A number of eminent psychologists have expressed the wish to develop a “periodic table of behavior change methods” or a "periodic table of personality." I want to explore this idea in the next two blog posts. If we could create a sort of "periodic table of behavior," it might allow for more rapid progress in what has historically been regarded as a "soft" science. (My kids used to tell me "you're not a real scientist," by which they mean a physicist or chemist). As you will see below, a classification scheme for behavior would also serve as a marker of progress, an agreement on fundamental properties or underlying mechanisms that are suitable as subjects of psychological study. Having such agreement on our terms would prevent a great deal of re-work and confusion as people study similar concepts under different names. Chemistry used to have these same problems, however. Before we can explore what a "periodic table of behavior" might look like, I want to talk through some of the properties of the actual periodic table in chemistry, and spend a little while reflecting on what makes it so useful.
History of the Periodic Table
The Problem of Classification
Our situation in behavioral science is analogous to the chemists’ of the 19th century, who had identified many different elements based on their chemical properties and reactions, but were unable to specify the full list, were sometimes surprised by a fact about an element that didn’t seem to fit, and sometimes argued about whether a particular substance was actually an element or just some combination of other substances. Dalton’s early periodic table, for instance, included “lime,” “soda,” and “potash,” three compounds with widely-known industrial applications. The actual chemical identities of these substances are, respectively, calcium carbonate, sodium carbonate, and potassium carbonate. Calcium, sodium, and potassium had not been isolated yet, so in some ways Dalton was just confusing the compounds for the elements. But carbon did appear on Dalton’s table as a separate element, so there was clearly some confusion. Even worse, the generic term “potash” was also applied to other substances such as potassium hydroxide (“caustic potash”), potassium chloride, potassium sulfate, or potassium nitrate (also called “saltpeter”). Potassium is part of many compounds because it so readily combined with other elements, and these compounds were initially mistaken for elements themselves.
Related Groupings of Elements
Beyond just classifying elements, steps toward chemistry’s periodic table were taken by alchemists in the Middle Ages who observed similarities in certain substances’ chemical properties. Calcium, barium, and magnesium, for instance, are all “alkaline earth metals” (Group 2 of the table), and share properties such as shiny whiteness and an ability to react with water. The “alkali earth metals” (Group 1) like potassium, sodium, and cesium have even more violent reactions with water, glowing as they emit hydrogen gas; they easily combine with other minerals to form salts (e.g., table salt, NaCl), and often have an effect on biological processes (e.g., neuroelectrical activity is only possible because of the movement of these minerals’ ions across the cell membrane in a “sodium-potassium pump”). Chlorine, bromine, and iodine — halogen elements found in Group 7 — were also known to share important properties. In the modern periodic table, each of these groupings can be seen as a column of elements: Reading down each column defines a chemical group, and knowing something about the behavior of one element in a group allows one to make predictions about the ways that other elements in the same grouping will react with substances or not — their affinity to combine or “valency” is their common characteristic.
Dmitri Mendeleev was the Russian chemist who in 1869 formalized these known groupings into a table that also ordered the elements by their atomic weight. These weights were newly discovered at the time, by chemist John Dalton who weighed samples of different gasses at constant volume, pressure, and temperature (using the universal gas law, pressure x volume = amount x the gas constant x temperature, or PV=nRT, plus the newly discovered "Avogadro number" of atoms in any given amount "n" of an element, and expressing the result as a proportion relative to hydrogen atoms). The idea of ascending atomic weights, with mostly-regular intervals between the columns or “periods” of the table, allowed Mendeleev to clear up some confusion (e.g., about the true atomic weight of beryllium, which had been in doubt). Even more convincingly, it allowed Mendeleev to predict the existence and properties of a yet-undiscovered element, gallium, which was first isolated from nature in 1871. Mendeleev’s predictions for this new element were spot-on, including its atomic weight, its ability to react with other elements, and some of the specific compounds that could be produced. Two other elements - scandium and germanium - were similarly predicted and later discovered. And when a whole new grouping of elements, the inert or “noble” gasses (Group 8), was discovered a few years later, it was immediately apparent where this grouping belonged at the right edge of Mendeleev’s table.
Underlying Mechanisms that Produce the Groupings
Impressive as the periodic table was, its modern form required three further developments: The first was provided in 1913 by British physicist Harry Moseley. Mosely was a student in the lab of Ernest Rutherford at Cambridge, who had first proposed the familiar “solar system” model of the atom that is mostly empty space. Moseley used cathode rays to bombard elements that he ran through a vacuum tube on a little train, and observed the different x-ray frequencies that they emitted in response. This allowed him to isolate the exact positive charge in each atomic nucleus - the number of protons - and to differentiate this from the atomic mass that also included neutrons and was watered down by the existence of various isotopes (atoms with different numbers of neutrons) of each element in nature. The measurement of elements’ atomic numbers confirmed Mendeleev’s order, cleared up a few areas where the atomic numbers ascended even though the atomic weights went down, and most importantly identified exactly seven missing elements in the table along with their exact atomic numbers. Another of Rutherford’s students, Niels Bohr, added to this a concept of electron “shells” based on the quantum theory of physics formulated by Max Planck in 1900. In Bohr’s view, electrons could occupy only discrete states around the atomic nucleus — all in or all out — and a quantum jump from one state to another required the input or release of a discrete packet of energy. This notion explained why there were only 2 elements in the first row of the periodic table (2 positions available for electrons in the first shell), 8 in the next two, 18 in the fourth and fifth rows, and 32 after that. It explained why the noble gasses were inert (their outer electron shells were full), and why elements toward the edges of the table were so eager to gain or lose an electron — the consistent behavior of elements in the halogen, alkali, and alkaline groups that had been known to the alchemists. And the behavior of electrons explained further phenomena like the conductivity of metals (their electrons moved freely across atoms); characteristic patterns of light that could be used to identify each element via spectrography (based on the specific wavelengths of light that an element could absorb to jump its electrons to the next shell); phosphorescence (based on energy being used to excite electrons to a higher shell and then released as they gradually fell back to their original shell); and even the shininess or hardness of some minerals (based on the amount of energy binding their electrons to the nucleus). Mendeleev’s classification scheme specified which elements should be listed on what order; Moseley and Bohr explained why. Again, an advance in theory allowed for a forward leap in prediction: Bohr was able to predict that element 72, hafnium, would be a heavier analogue of zirconium, and he told colleagues on Denmark to look for it in zirconium ores, where they very quickly found it. As an aside, Bohr's notion of electron shells also led to predictions about the geometric shape of atoms and by extension molecules, which turned out to be of crucial importance in understanding how a given chemical interacts with the body -- as in the example of the "spike protein" that allows the SARS-CoV-2 virus to bond with human cells via the ACE2 receptor and convert them into virus factories resulting in COVID-19 illness.
