Skip to main content

Chemistry: The Periodic Table

Chemistry: The Periodic Table

Introduction

How did scientists come up with the idea for the periodic table? How has it changed as we learn more about atomic structure? And exactly why does it summarize so much useful chemical information?

The ancient Greek philosopher Aristotle (384–322 BC), British natural scientist Robert Boyle (1627–1691), and French chemist Antoine Laurent Lavoisier (1743–1794) all played important roles in this story, but the scientist who discovered the periodic system was Dmitry Ivanovich Mendeleyev (1834–1907). Ironically, he had no idea of the true basis of his table. He thought the chemical characteristics of an element depended on its atomic weight, although we now know that its key concept is atomic number. Yet following the reasoning processes of people like Mendeleyev, even when their result is not exactly the right answer, explains much about scientific progress.

Historical Background and Scientific Foundations

The Greek Theory of Four Elements

Modern science has its roots in the way ancient Greek philosophers such as Aristotle thought about the natural world. Earlier civilizations like the Babylonians had made careful astronomical observations; ancient Egyptians developed elaborate metallurgical techniques for refining copper and making alloys of gold. To this practical knowledge, the Greeks applied a curiosity about the underlying causes of natural phenomena.

Aristotle tried to find the fundamental elements of materials, such as solids, liquids, and gases. In On Generation and Corruption, written in 350 BC, he expressed his goal this way: “we are looking for the ‘originative sources’ of perceptible bodies” to understand the origins of their properties and how changes in their appearances take place. He assumed that there would be only a few, although we now know there are over a hundred chemical elements.

Aristotle postulated that there were four elements: earth, air, fire, and water. There were also two pairs of contrary qualities: hot/cold and wet/dry. Each element was associated with two of these fundamental qualities. So, for example, fire was associated with hotness and dryness. Its opposite, water, was cold and wet. Adding water to fire would temper its hot, dry qualities; the result would be something in between. Aristotle emphasized that these four elements were usually not encountered in pure form. For example, both metals and sand were “earth's,” but a ductile metal such as gold probably was a combination of the pure element earth with the pure element water. Blood was a liquid, an indication of the presence of water, but because of its role in sustaining life, it probably contained an unusual amount of elemental fire.

Aristotle's theory seems terribly simplistic and fanciful today, but it dominated scientific thought for centuries. Nevertheless, his attempt to explain the multiplicity of chemical materials in terms of a limited number of active ingredients still guides research today. The Aristotelian account, however, also allowed for the possibility that one material could transmute (change) into another.

Consider the Aristotelian account of what today we would call the water cycle. When the sun shines on water, the cold, wet water is turned into hot, wet air, or steam. If heat is removed, we once again get water. On a very cold, dry day (rarely encountered in the Mediterranean), the water might be transformed into ice, a material that has the coldness and dryness characteristic

of earth. This idea of impressing qualities on matter formed the philosophical basis for alchemy, an early scientific quest to turn base metals, such as lead, into gold.

Alchemical Principles

As the Roman Empire disintegrated, much ancient learning and science was lost. For example, the great library in Alexandria that had been started by one of Aristotle's students was burned on several occasions and eventually destroyed. As the Arab successors of Mohammed (570–632) occupied much of the Roman and Persian empires, interest in natural philosophy shifted eastward and became centered in Islamic cities such as Baghdad. Important Greek texts were translated into Syriac and Arabic. Ancient texts that alluded to the legendary Philosopher's stone and the possibility of transmuting base metals into gold were rediscovered.

Alchemy led to many important chemical discoveries. Early alchemists designed elaborate glassware and learned how to distill liquids. They synthesized hydrochloric, sulfuric, and nitric acids and used them to dissolve metals and to react with soda, potash, and lime.

Early alchemists admired Aristotle and continued to work within his framework of four elements. But the Aristotelian qualities of hot & cold, wet & dry seemed far removed from the properties of the new chemicals they discovered. So they postulated three additional fundamental principles: sulfur, salt, and mercury.

