atom must therefore have a greater attraction for the bonding electrons, relative to H, than the other halogen atoms. Also, HF has the largest dipole moment of the four compounds (fact).

One can then suggest to the students that the concept of electronegativity is simply a relative scale for the unequal sharing of bonding electrons in a covalent bond between two atoms in a covalent compound. For more on the concept of electronegativity, you could have your students read through this descriptive account on the web⁶, and then have them come back to the next class ready for further discussion.

References

1. R.R. Perkins, Assumptions: Part 1 — Preparing solutions, Chem 13 News, November 2017.
2. N. Tro, T. Fridgen and L. Shaw, Chemistry – A Molecular Approach, 2014, pages 345-349.
3. J. Umland and J. Bellama, General Chemistry (2nd edition), 1996, pages 308-310.
4. www.science.uwaterloo.ca/~cchieh/cact/c120/ bondel.html
5. https://chem.libretexts.org/Core/Physical_httpsand_Theoretical_Chemistry/Physical_Properties_of_Matter/Atomic_and_Molecular_Properties/Dipole_Moments
6. https://en.wikipedia.org/wiki/Electronegativity