Periodic Pandemonium: Exploring Chemistry’s Rule-Breakers

When I tell people that I study chemistry, the response is usually some version of “You must be bright” or “I hated that class” or, put more simply, “Why?”  I’ve grown used to defending my love for chemistry, and I’ve often pointed to its straightforward nature as the source of my affection. I liked that the elements on the periodic table are arranged according to trends in their chemical properties, and that we can infer things about an element’s behavior by its position.   An element’s electronegativity (the tendency to attract electrons), for example, increases as you move from left to right and from bottom to top across the table. The size of the atom, meanwhile, increases from right to left and from top to bottom. In class, I labeled different chunks of the periodic table to signify alkali metals, transition metals, lanthanides, actinides, halogens and other categories of elements with neatly-defined criteria for membership.  As a straight-A student and a type-A personality, I appreciated how chemistry was so orderly, how there was always a right answer.   

But as I’ve continued my studies throughout college and graduate school, I’ve realized that the lines on the periodic table are blurrier than I once believed them to be.  When I made a statement about an element that I had learned by heart, my classmates would sometimes respond, “Well, that’s not true 100% of the time.” With a little bit of digging, I came across scientific research papers reporting chemical reactions that would have, in fact, earned me a big red “X” if I had written them on an exam in high school. Eventually, I learned that there is a growing collection of instances in which elements break the rules of the periodic table.    

Both external and internal factors can bring about unusual elemental behavior.  Externally, exposing an element to extreme environments can alter its chemical properties.  For example, we know that we can melt gold by heating it to very high temperatures, and we know that diamond, the hardest material on Earth, comes from applying extreme pressure to carbon, the same element that coal and graphite are made of.  In the same way, by varying the temperature and pressure of a reaction, scientists can change the way atoms typically interact. Using this strategy, researchers have observed chemical phenomena that they never would have expected based on the element’s position on the periodic table.   

One of these phenomena relates to the noble gases, which are the elements that make up the right-most column of the periodic table.  At room temperature and normal pressure conditions, noble gases do not participate in chemical reactions. This is because of the way their electrons are arranged inside the atom.

Imagine filling up a spice rack with bottles of spices.  If you have one empty slot, you will probably go out in search of another spice to fill it.  If you have one spice too many, you may convince yourself that you don’t really use cumin that much anyway and give it to a friend.  Atoms similarly pursue the satisfaction of having the exact number of electrons required to fill their “slots,” known as orbitals. Noble gases have perfectly full orbitals, so they don’t see the need to exchange electrons with other atoms. 

Although unreactive, noble gases are far from unimportant.  In fact, two of them—xenon and argon—are at the root of one of the greatest mysteries in the geosciences field.  The level of xenon on Earth is much lower than that on meteorites made up of the same raw materials. Similarly, the total amount of argon in the Earth’s crust, mantle, and atmosphere does not account for all the argon that is predicted to be on our planet.  In other words, there should be a lot more xenon and argon on Earth than we’ve observed.  

Scientists have long speculated that there is another source of xenon and argon somewhere on Earth, but that they just haven’t been able to find it.  Recently, a group of researchers delved into the case of the missing xenon and argon, searching for a clue that could crack it. They began to wonder if noble gases could, in fact, react with other atoms when exposed to the extreme environment of the Earth’s core and hide from us in the form of a yet-to-be discovered compound.  To find out, the group conducted experiments on argon and xenon under conditions that modeled those at the Earth’s core.

To achieve such high pressure and temperature, they crushed the atoms between two hard surfaces with a huge amount of force and fried them with a laser. When they did this, the researchers discovered that, remarkably, xenon and argon can react with metals such as iron and nickel.  And since the Earth contains major stores of iron and nickel, it is possible that the reaction they observed in the lab could also happen at the Earth’s core.  Therefore, a significant amount of xenon and argon might be located inside the core of the Earth as a metal compound. It seems we can no longer say that the noble gases ar-gon! (Get it?)

While an extreme external environment revealed a rebellious streak in the noble gases, one internal factor that can cause an element to display unexpected behavior is if it has too many protons.  Protons are subatomic particles just like electrons, but they are positively-charged, and the number of protons an atom has dictates what type of element it is.  For example, hydrogen atoms have one proton and plutonium atoms have 94 protons. Hydrogen and plutonium and all the elements that fall in between are found in nature, but the remaining 24 elements on the periodic table are not.  This is because scientists have added protons to the plutonium atom to create new synthetic elements. Containing so many protons causes the atom to weigh a lot (for an atom, anyway), so these elements are referred to as “superheavy.”  The large number of protons also produces a strong positive charge at the center of the atom, which can attract the orbiting, negatively-charged electrons and throw them off their path. Since electrons control atoms’ ability to interact, this effect—known as the relativistic effect—can impact the superheavy elements’ chemical properties. 

One of these properties is called “standard enthalpy of formation,” which refers to the amount of energy needed to form a certain amount of an element.  Copernicium (which has 112 protons) and flevorium (which has 114 protons) have standard enthalpies of formation that do not follow the trend established by the preceding elements in their column.  Additionally, flevorium behaves similarly to mercury, which is in a different region of the periodic table altogether.  Also, oganesson—the element most recently added to the periodic table and the heaviest element yet synthesized at 118 protons—is naturally quite reactive, despite falling into the “noble gas” column.  

Unfortunately, our understanding of the superheavy elements is limited by their instability.  Just as suitcases have a finite capacity, you can only fit so many protons into an atom, and the superheavy elements are approaching packing-for-a-month-long-international-getaway-in-a-carry-on territory.  Therefore, like a suitcase busting at the seams, superheavy elements will expel some of their contents, or decay, to ease the congestion. Because they decay so quickly, scientists haven’t been able to study the superheavies as thoroughly as other elements, and it’s unlikely that they can be used in any practical applications.  But the abnormal behavior observed in initial experiments serves as a reminder that we don’t already have it all figured out and that there is always more exciting chemistry to be explored.

While this reality would have been unsettling to me at the start of my journey, I’ve evolved as a scientist, and I now appreciate the fact that understanding chemistry is more than simply memorizing facts.  It might make a mess of my color-coded periodic table and force me to add asterisks and addendums to my list of rules, but the truth is, when atoms react with one another in interesting and unanticipated ways, it illuminates new paths toward solutions to science’s most pressing problems and discoveries that can improve our world.  And that ultimately makes me love chemistry even more.  

Other references:

“Superheavy: Making and Breaking the Periodic Table,” by Kit Chapman

Sarah Anderson is a PhD candidate in the chemistry department at Northwestern University.