Like Aristotle's elements, these, they believed, were never encountered in pure form. “Philosophical sulfur” was not the yellow powder used to fumigate unhealthy rooms, but it was responsible for the combustibility of materials. “Salt” was the principle that underlay the reactivity of acids (when acids are neutralized they form ordinary salts) and “philosophical mercury” accounted for metals' malleability and shiny surfaces.

Not surprisingly, ordinary mercury was a popular starting material in attempts to make gold. It is easy to laugh at stories of alchemists mixing up mercury, lead, and some bright yellow material such as saffron, burying it in manure, and then warming the whole ensemble in an oven for weeks. Yet there was a method behind this apparent madness.

It was widely believed that metals “grew” in mines, and gold nuggets were often found embedded in ore containing copper. Perhaps, they reasoned, copper or other metals were just immature forms of gold. If they could only speed up the growth process, alchemists reasoned, maybe they could duplicate in the laboratory what nature did more slowly in the bowels of the earth. Manure appeared to be a likely catalyst. After all, little insects and maggots were formed in manure, they believed, which gave it strong generative powers.

As Muslim conquerors moved across North Africa and entered Spain in AD 711, they brought scientific writings in Arabic, both translated texts of the ancient Greeks and innovations of their Arab successors. In the twelfth and thirteenth centuries, many of these writings were translated into Latin, often by Jewish scholars. The influx of these preserved and rediscovered Greek texts, along with Arabic contributions to natural philosophy, accelerated the rebirth of learning in European universities. The German-Swiss physician and alchemist Paracelsus (1493–1541) was particularly taken with the three new alchemical principles and promoted them as part of the new doctrines that would eventually supplant the old ancient Greek dogmas.

Paracelsus, sometimes called “the Luther of physicians,” was quite a radical. He once burned copies of Aristotle's and the Greek physician Galen's (129–c.216) books in a public bonfire at the University of Basel to show how dangerously outdated their theories were. Paracelsus introduced the concept of chemical medicines and recommended a mercury compound for the treatment of syphilis. (Luckily only small portions of this poison were prescribed!)

The Phlogiston Theory

Nearly 200 years after Paracelsus, German chemist and physician Georg Ernst Stahl (1660–1734) modified the three-principle theory and used it to systematize a whole series of important reactions. He renamed the sulfurous principle “phlogiston” (from the Greek phlogistos for “flammable”) and claimed that it was contained in every combustible material, and forcibly expelled in all combustion reactions producing heat and light. To explain why combustion won't occur unless there is lots of air around, Stahl claimed that it was necessary to carry away the phlogiston.

We now know that Stahl had it just backward—when materials such as charcoal or sulfur burn, they are not decomposing. Instead, something from the air is added. The air is a passive receptacle for something that is given off; it supplies oxygen, which combines with whatever is burning. The heat and light given off is energy produced by the rapid oxidation reaction.

Yet Stahl successfully used the phlogiston theory to systematize some important chemical reactions—a good example of how scientists can make progress even when their starting assumptions are wrong. Stahl's first insight was to note the parallels between what happens when metals are heated in air and turned into a powder (a process he called “calcination”) and ordinary combustion, such as when wood burns. Stahl noted that in both cases heat and light might be given off (perhaps you may have seen the sparks when iron filings are sprinkled into a flame) and that air was required to keep the reaction going. (He would have said phlogiston was escaping while we know that oxygen is being combined, but we also recognize the parallel between the combustion of metals and nonmetals. Both are oxidation reactions.)

Then Stahl used his system to describe what happened when metallic ores were reduced to metals. He noted that copper miners often mixed the minerals they took from the ground with charcoal and heated them while stirring the molten results with a green stick. Stahl reasoned that the metallic ores had lost their phlogiston. They were powdery just like the metal calxes that were the products of calcination. To turn them into shiny metals, one needed to add phlogiston. But he was convinced that charcoal was a rich source of phlogiston because it burned so well. So although they didn't realize it, the miners were transferring phlogiston from the charcoal into the metal ore and thus producing metals. Since the miners thought of the charcoal as only a fuel, not as an ingredient in a chemical reaction, they sometimes didn't add enough charcoal to reduce all of the metal present in the ore. So Stahl explained to them the necessity of doing so.

