History of the Periodic Table Part 2: What is Atomic Mass?

In the early 1800's, as Johan Wolfgang Dobereiner's triadic version of the periodic table was being developed, John Dalton, Thomas Thomson and Jons Jakob Berzelius were beginning to figure out the relative atomic masses of the elements. At the time, each element's mass was taken as a number relative to the lightest known element, hydrogen, which they called number 1. The logic behind this is that scientists believed that each element was built up of atoms of hydrogen. And at this time, knowing nothing about subatomic particles, they considered each atom to be an indivisible unit.

This relative mass idea is known as Prout's hypothesis. Scientists thought that the atomic mass of any element would always be an exact whole-number multiple of hydrogen's mass (1), but soon, to the shock of the scientists involved, this was proved to be wrong. Some measured masses weren't even close. In 1826, Berzelius (shown below right), a man devoted to careful measurement and fastidious lab work, developed a precise way to measure atomic mass through experiment.

He discovered that the atomic mass of chlorine, in particular, fell in between two whole numbers (its mass is approximately 35.45 u). In his Treatise on Chemistry, Berzelius described his procedure for measuring the atomic mass of chlorine:

"I established its [chlorine's] atomic weight by the following experiments: (1) From the dry distillation of 100 parts of anhydrous potassium chlorate, 38.15 parts of oxygen are given off and 60.85 parts of potassium chloride remain behind. (Good agreement between the results of four measurements.) (2) From 100 parts of potassium chloride 192.4 parts of silver chloride can be obtained. (3) From 100 parts of silver 132.175 parts of silver chloride can be obtained. If we assume that chloric acid is composed of 2 Cl and 5 O, then according to these data 1 atom of chlorine is 221.36. If we calculate from the density obtained by Lussac, the chlorine atom is 220 [relative to the atomic weight of oxygen]. If it is calculated on the basis of hydrogen then it is 17.735."

Notice that his measurement is almost exactly half of the modern measurement. The reason is that at the time scientists didn't know hydrogen existed as a diatomic gas so his hydrogen standard was off by half.

Atomic mass can be a confusing concept. Measured in unified atomic mass units, amu or just u, it is sometimes called atomic weight instead - the two terms are often used interchangeably, and most of the time that's fine. However, these are two of several terms in chemistry that create a lot of problems for students and wreak havoc on their teachers. Chemwiki has an excellent definition chart you can use to clear up the chaos. Strictly speaking, atomic mass is the mass of an individual atom at the microscopic scale, whereas atomic weight is the average atomic mass of an element. Isotopes are the reason why there is a subtle difference between the two definitions.

The Discovery of Isotopes: The Modern Atom Begins to Take Shape

Berzelius didn't know this, and in fact most high school chemistry students don't know this (yet), but you can measure the atomic mass of one sample of pure chlorine to fantastic precision and measure another sample of pure chlorine taken from somewhere else in the world and get a different number. Why?

Almost a century after Berzelius's work, around 1910, isotopes were discovered, and the unexpected discrepancies between measurements like the chlorine example I just mentioned, were shown to be due to an isotope effect, where the masses of elements may reflect a mixture of stable isotopes of those elements. This means that atoms of the same element (same number of protons) can have different numbers of neutrons in the nucleus, and that variation affects the atom's mass.

Most people credit the discovery of radioactivity to Henri Becquerel in 1896. He noticed that uranium salts blackened photographic plates, due to some kind of radiation, thought at first to be X-rays. It was not long until researchers such as Ernest Rutherford, Paul Villard and Pierre and Marie Curie realized that the radiation that Becquerel detected was more complex than first thought. And the implications were unsettling.

Around this time, researchers working with thorium, a radioactive element found in naturally occurring thorite minerals, discovered that the naturally found thorium in the mineral emits beta particles (electrons). A thin sheet of thorium in argon is shown below. Pure thorium is a silvery white lustrous metal but when some thorium oxide is present, as it usually is, it eventually tarnishes to black.

alchemist-hp;Wikipedia

They found that thorium isolated from decaying uranium emitted an entirely different particle - the alpha particle. Thorium has over 30 (all radioactive) isotopes, most of which decay by emitting an alpha particle but some isotopes decay through beta decay.

They didn't know what alpha and beta particles were but they could detect that they moved in opposite directions when placed in an electric or magnetic field, and they could detect that one kind (beta) always traveled a lot further than the other kind (alpha). It only took a few years for Becquerel to realize that the beta particle was an electron based on its mass to charge ratio, which was the same as Thomson's results. Otherwise, the two thorium samples were identical. These results flew directly in face of Dalton's atomic theory: If two atoms have the same number of electrons they must also have the same number of protons in their nuclei, and therefore they should behave identically. And yet, these researchers knew that something had to be different between these two thorium samples.

