The Periodic Table of Elements II: _{History and development}

Beyond hearts and spades

During the years Lottie presided over the roughest saloons of the American Wild West, Russian chemist Dimitri Mendeleev preferred the lower-stakes game of solitaire. The player’s goal is simple: organize a shuffled deck into clubs, hearts, spades, and diamonds using certain prescribed moves.

While Lottie was winning money from outlaws and lawmen, Mendeleev was using this time to ponder a chemistry problem. The 63 elements known at the time consisted of a puzzling mish-mash of characteristics. An atom of carbon, for example, weighs twice as much as one of lithium. Early chemists weren’t sure why. Today we know atoms are made of three particles – electrons, neutrons, and protons. But in Mendeleev and Lottie’s time, these particles wouldn’t be discovered for another 20 or so years.

Mendeleev did, however, understand atomic weight. Atomic weight is the average weight of atoms of a particular element. Clever experimentation with the element hydrogen (H) allowed scientists to measure atomic weights for many of the common elements. Mendeleev also understood – like other chemists of his time – that different elements had different chemical reactivity. While you could drop a piece of carbon (C) into a glass of water with little fanfare, its lighter cousin lithium (Li) reacts violently – releasing hydrogen gas (H2) and generating so much heat that the hydrogen bursts into flames.

While chemists could measure differences in elements’ atomic weight and reactivity, they just couldn’t connect the two – and not for lack of trying, either. But it seemed an atom’s weight had little, if any, say in how that atom reacted.

For example, the element sodium (Na), almost twice again as heavy as carbon, reacts violently with water, releasing flammable hydrogen gas, just like lithium. There must be a connection between the elements. But what that connection might be would remain a mystery for years to come. So early chemists wondered, how can we make sense of the knowledge we do have?

These chemists reasoned that there must be patterns among the 63 known elements. And even if they had no idea what caused those patterns, chemists felt tantalizingly close to at least identifying them and unlocking their predictive power.

So, Mendeleev made a new card deck – one card for each element. First, he grouped elements by reactivity and ended up with eight distinct clusters. Mendeleev numbered them Groups I - VIII. Then Mendeleev laid his element cards out like a game of solitaire. But instead of four suits, he had eight.

When he arranged the elements from lightest to heaviest while taking into consideration the groups he discovered, patterns emerged. Mendeleev was not the first to notice this. But he is credited with putting it all together in 1869, inventing the first useful periodic table (Figure 2).

Mendeleev didn't work alone

Developing Mendeleev’s periodic table required ultra-precise measurements of atomic weight. (Atomic weight is measured in grams and refers to the average weight of a certain number of atoms of that element. Atomic mass is measured in atomic mass units (amu) and refers to the mass of a single atom of that element.) That’s because the original table used atomic weight to determine the elements’ order. And Mendeleev himself did not do this tricky work. That credit goes to chemists of a subdiscipline called analytic chemistry.

One of these was Russian analytical chemist Julia Lermontova (Figure 3), one of the first women in the world to earn a doctorate in chemistry. Lermontova overcame many educational barriers on account of her gender. Petrovskaya Agricultural College in Moscow, renowned for its chemistry program, rejected her application. But Lermontova persisted, and on the advice of a female friend, Sofia Kovalevskaia (who became a mathematician), Lermontova applied and was partially accepted at Heidelberg University in Germany.

At first, Lermontova was only permitted to audit the chemistry lectures. But over time, Robert Bunsen (of Bunsen burner fame) admitted Lemontova to work in his lab. It was here that she helped out Mendeleev.

When Mendeleev asked for Lermontova’s help, it was because his periodic table project had stalled thanks to a pesky cluster of elements: the “platinum group metals.” This group includes ruthenium (Ru), rhodium (Rh), palladium (Pd), osmium (Os), iridium (Ir), and platinum (Pt).

At that time, no one could get accurate atomic weight measurements of these elements. Doing so required that the scientist obtain an ultra-pure sample of the element. A chemist must first collect a sample containing the element. Then the chemist must separate that element from all the other substances with which that element occurred. That separation step was hard. The platinum group elements presented a challenge because they all occurred together in the same mineral deposits and they reacted chemically in the same ways.

Imagine you’re tasked with separating a mixture of sand and salt. You’re given a tool: water. With water, you dissolve the salt and collect the saltwater. Voila! You’ve separated sand from salt.

Think of the salt/water separation as a simplified version of Lermotova’s methods. The problem Lermontova faced, however, was that the platinum group elements all dissolved equally in most solvents. She kept at it though. She tried different solvents and different preparation and separation methods until she was successful. Her work contributed to the determination of these atomic weights (Tiggelen & Lykknes, 2019). Despite this key contribution, the accounts of this work exist only in Mendeleev’s archives and in her correspondence with Mendeleev.

Mendeleev wasn’t the only chemist using such measurements to organize elements. Think of the endeavor as a sort of worldwide project. German chemist Lothar Meyer came up with a table very similar to Mendeleev’s. But Meyer’s table was published a few months later than Mendeleev’s. That’s why today Mendeleev receives much of the credit.