Stahl spoke of the reactions above as analyses and syntheses. In calcination, a metal was analyzed (separated) into a metallic powder and phlogiston. What the miners were doing was synthesizing the metal by combining the metallic ore with phlogiston that came from the charcoal. One of Stahl's students proposed the axiom that every analysis should be followed by a synthesis if one wanted to prove that one understood the composition of a material. Stahl certainly felt that he had proved that metals were compounds, not elements. Lavoisier was soon to use the same methods of analysis and synthesis to prove him wrong!

Lavoisier's Scientific Criterion for Chemical Elements

Aristotle thought of elements as “originative principles.” Nearly 2,000 years later in The Sceptical Chymist (1661), the Irish chemist Robert Boyle defined elements as “certain primitive and simple, or perfectly unmingled bodies; which not being made of any other bodies, or of one another, are the ingredients of which all those called perfectly mixt bodies are immediately compounded, and into which they are ultimately resolved.”

But how could chemists determine exactly what these primitive and simple materials were? Aristotle had proposed four elements: earth, air, fire, and water; the alchemists added salt, sulfur, and mercury. Yet none of these so-called elements could be put in a bottle or used in experiments. (Remember that the alchemists' third element was “philosophical mercury,” not the quicksilver that was used in barometers.)

It was the French chemist and public servant Antoine Laurent Lavoisier (1743–1794) who finally provided a workable definition of a chemical element in 1789. Since it was impossible to reach agreement about the invisible constituents of bodies, Lavoisier proposed in Traité élémentaire de chimie (Elements of Chemistry; 1789) that: “we apply the term elements, or principles of bodies, to express our idea of the last point which analysis is capable of reaching, we must admit, as elements, all the substances into which we are capable, by any means, to reduce bodies by decomposition.”

Lavoisier recognized that the list of elements might change as chemists figured out how to further decompose complex materials. Thus he tentatively included lime and magnesia on his list, but noted that they might not be simple substances. (We now know they were oxides of calcium and magnesium.) Still another step was required before chemists could unambiguously agree on a list of elements. Stahl was dead by the time Lavoisier proposed his definition, but had he been alive, he might have argued that metals were not elements—they contained phlogiston, which was then freed upon heating. So phlogistonists considered a metal calx to be a simpler substance than a metal.

Lavoisier provided a direct way to show that metals were indeed elements. He weighed the sample of metal, heated it in air, and then weighed the powdery product. Since the calx was heavier, something must have been added in calcination. Through more elaborate experiments Lavoisier isolated the gas from the atmosphere that was absorbed during both calcination and combustion and called it oxygen, from the Greek oxys, which means acidic, or sharp.

Chemists rapidly accepted Lavoisier's definition of an element and adopted the helpful strategy of tracking the weights of products and ingredients. When Lavoisier published his Elements of Chemistry, there were 33 elements on his list, including the common metals, sulfur, arsenic, phosphorus, oxygen, and nitrogen. Soon hydrogen and chlorine were discovered by means of electrolysis; by the time Mendeleyev began his textbook, there were over 60, each with a set of distinctive chemical properties.

Assigning Atomic Weights

Throughout history, natural philosophers and chemists speculated about the existence of atoms. Because both Aristotle and (much later) French mathematician

René Descartes (1596–1650) thought that a vacuum was physically impossible, they both rejected Democritus's (c.460–370 BC) theory of atoms in a void. However, virtually everyone was willing to entertain the idea that matter might be composed of tiny, invisible “corpuscles.” Yet as was the case with elements before Lavoisier, it seemed impossible to find out anything specific about these particles. It was John Dalton (1766–1844), a British meteorologist and schoolteacher, who came up with the crucial inferences that made modern atomic theory viable.

In a book called A New System of Chemical Philosophy (1808), Dalton noted that it was well established that when elements combined to make a new compound, they almost always did so in the same proportion by weight. (Alloys were the exception.) So it seemed reasonable to assume that the compound always contains the same ratio of atoms. Dalton put it this way:

Therefore we may conclude that the ultimate particles of all homogeneous bodies are perfectly alike in weight, figure, &c. In other words, every particle of water is like every other particle of water; every particle of hydrogen is like every other particle of hydrogen, &c.