There are many different kinds of radioactive decay and some of them change or transmute one kind of atom into another kind by changing a neutron into a proton or vice versa. Protons were discovered in 1917 when Rutherford expanded on Prout's idea that hydrogen was a standard building block of all heavier elements. Hydrogen contained only one of the newly discovered positively charged particles  emitted through some kinds of radiation (proton emission), while other atoms contained more of these particles. Rutherford named these positive particles protons in 1920.

Along with the isotope mystery, something about atoms was way off. Scientists, looking at the various elements, knew that the relative atomic mass of an atom always seemed to be a bit more than double the atomic number, Z. They also knew that almost all the mass of an atom was concentrated in a tiny volume in the centre (thanks to Rutherford). The atom, as far as they knew,  consisted only of protons and electrons and the atomic mass data meant there had to be twice as many protons as electrons. This was a mystery because they also knew that atoms are electrically neutral. They thought that perhaps half the electrons were bound up in with the protons, cancelling their positive charge somehow. It wasn't a very satisfactory explanation, and the newly formulated uncertainty principle implied that there wasn't nearly enough energy in the atom to confine the (electrically repulsive) electrons inside the positive nucleus.

Neutrons were discovered in 1932 by James Chadwick. It's a bit of a story and the link explains how he did it. His discovery of the neutron finally put the lingering isotope mystery on firm conceptual ground. Isotopes have the same number of protons but different numbers of neutrons in their nuclei. The alpha particle was finally found to be a helium nucleus, consisting of two protons and two neutrons.

Relative atomic mass, which used to be called standard atomic weight, is now calculated as 1/12 the mass of carbon-12 (which happens to be almost exactly 12 u; 12.0107 u to be more exact*). It is the average atomic mass of both stable isotopes of the element carbon (carbon-12 and carbon-13) based on their relative abundance on Earth. However, the isotopic makeup of samples from different sources on Earth can vary quite a bit, and this turns out to be both a little problem and a very useful scientific tool. On the plus side, you can pinpoint the original location of archeological samples of bone, teeth, iron tools, glass and lead-based pigments based on their isotopic profile. On the minus side, this can lead to inaccuracy when relying on the relative or average value for mass. Calculating the geographical variance of the isotopic profile for various elements is still a work in progress, as scientists work out with increasing precision the relative isotopic abundance of elements not just at various locations on Earth, but in the universe as a whole.

As a result, in 2010 the International Union of Pure and Applied Chemistry (IUPAC) changed the formal definition of atomic mass. The atomic masses of hydrogen, lithium, boron, carbon, nitrogen, oxygen, silicon, sulfur, chlorine and thallium are now written as intervals rather than as single numbers. Some modern applications require very precise atomic mass, so this change is necessary for accuracy. *Carbon is now listed as 12.0107 +/- 0.0008 u, a reflection of the varying abundance of two stable isotopes - carbon-12 and carbon-13 depending on the geographical origin of the sample. Only ten elements have so far been updated either because others only exist as a single stable isotope or because the upper and lower mass limits of the element haven't been measured yet. The IUPAC also regularly updates atomic masses as measuring precision improves.

Below is a screenshot from Wikipedia showing the average relative atomic masses of the elements. It doesn't reflect the new mass intervals of ten elements.

Lead (Pb, Z=82, atomic weight = 207.2) is the heaviest stable element. All elements with atomic numbers over 82 have no stable isotopes. The atomic mass of these elements is taken as the mass of the longest-lived isotope, and for some very short-lived elements it is an estimate only.

Mass Defect

When isotopes were discovered, scientists figured that a pure isotope (no mixture) should still have a relative atomic mass that is an exact multiple of 1 to within 1%. In other words, it should have a mass equal to adding up the individual masses of its protons, neutrons and electrons (proton and neutron masses are almost identical). However, this too is now known to be incorrect, and we will use the element helium as an example to explain why.

Helium exists almost entirely as helium-4 on Earth. Based on what we now know about relative atomic mass, we expect its mass to be almost exactly 4.000 - 2 neutrons + 2 protons = 4 (plus a small mass contribution made by electrons). We find by Googling helium that its average measured atomic mass (standard atomic weight) is 4.003.

However, helium-3 is also present in trace amounts on Earth (it is the only other stable isotope of helium). Looking up helium-3, we find it has an isotope mass of 3.016. Wait a minute. We expect it to be almost exactly 3.000 because it is a pure isotope. Why so much off?

This leaves us with two questions: (1) why isn't helium-3, a single isotope, almost exactly 3.000 and (2) why is the average atomic mass of helium slightly HIGHER than 4.000 (4.003) rather than lower, since helium-3 (smaller mass) makes an, albeit, small contribution?