Inter-Relationships between One Element and Another
The final missing ingredient of the puzzle is a hidden structure of the periodic table that was provided by Pierre and Marie Curie, who discovered the element radium in 1898 but took years to understand it. What the Curies found most intriguing about radium was its ability to generate energy in the form of heat and light, which continued even when the element was isolated from outside sources that might replenish the energy supply (unlike, for example, phosphorescent minerals that would gradually stop glowing when removed from a light source). What gradually became clear in this research was that radium was a point in a radioactive "decay chain" of elements that started with uranium. Any given atom of uranium would eventually give off a helium nucleus (also called an "alpha particle") and decay into thorium -- at a glacially slow rate of about half the uranium atoms decaying over 245 thousand years. Thorium in turn would give off an alpha particle and decay into Curie's radium, with a half-life of 75 thousand years. And radium was even more strongly radioactive, with a half-life of just 1,600 years, which led to its more obvious physical effects. After radium, the chain goes through radon, polonium, and finally reaches a stable state when the element changes into lead. Each step of this decay chain moves left by two positions on the periodic table, and works by removing a helium atom, the element with an atomic number of 2. The existence of radioactive decay chains, and their regular steps down the list of elements, provides further evidence that the periodic table is not just a handy picture, but actually an illustration of the hidden structure of the universe. Understanding this also provided the basis for understanding how elements could be formed from the basic building block of hydrogen atoms, including the natural creation of all the elements up to iron in the hearts of stars, the production of heavier elements in supernova explosions, and the artificial creation of elements that do not exist on Earth by human scientists. Marie Curie eventually received Nobel prizes in both chemistry and physics for her work, and died from complications of radiation exposure.
The addition of nuclear fission (radiation) and nuclear fusion to our understanding of the periodic table actually brings this organizing scheme back to one of the ancient goals of alchemy, which was the transmutation of one element into another. By splitting atoms, scientists have actually succeeded in changing lead into gold (it's possible with particle accelerators, just not cost-effective)!
Properties Needed for a Table of Behavior
To return to the place we began, let's examine the fundamental properties of the periodic table that might serve as models for a periodic table of behavior. Some people looking at this question have tried to figure out what the "rows" and "columns" in a periodic table of behavior should be, but the shape of chemistry's periodic table is not fundamental. It could, for instance, be drawn like this with equal validity: Elements in this version march out from the center in concentric rings, with the families of elements along each radial arm and special, extra-long protrusions for the transition metals, the lanthanide/actinide series, and a further hypothetical electron shell that would show up starting in the 8th row of the table.
1. Classification - at the beginning of the modern era of chemistry, there was not yet any agreement about what "counted" as an element. Some compounds had historically be considered elements, and some elements had two or more names. To make any progress at all, chemists had to agree on the list of known elements; the new technology for establishing atomic weights helped to standardize that list.
2. Periodicity - in other words, regular groupings of things that have common properties and characteristics. Knowing something about one element of any of these groups allows you to make predictions and inferences about other elements in the same group. There is room for some level of individuality in the elements of a group -- for instance, Oliver Sacks points out that the "chalcogens" in Group 6 of the periodic table are usually noxious-smelling gasses (sulfur, selenium, tellurium), with the exception of the top element in the column which is pure oxygen!
3. Underlying Mechanisms - the reason for periodicity will be an application of some underlying property or law. For the periodic table this is provided by the concept of elements' valence -- that's what makes them reactive or non-reactive. The underlying mechanisms should be stable and invariant, even if the properties of the individual elements in a grouping vary somewhat.
4. Inter-Relationships - elements in the table can't be a simple list; they have to be related to each other based on underlying structures of some kind. In the case of the periodic table the underlying structures are the number of protons (atomic number), the number of electrons in ordered shells (which produces an atom's shape and charge), and the stability or instability of the atom (leading to radioactivity and the transition of one element into another). An element's position on the periodic table suggests not just its properties based on column grouping, but also its linkages to other elements based on their positions relative to one another.
Stay tuned in two weeks: My next blog post will examine how we are doing on each of these properties in behavioral science, and some developments that suggest we are finally moving in the right direction.
Thanks this week to my colleague Kai Larsen at CU Boulder who first introduced me to the idea of a "periodic table of behavior." Thanks also to Oliver Sacks's enjoyable book "Uncle Tungsten: Memories of a Chemical Boyhood" for the history lesson on chemistry, and to Mr. Butler who taught chemistry at Rush-Henrietta Senior High School in the 1980s. At the time I didn't properly appreciate just how fun all this science can be!
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