Knowing the formula for a compound, Dalton posited, would allow him to calculate the relative weights of the atoms. But how could he determine the formula? To solve this problem Dalton took the bold step of postulating rules for making conjectures about the formulae for compounds:

The following general rules may be adopted as guides in all our investigations respecting chemical synthesis.

1st. When only one combination of two bodies can be obtained, it must be presumed to be a binary one, unless some other cause appear to the contrary.

2d. When two combinations are observed, they must be presumed to be a binary and a ternary.

For example, if carbon and oxygen form only one compound, Dalton assumed it contains only two atoms. Hence its formula should be CO. So if a sample of carbon monoxide (using our name for it) consists of 3 grams of carbon combined with 4 grams of oxygen, we presume the ratio of their atomic weights is 3 to 4. If, however, carbon and oxygen form two different compounds (which is actually the case), then one of the combinations is CO, while the other consists of three atoms. Hence its formula should be either CO2 or C2O. If it analyzes into 3 grams of carbon for every 8 grams of oxygen, then we know the correct formula is what we now call carbon dioxide.

This example illustrates the way Dalton reasoned, but it makes the job of coming up with a comprehensive, consistent system of relative atomic weights look much easier than it actually was. First of all, the experimental data on combining weights were not very accurate. Secondly, Dalton got off on the wrong foot because since only one compound of hydrogen and oxygen was known at the time, his rules led him to assign water the formula HO instead of H2O, and when he set the atomic weight of hydrogen as 1, oxygen came out to be half of its correct value!

Nearly 50 years later, at an 1860 conference in Karlsruhe, Germany, Italian chemist Stanislao Cannizzaro (1826–1910) finally presented a method for determining atomic weights that all chemists found acceptable. Attending that conference was a young Russian chemist studying in Germany. His name was Dimitri Ivanovich Mendeleyev (1834–1907).

Mendeleyev's Periodic Table and Periodic Law

Mendeleyev was born in a small Siberian town 1,300 miles (2,092 km) from Moscow in 1834. After his father died, Mendeleyev and his mother hitchhiked to the capital city so he could go to the University of Moscow. Unable to get a scholarship, they traveled another 400 miles (644 km) east to St. Petersburg. Soon after graduating first in his class, Mendeleyev went for postgraduate studies at Heidelberg University in Germany.

Because Germany at that time was home to many famous chemists, while Russia was something of a scientific backwater, when Mendeleyev returned to Russia to take up an academic position at the University of St. Petersburg, one of the first things he did was to write a 500-page chemistry textbook. It was a huge success both financially and intellectually, and Mendeleyev began to write another to cover the chemistry of all the elements.

He began the manuscript intending to write a chapter that discussed the chemical properties of each element. But it soon became apparent that he couldn't do so for each of the 60-odd elements known at the time. Furthermore, he needed some rational way to order the discussions. At the Karlsruhe Conference, Mendeleyev had become convinced that a key property of an element was its atomic weight. So he made a note card for each element and started arranging them on his desk in order of atomic weight, then grouped those that had similar chemical properties. (Legend has it that Mendeleyev liked to play the card game solitaire, in which cards end up in rows organized by both suit and numerical order, and that this experience helped him make his discovery.)

Others before Mendeleyev had noticed that some elements could be grouped into families. The German chemist Johann Wolfgang Döbereiner (1780–1849) had even pointed out that in triads such as lithium (7),

sodium (23), and potassium (39), the atomic weight of the middle member was about halfway between the atomic weights of the extremes. In a similar vein, British chemist John Newlands (1837–1898) proposed the law of octaves because he noted that, when arranged by atomic weight, the elements' chemical properties repeated themselves in periods of eight—just like the notes in a musical scale. Few paid much attention to this pattern, however, in part because it broke down at elements of higher atomic weight.