The answer is called mass defect. To illustrate what mass defect is, let's add up the known masses of all the subatomic particles in an atom of helium-4. It has two electrons, two protons and two neutrons. The masses of these particles (in amu) are all known to at least six significant digits so:

2 X proton (1.007276) = 2.014552
2 X neutron (1.008665) = 2.017330
2 X electron (0.000549 = 0.001098

Total = 4.032980

Why isn't this value 4.003, the measured mass of helium? First, let's recap what we know: 4.000 is helium's relative atomic mass (relative to carbon) but it is not its unified atomic mass, the value that reflects its actual average measured atomic mass as it is found on Earth. If we look up helium-4's unified atomic mass, it is 4.002603 amu (or just u), rounded up to 4.003 as shown above. This number is quite a bit different from both its relative atomic mass (4.000) and the value we got by adding up its components (4.032980).

The difference between the sum of the components and the measured amu value is called the mass defect. Energy is released when the helium nucleus is assembled from its protons, neutrons and electrons, so the helium atom has lower potential energy (reflected by 4.003 rather than 4.033). The mass difference, 0.030377 u, is the mass equivalent of the energy that is released. This released energy is called the binding energy. The helium nucleus has lower potential energy than it's separate nucleons but it has higher binding energy, and this is what makes the formation of helium atoms thermodynamically favourable. If we wanted to split a helium-4 atom, we would have to add that energy back.

You might wonder why we have to add energy to helium to get it to break apart. After all, doesn't the fission going on in nuclear plants release energy? The answer is in the size of the atom. Iron is the most stable atom of all atoms. It has the highest binding energy and therefore it doesn't fuse or split. Atoms smaller than iron release energy when they fuse into larger atoms. The larger atom will have higher binding energy and lower potential energy. Atoms larger than iron (including those used in nuclear fission reactions, such as uranium) release energy when they are split apart.

If we add up the components of helium-3, we get:

2 X proton (1.007276) = 2.014552
1 X neutron (1.008665) = 1.008665
2 X electron (0.000549 = 0.001098

Total = 3.024315

Helium-3's isotopic mass is 3.0160293 u, so 3.024315 - 3.0160293 gives it a binging energy of 0.0082854 u, about a quarter of helium 4's binding energy of 0.030377 u.

ALL atoms, including those with radioactive unstable nuclei, have at least some positive binding energy. However, the binding energy in some atoms is not strong enough to hold the nucleus together indefinitely. These atoms will lose neutrons or protons (decay) until they reach a product that is stable. Very unstable nuclei may decay in microseconds while almost stable nuclei may take up to billions of years to decay.

The real measured atomic mass of an element therefore depends not only on its isotopic makeup but on its mass defect as well, and that depends on the particular atom's binding energy. An atom's binding energy consists of its nuclear binding energy (huge contribution since the strong force is involved) as well as its electron binding energy, better called ionization energy (a much smaller contribution since the far weaker electromagnetic force is involved). Ionization energy is the energy required to strip the atom of its electrons, to ionize it in other words.

Atoms with especially stable arrangements of neutrons and protons have especially high mass defects, especially low potential energy, and especially high binding energy. Atoms with unstable nuclear arrangements are radioactive. Calculating stability compared to decay rate is a bit complicated; it is an example of a many-body problem in physics. Physics Stack Exchange provides a really good explanation of it, if you are curious. Radioactive nuclei will go through a decay process. There are three basic decay possibilities: A nucleus will change a proton into a neutron or the reverse (beta decay), it will eject either an alpha particle (helium nucleus) or a proton, or, third, it will eject an even larger element nucleus. These decays either result in a new isotope of the element (if only the number of neutrons is altered) or a whole new element (if the number of protons changes). The latter process is called transmutation and it is the only real way, along with fusion, to change one kind of element into another kind. This is the real-life version of the philosopher's stone mentioned in the previous article.

Helium, with its two neutrons, two protons and two electrons, forms an unusually stable atomic arrangement. The graph below compares binding energy with nuclear size.

Iron (Fe), mentioned earlier, has the highest binding energy (graph peak) while helium-4's binding energy forms an unusually sharp upward spike at the left end of the graph. Helium-4 has a remarkably stable nucleus giving it a significantly lowered atomic mass. If you look at helium-3 above you see that it's right in line of where it's supposed to be. It's nuclear arrangement of two protons and just one neutron gives it an average binding energy.

Helium is extremely inert, which means it is chemically unreactive under all normal conditions so it won't form any compounds. It is also almost always a monatomic gas, condensing to a liquid only at the very cold temperature of 4.22 K, or -269°C, that's very close to the temperature of the vacuum of outer space (2.73 K). Like all elements, these chemical properties of helium come from its electron configuration, which is influenced by the electron's interactions with each other and with the nucleus. Helium's unusually stable nucleus and high binding energy is why element formation in the early universe pretty much stopped after helium nuclei formed (all larger atomic nuclei have been created in stars).