Both Meyer and Mendeleev did more than just group elements. They also predicted the existence of undiscovered ones. The two scientists left gaps in their periodic tables for where these elements might land.

But Mendeleev took his predictions one step further. Like a card game played with an incomplete deck, Mendeleev saw that some suits were missing elements. And he went on to predict the properties of these missing elements. He accurately described the elements scandium (Sc), gallium (Ga), and germanium (Ge) many years before their discovery.

The modern periodic table

The modern periodic table (Figure 4) looks quite a bit different from Mendeleev’s early version. But the premise is the same: Elements are listed from left to right in order of increasing weight, and elements in the same column share similar reactivity. For example, lithium, sodium, and other elements in the same column all react when placed in water.

Reading from left to right like a book, chemical properties gradually shift. That means that as columns get farther apart, differences between elements in those column also widen. One basic change is in atomic number. Atomic number (shown above the element’s symbol in Figure 5) tells you how many protons atoms of that element contain. And each element contains one more proton when reading from left to right. Atomic number increases across the periodic table, starting with the 1-proton element (hydrogen, H) all the way to the 103-proton element (lawrencium, Lr). And maybe beyond.

Recall that Mendeleev knew nothing about atomic numbers in his time. He didn’t even know protons existed! Chemists after him figured out what atoms were made of. Discovering that each element had a unique proton number was a big deal – it means that atomic number alone tells you what element you have. An atom with one proton is always hydrogen, and one with six protons is carbon – it’s as simple as that.

As you can read in our Atomic Theory I and Atomic Theory II modules, atoms contain neutrons and electrons in addition to protons, and these numbers can vary. Not so with atomic number. If the atomic number changes, so does the element.

Beyond atomic number

The other important number that describes each element on the periodic table is atomic weight. You can find atomic weight at the bottom of an element’s block (refer to Figure 5 above).

You might expect atoms to get heavier as atomic number increases, and you would be right. But atomic weight doesn’t increase by 1 like atomic number. That’s because atomic weight is determined by the combination of weights of all of the subatomic particles in an atom combined, not only protons.

An element’s atomic weight gives a rough idea of the total number of protons plus neutrons in atoms of that element. You can see Atomic Theory II for more details. But to summarize, protons and neutrons weigh about the same: 1 amu (atomic mass unit). Most of an atom’s mass comes from these protons and neutrons. Teeny-tiny electrons weigh about 2,000 times less than a proton or neutron. Therefore, we can ignore their contribution to atomic weight and atomic mass. (They are, however, crucial for understanding chemical bonds.) So, for example, Figure 5 shows that carbon has an atomic number of 6 and an atomic weight of 12.01. This tells us that atoms of carbon contain 6 protons (from the atomic number) and the typical carbon atom contains 12 minus 6, or also 6 neutrons; the atomic weight is rarely exact because of the existence of isotopes – a concept we discuss in Atomic Theory II.

Let’s step back from such details as atomic weight and atomic numbers. Instead, let’s shift focus back to the big picture – the periodic table overall. On first glance, the periodic table’s odd shape gives rise to questions like “Why all the gaps?” A wide empty space separates hydrogen (H) and helium (He) (Figure 6). Why not place them next to each other since they are only one number different in atomic number?

The answer: Helium’s chemical reactivity is nothing like hydrogen’s. So instead, helium is placed on the far right with elements like neon (Ne) – a sister element in terms of reactivity. And the spaces between hydrogen and helium are left blank.

The periodic table’s structure keeps elements that react similarly together. (Hydrogen is the one exception, as it reacts similarly to elements on the right-hand side as well as the left.) These groupings are called chemical families. Chemical families group together elements that are similar in how they react, as well as in many of their physical properties. Most periodic tables include eight chemical families, though some families get sub-grouped occasionally.

You can read more about the characteristics of these chemical families in Periodic Table III: Chemical Families. But in summary, the families include alkali metals, alkali earth metals, transition metals, metalloids, reactive nonmetals, halogens, noble gases, lanthanides and actinides. Together, these chemical families make up the periodic table of elements. The beginning of those families trace back to 1869 – to Mendeleev and his deck of handmade element cards.

In Mendeleev’s Group I, we see lithium (Li), sodium (Na), potassium (K), rubidium (Rb), and cesium (Cs), all of which make up the modern-day alkali metals. Mendeleev’s Group 7– containing fluorine (F), chlorine (Cl), and bromine (Br) – matches closely what we today call the halogen family.

But many differences exist as well. For example, francium (Fr) – an alkali metal – is absent from Mendeleev’s table. That’s easy to explain. Francium had not yet been discovered. But other differences are trickier. For instance, copper (Cu), silver (Ag), and gold (Au) appear in the Mendeleev’s alkali metals. Today we group them in the slightly chaotic family known as transition metals. What’s more, some elements, such as silver, appear multiple times on Mendeleev’s table. These differences speak to what was still unknown at Mendeleev’s time. In later years, increasing knowledge of atomic structure would illuminate gaps in knowledge, and this would give rise to the modern periodic table we use today.