Mendeleyev was convinced that there was a periodic law that governed the properties of elements. Because mass played a central role in Isaac Newton's (1642–1727) laws, Mendeleyev believed that atomic weight determined each element's chemical properties, and that it was no coincidence that there were periodic similarities among them. (We now know that indeed it is not a coincidence—each row on the periodic table represents a filled shell of electrons; in Mendeleyev's time, however, no one knew anything about the atomic structure.) He was so convinced that a natural law underlay his table that he made several bold predictions based on it.

First, he predicted the existence of three new elements based on gaps in his table. Mendeleyev was convinced that there would be two new elements between zinc and arsenic because the properties of arsenic showed that it should be in the phosphorus family. Furthermore the jump in atomic weight was unusually large. Using the known properties of the missing elements' neighbors, Mendeleyev predicted not only their atomic weights, but their densities and melting points as well. When gallium was discovered a few years later, its actual properties were extremely close to Mendeleyev's predictions. When other gap elements were discovered in short order, chemists all over the world were convinced that the periodic table reflected a deep truth about the world.

Mendeleyev also used his table to predict that some current experimental values for atomic weight, such as the 128 initially assigned to tellurium (Te), were incorrect. If he had arranged the elements in that row strictly according to increasing atomic weight, iodine (I) with atomic weight 127 would have come before tellurium (128). But Mendeleyev knew that because of its chemical properties, iodine belonged in the same horizontal row with fluorine, chlorine, and bromine, while tellurium was more like its neighbor to the left, selenium. So he placed these elements where they belonged, based on their chemistry, and challenged chemists to redetermine the values of atomic weights.

In many cases, Mendeleyev's prediction that an atomic weight was a little off was correct. But the so-called reversed pair of tellurium and iodine remained an anomaly. Look up today's values for the atomic weights of tellurium (127.6) and iodine (126.904), remembering that we now put families such as fluorine, chlorine, bromine, and iodine in vertical columns while Mendeleyev put them in horizontal rows.

Chemists continued to find new elements. Some filled gaps in Mendeleyev's table but others didn't seem to fit at all. Helium was first discovered through spectrographic analysis of sunlight and later isolated from uranium ore. (We now know that uranium emits alpha particles, which are helium nuclei.) At first Mendeleyev didn't think there was a place for helium in his table, but after other inert gases such as argon and neon were found in the atmosphere, he simply inserted a whole new family into the table. These gases formed a useful buffer between the highly reactive alkali metals such as lithium and sodium and the halogen gases such as fluorine and chlorine gases; because they do not form compounds he assigned them a zero valence. But Mendeleyev was never able to figure out what to do with the so-called rare-earth metals, now called lanthanides.

According to their atomic weights, these metals should all come immediately after barium. But the problem was that each of these 14 elements have the same valence (+3) and very similar chemical properties. This violated the principle of periodicity on which Mendeleyev's table was based. He eventually concluded that the rare-earth metals “broke” the periodic law.

On the modern periodic table, both the lanthanides and actinides are displayed somewhat like a postscript below the table—even though according to atomic number they should appear above! Mendeleyev's contemporaries had no idea of how the rare earth's could fit into the pattern represented by the periodic table. We now know that the quantum number theory predicts their existence, and they do fit smoothly into the modern theory of the elements that is based on atomic number.

By the end of the nineteenth century, the very concept of chemical element that had been developed by Boyle, Lavoisier, and Dalton and that underlay Mendeleyev's system was to undergo profound changes. With the discovery of radioactivity, scientists learned that atoms were not the smallest pieces of matter. And the discovery of isotopes proved that not all the atoms of a given element were alike. Yet despite these revolutionary changes, the basic shape of the periodic table remained intact. Let us see how this was possible.

From Atomic Weight to Atomic Number

Mendeleyev first learned about cathode rays, x rays, and radioactivity a few years before his death in 1907. However, he was not very receptive to the news, in part because reports of glowing fluorescent screens caused by invisible rays reminded him of the false claims spiritualists at that time made about occult phenomena. Mendeleyev had, in fact, debunked the practices of several Russian mediums, who claimed that they could communicate with the dead, uncovering their deceptive and unscrupulous methods in the process.