Why isn't hydrogen chemically inert too? Iron, even though it has an extremely stable nucleus, is fairly chemically reactive. Because of its particular electron configuration, it can either lose or accept electrons (usually from water or oxygen) to form various ionic compounds, such as iron oxide (rust). It is the electron configuration, not nuclear stability, which influences the physical and chemical properties of the elements, aside from their radioactivity of course. Chemical bonds between atoms was explained by Gilbert Newton Lewis in 1916, as an interaction between the electrons of the atoms involved.

Between 1800 and the early 1900's, the idea of what an atom is evolved at an explosive rate because there were so many great minds at work on the concept of the atom. Around the same time as the proton and neutron were discovered, the electron configuration of the atom was being sorted out. Bohr's model of the atom hinted that electrons have specific energies within the atom. They can gain or lose only discrete packets or quanta of energy. In the 1920's the quantum mechanical model of the atom was formulated. The evolution from Rutherford's atomic model to the modern quantum mechanical model marks one of the greatest breakthroughs ever in both chemistry and physics, with spin-off progress in biology, engineering, geology, practically every other scientific discipline there is.

In 1817, when Johan Wolfgang Dobereiner was putting together his law of triads, none of these things were known - isotopes, nuclear binding energy and ionization energy. No one knew that mass and energy were equivalent. No one knew exactly how atoms interact with each other, why they give off light and other radiation and how they transmit heat. No one knew how the fundamental forces make atoms what they are. All they knew was that the Earth seemed to be composed of a growing list of various simple substances, substances that seemed at the time to be fundamental, meaning they can't be broken down into anything smaller, and that some substances reacted with other substances to make yet different substances and others did not react at all. Unknown to researchers of this time, they were taking the first steps toward an amazing new era of science.

History of the Periodic Table Part 1: From Alchemy to Mendeleev

Memorizing the periodic table and old dead guys is a quick way to turn people off chemistry. But, like many things, the background and context bring it to life. For teachers, I hope this is a refreshing second look. I this six-part series of articles, we will explore not only the timeline itself but some aspects that are less often discussed, such as how the concept and measurement of atomic mass was developed, how X-ray spectroscopy works and what it has to do with the periodic table, what actinides and lanthanides are and why they are kept separate. We will also explore the rare earths in particular and why their story is currently newsworthy, Finally we will close with a glimpse into the table's future.

There are fascinating tales of mystery, intrigue, competition and greed behind the evolution of chemistry, and the history of the periodic table is really about that. It is about how magic evolved into science, and yet the elements themselves are interesting in their own right. They are puzzles and their relationships to one other offer clues into the deepest nature of the atom itself. We will explore this in detail in this series. We may be surprised to to find that the periodic table we all had to learn may be about to change.

Here is Wikipedia's classic periodic table of elements:

DePiep;Wikipedia

My favourite table comes from the Royal Society of Chemistry (RSC). If you click on any element there you will get a scroll-down index of all kinds of information about it. Wikipedia also does a good job if you want to search an individual element by name. If you look up iron for example, you will see a photo of it and a useful chart at the right with electron configuration (we'll find out why this is so important), melting point, oxidation states, etc.

The following 11-minute video is a nicely done introduction to the history of the periodic table created by CrashCourse.



It Begins with Gold and Black Magic

The first known elements were those that that stood out visually and were accessible. Ancient people found ingots of gold where they lay scattered about on the soil or just beneath. Gold, along with copper, was collected, melted down and shaped into decorative objects as early as about 6000 BC. It wasn't until about 330 BC, however, that people such as Aristotle looked at these pure materials and began to wonder if these and other less visually distinct materials ultimately come from some basic "prima materia" or first matter. At around this time, Plato, thinking along the same lines, suggested there are four basic building blocks of all matter - earth, water, air and fire - and he named them "stoicheia," the Greek word for elements. For many centuries afterward, no one knew how these elements formed all the different materials on Earth, but it was thought that everything stemmed from a single mysterious formless source variously named chaos, quintessence or the aether.

Ingots of gold are beautiful in their own right and it's not surprising that people were not only curious about how these quite rare materials formed but also how they could get more of them. In the meantime, between 6000 BC and 750 BC, a variety of other useful metals were discovered. In order of discovery they were silver, lead, iron, tin and mercury. The craft of smelting - getting the metal to melt and separate from its ore - improved over time and made many of these discoveries possible. While gold, copper and silver can be found in their native form, other metals like lead, tin and iron (except for pure iron found in meteors) are not. Metals, hard and malleable, were a huge boon for ancient people. Metal tools and weapons were vastly superior to those made of wood, rocks or bone. Unknown to them, these metals were the first elements to be identified, extracted and purified.

