Professor Brown describes bonds in terms of their electrostatic flux. This is not the same as

Water molecule, (A) isolated, (B) in condensed phase showing atomic and bond valences.
Electrostatic flux lines
Professor Brown describes bonds in terms of their electrostatic flux. This is not the same as trying to work out where the negative charge in a bond is most likely to be found since the flux can be calculated from the valence characteristics of atoms and number of bonds they form as mentioned previously. Large bond fluxes correspond to shorter and stronger bonds and, for molecules where the bond lengths have been experimentally measured, the flux can also be back calculated from the bond length.
Often what chemists want to know is the answer to the question ‘will they bond’?
Balancing acts

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Another application of modern bond valence theory that Professor Brown has highlighted is the ability to predict the structure and shapes of molecules. Sometimes molecules which contain exactly the same atoms may have different structures. While some of these additional structures may also be stable, often what is of interest is the most stable, or equilibrium, form of the molecule.
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Modern bond valence theory is a simple and efficient tool that allows for quick and intuitive insights.
Through Professor Brown’s models, the equilibrium structure can be identified by finding the combination of linked atoms in which all of the fluxes of the bonds are most nearly equal. This then gives rise to the lowest energy and most likely structure of the molecule. The bond flux picture describes not just ‘traditional’ chemical bonds but all types of chemical bond between atoms, including ionic bonds and the more exotic hydrogen bonds. Hydrogen bonds are an unusual type of bond as they are very long-range interactions, so molecules capable of hydrogen bonding with each other will interact over large distances. While distances of up to 3 Angstroms might seem insignificant to us, these are huge on a molecular scale and over two times the length of a more standard chemical bond. The most famous example of a hydrogen bonded system is water and it is these unusual bonding interactions that cause water’s sometimes unintuitive behaviour, such as expanding when it freezes (nearly all other liquids will contract).
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Professor Brown describes bonds in terms of their electrostatic flux.
General models
A large amount of the success of modern bond valence theory comes from its ease of use and general applicability. Professor Brown has shown that the model works well for inorganic compounds, which typically are difficult to describe with both quantum and older classic models. To help develop and test the model, he also helped established the Crystallographic Information Framework (https://www.iucr.org/resources/cif/software) and the Inorganic Crystal Structure Database so these models can continue to be refined and extended for future use.
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Personal Response
What are some of the key predictions or insights that have come from modern bond valence theory?
The model predicts that bonds will only be formed between atoms of similar bonding strength; strongly bonding atoms do not bond to weakly bonding atoms. This important property allows one to generate plausible bond networks.
The model shows that the unusual geometries found in hydrogen bonds arise from the sizes and valences of the atoms involved. They are not electronic in origin as often supposed.
The model can predict when the bonds will be strained, and whether these strains will result in unusual physical properties such as ferroelectricity.
The simple description of the structure in terms of atoms linked by bonds can be used to model the structures found in liquids, solutions, and surfaces, as well as in solids, and from these structures one can explore chemical and biological properties, such as solubility in water and the mechanisms of catalysis, and physical properties, such as ionic conduction in the solids used in modern batteries.
Article References
Brown, I. D. (2000). The Bond Valence Model as a Tool for Teaching Inorganic Chemistry: The Ionic Model Revisited. Journal of Chemical Education, 77(8), 1070–1075. https://doi.org/10.1021/ed077p1070
Brown, I. D. (2019). Another look at bonds and bonding. Structural Chemistry, 31(1) 1–5. https://doi.org/10.1007/s11224-019-01433-7
Brown, I. D. (2009). Recent developments in the methods and applications of the bond valence model. Chem. Rev. 109, 6858–6919. http://dx.doi.org/10.1021/cr900053k
Behind the Research
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Prof David Brown
David Brown is Professor Emeritus at McMaster University, Ontario, Canada. He has shown that the theory of chemical structure is classical and does not require quantum mechanics. To assist in this work he helped found the Inorganic Crystal Structure Database and develop the Crystallographic Information Framework (CIF): https://www.iucr.org/resources/cif/software