Conclusive evidence soon emerged, however, that atoms consisted of positive particles and negative electrons. The New Zealand-born British physicist Ernest Rutherford (1871–1937) further showed in 1911 that the positive particles were concentrated in a nucleus. In 1913 his British student Frederick Soddy (1877–1956) found that not all atoms of a pure element were identical. For example, although a material then called mesothorium(because it was found in thorium ores), had an atomic weight of 228, its chemical properties were identical to those of radium, whose atomic weight was 226. (We now know that thorium undergoes radioactive decay to produce radium.) Soddy called atoms that were chemically identical but had different atomic weights “isotopes.” Using the mass spectrograph, scientists were able to separate all sorts of elements into their constituent isotopes.

These discoveries forced chemists to revise their definition of chemical elements. Boyle had stated that elements were the simplest substances out of which all materials were made, but it was now clear that the fundamental building blocks of matter were subatomic particles such as protons and electrons. Lavoisier had

defined elements as the last products of analysis, but scientists could now divide samples of what had been thought to be homogeneous materials into isotopes.

In 1920, in fact, British physicist Francis Aston (1877–1945) reported that gaseous neon consisted of Ne-20, Ne-21, and Ne-22 while chlorine consisted mostly of Cl-35 and Cl-37 with traces of Cl-36 and Cl-38. John Dalton would have been shocked to learn that all of an element's atoms do NOT have the same atomic weight. And if elements' identities did not depend on their atomic weights, then order by atomic weight could hardly provide the foundation for Mendeleyev's periodic table.

The solution, which may seem obvious today, is to define chemical elements by their atomic number. Chemical properties depend on the nucleus's positive charge and the corresponding number of electrons around it. Atoms of the same atomic number have different atomic weights if they have a different number of neutrons in their nuclei. Mendeleyev's periodic table worked as well as it did because there is a very close correspondence between order by atomic weight and order by atomic number. Sometimes the order is not the same (remember the reversed order of tellurium and iodine) but because nuclei that have many more neutrons than protons are unstable, Mendeleyev's idea of ordering elements by weight turned out to be a very close approximation of the true ordering principle: atomic number.

Some chemists balked at accepting this new definition, arguing that each isotope should be considered to be a separate element. After all, a uranium-235 atom is much more highly radioactive than a uranium-238 atom. There are also chemically significant differences between ordinary hydrogen H-1 and deuterium H-2. In fact deuterium oxide (so-called heavy water) is poisonous to animals. Following this logic, atomic weight would have remained the defining property.

The chemical community, however, wisely decided to use the atomic number. Can you imagine how complicated the periodic table would be if each individual isotope got its own box? Yet even with all this new information about atoms, one major puzzling feature of the periodic table remained: its periodicity.

Why do chemical properties repeat after every eight, or in some cases 18, elements? Why does the numerically ordered string of elements wrap so that if sodium (Na) is placed under lithium, where it obviously belongs, then chlorine automatically takes its place under fluorine? Why are some periods longer than others? And what about those pesky rare earth's, where the valence remained +3 even as the atomic number kept increasing? The answers were provided in the 1920s by quantum theory.

Quantum Numbers and the Periodic Table

When heated, each element gives off light with a characteristic spectrum. (If you sprinkle ordinary table salt into a blue flame you will see the characteristic reddish-orange color of sodium.) When this emitted light is analyzed with a spectroscope, the spectrum breaks into sharp lines. Because sunlight contained lines that did not correspond to any element known on Earth, spectroscopers realized that the sun contained a previously unknown element, which they named helium, from the Greek helios, for sun.

But why would only certain specific wavelengths be emitted? This and many other questions led physicists such as Niels Bohr (1885–1962) to conclude that the electrons surrounding the nucleus occupied orbits at certain discrete distances. When an individual atom absorbed energy, say from a flame, an electron in one of those shells jumped up to a new orbit that was farther away from the nucleus. When it eventually fell, it produced light of a single wavelength, determined by the energy difference between the two orbits. In the case of a very simple atom such as hydrogen, Bohr was even able to calculate from his theory what the spectrum of excited hydrogen should be.