At around 300 CE, or possibly even before this, what would become the legend of the philosopher's stone took shape. People thought that some alchemical substance with magical properties could turn more common base metals such as lead and iron into more rare and highly valuable gold or silver, and thus began the alchemical race for the philosopher's stone. This might seem stupid to us living in the quantum age, but back then, without scientific knowledge to rest on, it would have been only logical to conjure up magical origins for phenomena we don't understand. There was a significant magical component to very early medicine as well. And, if you look around you, you will come across many vestiges of magical thinking today. Consider superstitions.

This stone was not only considered to be a physical material but a symbol of perfection and an elixir of life and immortality as well. Prima materia was thought to be a starting ingredient in a recipe for the philosopher's stone. The quest to find that recipe went on for many centuries.

Along the way, scientific progress was made, if by sheer trial and error. An alchemist in the 8th century, Jabir ibn Hayyan, surmised that every metal must be made of a combination of four principles. Elaborating on the four original elements, these principles were hot/dry, cold/dry, hot/wet and cold/wet. By rearranging these principles and applying some kind of elixir (the philosopher's stone), he thought he could transmute one metal into another metal. Although gold is generally found in its native pure form as nuggets, he thought that yet more gold resided hidden within different principle mixtures (we call them alloys and ores today) and it was just a matter of treating those mixtures with an elixir to release them. Smelting (roasting the ore or mineral over fire to release molten metals such as copper, lead, silver, tin, iron and mercury) was known since 6000 BC, but chemical metallurgy was not. By the 14th century, one such chemical treatment, called aqua regia (royal water), was discovered. Not the pretty solution you might expect an elixir to be, it was a fuming highly corrosive orange/yellow mixture of powerful acids that can dissolve certain metals such as gold from minerals and alloys so they can be recovered in pure form. This was the first chemical extraction of an element, as opposed to the physical (heat) extraction method called smelting.

Henning Brand - From Pee to Phosphorus, or, From Alchemy to Chemistry

Still, as of 1649, alchemists remained on the hunt for the elusive philosopher's stone, a substance that could do one better than extract a substance - it could create it anew. A German merchant called Henning Brand, trying to find the stone, ran experiments on distilled human urine, and discovered not gold but a white substance that glowed pale green in the dark, which he named phosphorus, the name owing itself to the Greek word for "light bearer."

Of all substances, why urine? At the time philosophers believed that man's body is a microcosm of the universe, so bodily fluids should contain, like the world itself, gold among all other materials. It is reported that he eventually came across a recipe in a then fairly recent tome called "400 Auserlensene Chemische Process" that called for using a mixture of alum, salt peter and concentrated urine to turn base metals into silver (it didn't work). So, I imagine that book's promise placed gold in the realm of possibility too. I can't help but chuckle here. Men.

This discovery would have been awesomely horrific: imagine a giant cauldron, with a fire roaring underneath it, boiling with pee. Eventually the urine concentrates into syrup. A glowing liquid trickles out the bottom spigot - itself entirely aflame.

The chemistry of the process is this:

Urine is rich in potassium salts. Evaporating it produces, among other salts, ammonium sodium hydrogen phosphate, or (NH₄)NaHPO₄. Heating the evaporate decomposes it into sodium phosphite, ammonia and water:

(NH₄)NaHPO₄ → NaPO₃ + NH₃ + H₂O

Heating sodium phosphite with charcoal decomposes it into carbon monoxide and white phosphorus (not an especially riskfree reaction):

8NaPO₃ + 10C → 2Na₄P₂O₇ + 10CO + P₄ (white phosphorus)

This experiment (as well as Boyle's famous phosphorus experiment to follow) is recreated in the first episode of the BBC series Chemistry: A Volatile History (a link to watch it can be found at the end of this article).

Though not precious gold, I am sure he reveled in his man-made product, aglow with some kind of mysterious life force! Below is a sample of white phosphorus (as a solid) under water.

BXXXD; de.wikipedia

Not a substance to mess around with, white phosphorus, also known as Willie Pete, is highly reactive (hence the water) and volatile. It has often been used as an incendiary weapon of war. Set it aflame and everything it touches including humans lights on fire, stays on fire, and creates a thick smothering smoke. It is also what makes matches burn as long as they do.

Brand's journey was captured in the famous 1771 painting "The Alchemist in Search of the Philosopher's Stone," shown below.

This painting romanticizes the actual process that was described in 1730 as requiring 50-60 pails of urine that was both putrid and "bred worms" (chuckling again).