According to quantum theory, the orbits available to any electron in an atom are described by integers called quantum numbers. In an unexcited atom (one that is not being bombarded by an outside energy source), the electrons gravitate to the orbits of lowest energy. But only two electrons can occupy a given orbit. If we apply the principles of quantum mechanics to the periodic table, we find that there can only be two elements in the first row because there is only space for two electrons in the first shell.

Quantum mechanics, however, tells us that there can be eight electrons in the second shell because it contains four available orbits—and we indeed find eight elements in the second row of the periodic table. Lithium (atomic number 3) has two inner electrons and one electron in its outer shell. In chemical reactions it may lose that electron and become a Li+1 ion. Fluorine (atomic number 9) has two inner electrons and seven outer electrons. Because a full outer shell is an especially stable arrangement (that is why the rare gases are inert), in reactions fluorine atoms often gain an electron and become negative F1 ions.

As we move to higher numbers of shells, describing the number of orbitals and their relative energies becomes more complicated, but the existence of horizontal families of elements (like the lanthanides and actinides) is predicted from the theory. All lanthanides have three electrons in their outermost shell (hence the +3 valence) but they differ in the number of electrons in the shell just beneath. Hence as the atomic number increases there is little variation in the properties of the outside electrons, which are those that typically enter into chemical bonding.

Modern Cultural Connections

Look on the walls of chemistry classrooms or laboratories in any country of the world and you will see a table of roughly 100 elements arranged to look something like a bed. Boxes containing the letters H (hydrogen) and He (helium) are stacked like bedposts on top of columns labeled as I and VIII, respectively. The atomic number of each element tells how many protons are in the nucleus; the Roman numeral of the column is an indication of the valence (or combining power) of the element. Since lithium is in column I (under hydrogen) but beryllium is in column II, then we expect the formula of lithium chloride to be LiCl while beryllium chloride is BeCl2. We also expect family resemblances between elements in the same column. For example, in column II a couple of squares under Be we find Ca for calcium and, right under it, Sr for strontium. Knowing that calcium is an important component of bones, we aren't surprised to hear that its neighbor strontium is also found in the skeleton. This family relationship explains why during the period of atomic bomb testing, scientists were very concerned that a radioactive isotope of strontium (called Sr-90) might end up in milk just like calcium and be absorbed into the bones of growing children.

Primary Source Connection

A World made by Atomes

Small Atomes of themselves a World may make, As being subtle, and of every shape: And as they dance about, fit places finde, Such Formes as best agree, make every kinde. For when we build a house of Bricke, and Stone, [5]

We lay them even, every one by one: And when we finde a gap that's big, or small, We seeke out Stones, to fit that place withall. For when not fit, too big, or little be, They fall away, and cannot stay we see. [10] So Atomes, as they dance, finde places fit, They there remaine, lye close, and fast will sticke. Those that unfit, the rest that rove about, Do never leave, untill they thrust them out. Thus by their severall Motions, and their Formes, [15]

As severall work-men serve each others turnes. And thus, by chance, may a New World create: Or else predestined to worke my Fate.

Margaret Cavendish

cavendish, margaret. the atomic poems of margaret (lucas) cavendish, duchess of newcastle, from her poems, and fancies, 1653. tr anscribed and edited by leigh tillman partington. atlanta: women writers resource project, emory university, 1996.

See Also Chemistry: Chemical Bonds; Chemistry: Chemical Reactions and the Conservation of Mass and Energy; Chemistry: Molecular Structure and Stereochemistry; Chemistry: States of Matter: Solids, Liquids, Gases, and Plasma; Chemistry: The Practice of Alchemy.

bibliography

Books

Ball, Philip. The Elements: A Very Short Introduction. Oxford and New York: Oxford University Press, 2004.

Ball, Philip. The Ingredients: A Guided Tour of the Elements. Oxford and New York: Oxford University Press, 2002.