Joking aside, this was the first element to be chemically discovered and its discovery was the catalyst (sorry) that ushered in the age of modern chemistry as we know it. There was a lot of fascination around this new product and Brand sold his secret recipe to anyone willing to meet his price. At the time, alchemy was a shadowy secretive world, filled with arcane symbols and recipes and procedures that were rarely shared. The RSC Periodic table has an alchemical version where you can see the (very beautiful) alchemical symbols for various elements that were known at the time. They are quite fascinating. For example, the symbol for iron (right) is also the symbol for Mars and for masculinity.


An ironically fun fact here: phosphates (salts of phosphorus) are one of three essential nutrients (nitrogen, potassium and phosphorus) for plants. There are some movements underway around the world to once again harness the phosphates in urine to use in fertilizer, as geological supplies of phosphate rock dwindle.

Phosphorus, later rediscovered by Robert Boyle, led many others to wonder what exactly an element is. In 1661, Boyle opened up current alchemical knowledge to the world by publishing a book called The Sceptical Chymist in plain English. He defined an element as "any substance that can't be broken down into a simpler substance by a chemical reaction," a good working definition that serves well even after the discovery of subatomic particles (particularly electrons) in the late 19th century with the work of J.J. Thompson and others.

The transition from the Middle Ages to the Age of Enlightenment marked the gradual transition from alchemy to chemistry, as notions of transmutations and the philosopher's stone gave way to the hunt for new "simple substances."

Antoine-Laurent de Lavoisier - the Father of Modern Chemistry

Antoine-Laurent de Lavoisier was the first person to categorize a list of all then known "simple substances." He placed it in a book called Traité Elémentaire de Chimie (Elementary Treatise of Chemistry). It's quite the volume, over 500 pages. You can see a translation of it online created by Project Gutenberg. He brought the concepts of balancing equations and the conservation of mass to chemistry. The law of conservation of mass is the rationale behind balancing a chemical equation. This is his formulation of the law translated into English:

'We may lay it down as an incontestible axiom, that, in all the operations of art and nature, nothing is created; an equal quantity of matter exists both before and after the experiment; the quality and quantity of the elements remain precisely the same; and nothing takes place beyond changes and modifications in the combination of these elements.'

He also explained combustion in terms of combination with oxygen, a breakthrough over an earlier theory in which combustible substances were filled with a fire-like liquid called phlogiston that was released when the substance burned. Below is an elegant portrait of him and his wife (who was also a chemist) painted in 1788.

This was the first chemistry book ever written and it included elements such as oxygen, nitrogen, hydrogen, phosphorus, mercury, zinc and sulphur, among others. When Lavoisier's chemistry book was published (1789), over 20 elements were known. They also included light and "caloric," which were thought at the time to be basic material substances too. Light is, well, light and caloric is what they thought heat was made of - a liquid that flowed from hot objects toward cold objects. This book's chemical classifications were simple but classic: metals and nonmetals.

During the next century many more elements were discovered - 56 by the year 1850. The RSC period table has a history version where you can simply plug in any date from 1 CE to 2014 and all elements known at that date are highlighted. You can click on each element to read a short story about its discovery. Wikipedia has a good timeline of element discoveries as well.

In the 1800's, people were beginning to wonder how this wide variety of substances related to one another. There was a human need to get a handle on them by comparing their physical and chemical properties and categorizing them. In 1817, Johan Wolfgang Dobereiner noticed that there were trends in the properties of elements, so he came up with a way to classify them accordingly. He organized the elements into groups of three so that each element in a group shared related known properties.

John Dalton - the Atom Within the Element

Just a few years prior to Dobereiner's work, John Dalton and others achieved a huge breakthrough that would be tremendously helpful in categorizing these substances. Dalton figured out that substances are made of atoms. This first atomic model was Dalton's model. It was pretty simple - an atom was just a small indivisible object. You can think of it as a tiny solid sphere. Dalton, a physicist (shown right), revolutionized chemistry in 1802 with a series of papers on his atomic theory.

Based on his work with gases, he realized that each pure element is made up of identical atoms, and that atoms of different elements combined with each other in fixed ratios, a revolutionary idea. He also discovered that atoms could be told apart from each other by their unique relative atomic weights. Below is a scan of the first page of his "A New System of Chemical Philosophy" published in1808.

Its quality isn't great but you can see at the bottom of the diagram how much work he did figuring out how atoms combined with each other to make various compounds. The names of many of Dalton's atoms reveal a work in progress. For example, elements 8, 9 and 10 are called lime (which is actually a group containing many calcium-containing compounds), soda (there are actually many sodium salts) and potash (there are many potassium salts). We now know that number 37 is not a 'septemary' element composed of atoms of sugar, with each atom composed of 1 atom alcohol (33 left) and 1 atom carbonic acid (28 left). We know that sugar is a compound, composed of carbon, hydrogen and oxygen atoms. But Dalton based his results on what he knew from experimentation - that sugar ferments into two products - acid and alcohol. He knew that water (21) is made of oxygen (4) and hydrogen (1), except that he thought the ratio was 1 oxygen to 1 hydrogen because at the time no one knew hydrogen was a diatomic gas, and he also thought of water as binary atom, rather than a compound.

Wolfgang Dobereiner was the first person to try to classify the elements based on Dalton's work. He grouped them into clusters of three, arranging them in increasing relative atomic mass so that the mass of the middle element was close to the mean of the outer two elements. The next article in this series, History of the Periodic Table Part 2, explores in detail how the concept of atomic mass was developed and how its measurement was, and continues to be, refined.

Examples of Dobereiner's Triads (numbers are relative atomic masses) are shown below:

An example of one of his triads - lithium, sodium and potassium - is shown below left.

These three elements have a lot in common. They are all a silvery colour. You can squash them with a knife. They float on water. They melt very easily. And, they are all highly reactive. They have to be stored under oil but even under oil they gradually corrode, reacting with traces of oxygen that eventually seep through the oil. Now we group these elements as alkali metals.

This particular grouping works but in general the triad concept wasn't very successful. Many similar elements (judged by appearance and reactions) such as the transition metals (nickel, copper, chromium, zinc, platinum, etc.) can't be arranged in triads, while chemically dissimilar elements can be placed in triads. However, the triad arrangement did provide a clue about the relationship between element appearance/behaviour and atomic mass.

Several more attempts were made to categorize the elements, with none having much more success than the triad table. In 1865, John Newlands devised a 'law of octaves' for the elements (56 of them known by then). He noticed that many pairs of similar elements existed but they differed by a multiple of eight in their mass number, so he organized the elements into eight rows and gave each element an atomic number. A scan of his original list is shown below.




(credits for element photos respectively: Tomihahndorf;Wikipedia, Dnn87;en.wikipedia, http://images-of-elements.com/potassium.php;wikipedia)

For example, lithium (Li, atomic number, Z=3) and sodium (Na, Z=11) (his atomic numbers are a bit off) have similar properties as we just saw. Both are soft and can be cut with a knife, they are good electrical conductors, they both react exothermically with water.

Beryllium (Be, atomic number Z=4) (below left) and magnesium (Mg, Z=12) (below right) are also fairly similar to each other in appearance and chemistry. Both are strong rigid metals and both form a thin layer of oxide in air.

(Alchemist-hp;Wikipedia and Wanut Roonguthal;Wikipedia)

The problem with Newlands' scheme is not that it was fundamentally wrong. The problem was that he compared it to an octave musical scale. The fairly recently formed Chemical Society of London thought his musical note theory was far too ridiculous to publish until the octet theory of chemical bonding was established in 1916, and its importance was finally recognized. The octet rule we know is basically a modern rewrite of Newland's rule of octaves with a few corrections. Both systems have limitations. For example, they work well only for second period elements (lithium (Z = 3) to neon (Z = 10)).

Dmitri Mendeleev, the Father of the Periodic Table

It was not long afterward (1869) that the man we know as the father of the periodic table, Dmitri Mendeleev (below), published his periodic table in an obscure Russian journal.

His arrangement provided spaces for elements that were not yet discovered and it could predict some of these unknown element's characteristics, based on their locations in the table. The only group missing was the noble gases, a group of odourless colourless almost nonreactive gases that were discovered later. And thallium, lead, mercury and platinum were in the wrong groups.

This is what his handwritten version looked like (below right), not the neat and colourful castle turret block diagram we know today:


X-ray Spectroscopy Refines Mendeleev's Table

In 1914, X-ray spectroscopy was a new and very au courant investigative tool thanks to the work of W.H. and W. L. Bragg and Maurice de Broglie. Physicist Henry Moseley took advantage of this tool to find a relationship between the X-ray wavelength of an element and its atomic number, Z. This was the first logical order imposed on the elements that is based on physics, and it also refined Ernest Rutherford's model of the atom, which was published just a few years earlier, in 1911. Rutherford discovered that the atom, rather than being a uniform sphere, consisted of an intense central positive charge concentrated in a tiny volume at the center surrounded by a more diffuse but equal negative charge. Moseley's work suggested that the central positive charge was not a single positive charge but actually up to several positive charges instead, the number of which was equal to the element's atomic number in the table. To explore in detail how Moseley's work brought the periodic table into the quantum age, see History of the Periodic Table Part 3: Spectroscopy and the Quantum Atom.

Mendeleev's periodic table doesn't much look like our modern version. I created a block version of his table, shown below, with Dobereiner's triads marked out as yellow squares, as well as the elements known to the ancients (squares with dark green dots). The green squares are elements, such as noble gases, unknown in Mendeleev's time but he left the green gaps to be filled in later. Approximate atomic masses are shown beneath each element's symbol.