Boyle, Robert. The sceptical chymist; or, Chymicophysical doubts & paradoxes, touching the experiments whereby vulgar spagirists are wont to endeavour to evince their salt, sulphur and mercury, to bethe true principles of things. To which in this editionare subjoyn'd divers experiments and notes aboutthe producibleness of chymical principles. Oxford: 1680.

Brock, William H. The Norton History of Chemistry. New York and London: W.W. Norton & Co., 1992.

Cavendish, Margaret. The Atomic Poems of Margaret(Lucas) Cavendish, Duchess of Newcastle, from her Poems, and Fancies, 1653. Transcribed and edited by Leigh Tillman Partington. Atlanta: Women Writers Resource Project, Emory University, 1996.

Dalton, John. A New System of Chemical Philosophy. New York: Philosophical Library, 1964.

Gordin, Michael D. A Well-Ordered Thing: Dmitrii Mendeleev and the Shadow of the Periodic Table. New York: Basic Books, 2004.

Lavoisier, Antoine Laurent. Elements of Chemistry. New York: Dover Publications, 1965.

Mendeleyev, Dimitri. Principles of Chemistry by D. Mendeléeff. Translated from the Russian by George Kamensky. 6th ed. Edited by T.A. Lawson. London, New York: Longmans, Green, and Co., 1897.

Morris, Richard. The Last Sorcerers: The Path from Alchemy to the Periodic Table. Washington, DC: Joseph Henry Press, 2003.

Strathern, Paul. Mendeleyev's Dream: The Quest for the Elements. New York: St. Martin's Press, 2001.

Web Sites

AUS-e-TUTE. “History of the Periodic Table of the Elements.” http://www.ausetute.com.au/pthistor.html (accessed June 8, 2006).

Chemical Heritage Foundation. “Chemical Achievers: The Human Face of the Chemical Sciences.”http://www.chemheritage.org/classroom/chemach/index.html/ (accessed June 8, 2006).

Royal Society of Chemistry. “History of the Periodic Table.” http://www.chemsoc.org/networks/learnnet/periodictable/pre16/develop/mendeleev.htm (accessed June 8, 2006).

Noretta Koertge

Cite this article
Pick a style below, and copy the text for your bibliography.

  • MLA
  • Chicago
  • APA

"Chemistry: The Periodic Table." Scientific Thought: In Context. . Encyclopedia.com. 16 Nov. 2018 <https://www.encyclopedia.com>.

"Chemistry: The Periodic Table." Scientific Thought: In Context. . Encyclopedia.com. (November 16, 2018). https://www.encyclopedia.com/science/science-magazines/chemistry-periodic-table

"Chemistry: The Periodic Table." Scientific Thought: In Context. . Retrieved November 16, 2018 from Encyclopedia.com: https://www.encyclopedia.com/science/science-magazines/chemistry-periodic-table

Learn more about citation styles

Citation styles

Encyclopedia.com gives you the ability to cite reference entries and articles according to common styles from the Modern Language Association (MLA), The Chicago Manual of Style, and the American Psychological Association (APA).

Within the “Cite this article” tool, pick a style to see how all available information looks when formatted according to that style. Then, copy and paste the text into your bibliography or works cited list.

Because each style has its own formatting nuances that evolve over time and not all information is available for every reference entry or article, Encyclopedia.com cannot guarantee each citation it generates. Therefore, it’s best to use Encyclopedia.com citations as a starting point before checking the style against your school or publication’s requirements and the most-recent information available at these sites:

Modern Language Association

http://www.mla.org/style

The Chicago Manual of Style

http://www.chicagomanualofstyle.org/tools_citationguide.html

American Psychological Association

http://apastyle.apa.org/

Notes:
  • Most online reference entries and articles do not have page numbers. Therefore, that information is unavailable for most Encyclopedia.com content. However, the date of retrieval is often important. Refer to each style’s convention regarding the best way to format page numbers and retrieval dates.
  • In addition to the MLA, Chicago, and APA styles, your school, university, publication, or institution may have its own requirements for citations. Therefore, be sure to refer to those guidelines when editing your bibliography or works cited list.