Beyond Mendeleev's Table

Many more elements are known to us today. Our modern version of the periodic table organizes elements into blocks, shown below.

Roshan220195;wikipedia)

This block arrangement reflects a deeper quantum mechanical understanding of the different atoms. Many scientists, working at the beginning of the 20th century, ushered in a new era called quantum physics. With this came a new atomic model, called the quantum mechanical model. Beginning with J.J. Thomson's discovery of the electron in 1897, and working from the Bohr and Rutherford models in which electrons orbit a positive nucleus in concentric orbits with increasing energy as one goes outward, rapid progress was made in determining exactly how electrons behave inside an atom. Physicists Albert Einstein, Louis de Broglie, Erwin Schrodinger, Werner Heisenberg, and Max Born all contributed to the modern model of the atom in which electrons orbit in orbitals defined by a probability distribution that is determined by the electron's angular momentum.

Blocks s, p, d and f are derived from the orbital configurations of the electrons in the atom, which themselves derive form the quantum angular momentum numbers of the electrons. To learn about orbitals and subshells and to review them, read Atoms Part 4A: Atoms and Chemistry - Atomic Orbitals and Bonding. Orbital configuration will also be explained in detail in History of the Periodic Table Part 4: Lanthanides and Actinides - Elemental Misfits?

The s-block contains alkali earths and alkali metals. The p-block contains all the nonmetal elements with the exception of hydrogen and helium. The d-block contains transition metals. The f-block contains inner transition elements - the lanthanides and actinides.

S-block elements exhibit well-defined trends in their physical and chemical properties, which can be explained by the increasing number of valence electrons filling up the s-subshell. The valence (these are highest energy outermost electrons involved in chemical interactions) electrons in the P-block elements fill up the p subshell. Valence electrons in the d-block elements fill up the d subshell, and f-block elements have valence electrons that fill the f subshell. What sets the f-block elements apart is that after the first element in each series (for example, lanthanum, Z=57, in the lanthanide series), the energy of the outer 4f subshell falls below that of the inner 5d subshell, so electrons from cerium (Z=58) onward start to fill up the 4f subshell before any more electrons are added to the 5d subshell. This places them in the f-block even though technically they fit into the d-block (see the light and dark blank pink squares in the table at the the beginning of this article). We will explore this in greater detail in History of the Periodic Table Part 4.

An additional hypothetical g-block would consist of elements with atomic numbers higher than 117. These elements would be filling g orbitals and all would be unstable and therefore radioactive. In fact, all elements above Z = 82 (lead) are unstable and the trend is that the half-lives of elements above lead generally decrease (meaning they are increasingly unstable) as atomic number increases. No elements larger than Z = 118 (ununoctium) have been created or discovered. Ununoctium has a half-life of just 0.89 milliseconds, and only a few atoms of it have ever been created in a collider. However, researcher Walter Greiner predicts there are more elements to discover and, in  fact, there may not be a highest possible element, at least theoretically.

Some elements may exist in what Glenn Seaborg, a nuclear physicist, called an island of stability: a group of elements with a neutron number (N) around 178 and an atomic number (Z) around 118 that would be unusually stable with half-lives at least minutes long and perhaps far longer, demonstrated below in a 3-dimensional map.

InvaderXan;Wikipedia

This hypothesis is built on the nuclear shell model in which the atomic nucleus is built up of energy shells, analogous to electron energy shells in which completely filled shells impart extra stability to the atom. This arrangement implies that certain numbers of nucleons in the nucleus are magic numbers - they create extra stable nuclei. Furthermore, researchers expect the shells for protons and the shells for neutrons to be independent of each other, so there are doubly magic numbers of nucleons imparting even greater stability to the nucleus. One such doubly magic number is around 178 neutrons and 118 protons, the island of stability, above.

These additional elements could be placed in an extended version of the periodic table, shown below. This is hard to see so you can find a large version of it here.

The extended periodic table would include additional blocks of elements such as super-actinides and eka -super-actinides.

In History of the Periodic Table Part 5, we will focus on the lanthanide series in the f-block elements - the rare earths. These elements have unique properties that are essential to the high-tech industry, a story that is gaining political energy. Part 4 and 5 articles continue our story of the history of the periodic table and bring it into the quantum mechanical age.

Finally, in History of the Periodic Table Part 6, we will explore the future of the periodic table further. In light of quantum mechanics, some experts believe that the periodic table is long overdue for makeover. The extended table above is a hint at what may come.

Finally, I recommend a BBC series of three 1- hour documentaries called Chemistry: A Volatile History (released 2010) as a must watch for those of us fascinated by the history of the periodic table and of chemistry in general. It is not easy to find and doesn't seem to be available for purchase, but you can watch the series (complete with episode descriptions) here at brainpickings.org.