Everything in Its Place - THE BIOLOGY ROOM



Everything in Its Place

Oliver Sacks story about the history of the periodic table highlights some of the important perspectives and approaches that characterize how knowledge is created in Science.

1. Read the story and using examples from the text describe the following aspects of the scientific approach or method.

a) Use of empirical evidence – observations made using the five senses

b) Classification of the natural world

c) Scientific knowledge can be both reliable and tentative and both can exist together simultaneously.

d) Creativity, imagination and risk taking

e) Scientists working collaboratively and independently

f) Production of laws

g) Production of theories

h) Use of inductive and deductive reasoning

i) New knowledge being produced both gradually through evolutionary changes and during revolutions or suddenly.

j) The production of new knowledge is dependent on the development of new technologies.

2. Are there any important aspects of a scientific method that the article does not mention or emphasize?

Everything in Its Place

One man's love affair with the periodic table.

By OLIVER SACKS

New York Times, 1999

| | |

| |John Dalton, a Quaker teacher and the first to |

| |assign atomic weights to elements, created this |

| |table in 1808. |

| |Photomontage by Amy Guip |

| |[pic] |

It used to be said, when I was a boy, that there were 92 elements, each with its own unique characteristics. These elements, which could combine with one another to form millions of compounds, were "the building blocks of the universe."

One knew, or suspected, that some of them were related. Tin and lead, for example, were both soft metals, easily melted; copper, silver and gold -- the "coinage" metals -- could all be beaten into foils so thin that they transmitted green or blue light.

But I am not sure that it occurred to me that all the elements might be related to one another until I went, at the age of 12, to the Science Museum in London (newly reopened after the end of the Second World War) and there saw an enormous cabinet labeled "The Periodic Table" hanging at the head of the stairs. Seeing the table, with its actual samples of the elements, was one of the formative experiences of my boyhood and showed me, with the force of revelation, the beauty of science. The periodic table seemed so economical and simple: everything, the whole 92-ishness, reduced to two axes, and yet along each axis an ordered procession of different properties.

Chemistry started to emerge from its alchemical roots in the 18th century, partly with the discovery of new elements: between 1735 and 1826, no fewer than 40 were added to the 9 known to the ancients (copper, silver, gold, iron, mercury, lead, tin, sulphur and carbon) and the few discovered in the Middle Ages (arsenic, antimony and bismuth). The discovery of these new elements forced certain questions on every chemist: How many elements were there? Was there any limit to their number? Were they all related somehow? And if so, how could they be classified?

Kinships were recognized among some. Chlorine, bromine and iodine -- all colored, volatile, hungrily reactive -- seemed a natural family, the halogens. Calcium, strontium and barium, the alkaline earth metals, were another family, for they were all light, soft, readily set alight and strongly reactive with water.

In 1817, a German chemist, Johann Dobereiner, observed that the atomic weights (now called atomic mass) of the alkaline earth metals formed a series, the atomic weight of strontium being just midway between those of calcium and barium. He later discovered other such triads, as well as triads in which the elements had similar properties but almost identical atomic weights.

Dobereiner's triads convinced many chemists that atomic weight must represent a fundamental characteristic of all elements. But confusion about the basics remained -- about the difference between atoms and molecules and about the combining power, or valency, of atoms. As a consequence, many accepted atomic weights were wrong. Dalton himself -- the originator of the atomic hypothesis -- assumed, for instance, that the formula of water was HO and not H2O, giving him an atomic weight for oxygen that was only half the correct number.

In 1860, the first international gathering of chemists was convened at Karlsruhe, Germany, for the express purpose of clearing up this confusion. Here, Stanislao Cannizzaro proposed a reliable way of calculating atomic weights from vapor density, and his beautifully argued presentation carried the day, leading to a consensus: now, at last, with corrected atomic weights and a clear idea of valency, the way was open for a comprehensive classification of the elements.

It is a remarkable example of synchronicity that no fewer than six such classifications, all pointing toward the discovery of periodicity, were independently devised in the next decade. Of these, Dmitri Ivanovich Mendeleev's system was the most comprehensive, and also the most audacious, for it ventured to make detailed predictions of elements as yet unknown.

| |[pic] |

| |John Dalton, a Quaker teacher and the first |

| |to assign atomic weights to elements, |

| |created this table in 1808. |

| |[pic] |

Mendeleev (whose name and wild bearded face were known to every school age child of my time) was a figure of heroic proportions. He was Russia's chief scientific adviser and closely involved with industry and agriculture, from coal and oil to cheese and beer. He was the author of the most delightful and vivid chemistry text ever published, "The Principles of Chemistry," and he had brooded since 1854 on how the chemical elements might be classified.

With the old, pre-Karlsruhe atomic weights, one could get, as Dobereiner did, a sense of local triads, or groups. But one could not easily see that there was a numerical relationship between the groups themselves. Only when Cannizzaro showed that the proper atomic weights for the alkaline earth metals, calcium, strontium and barium, were 40, 88 and 137 did it become clear how close these were to those of the alkali metals, potassium (39), rubidium (85) and cesium (133). It was this closeness, and the closeness of the atomic weights of the halogens -- chlorine, bromine and iodine -- that incited Mendeleev in 1868 to make a small, two-dimensional grid juxtaposing the three groups:

|Cl |35.5 |K |39 |Ca |40 | |

|Br |80 |Rb |85 |Sr |88 | |

|I |127 |Cs |133 |Ba |137 | |

And it was at this point, seeing that arranging the three groups of elements in order of atomic weight produced a repetitive pattern -- a halogen followed by an alkali metal followed by an alkaline earth metal -- that Mendeleev felt this must be a fragment of a larger pattern and leapt to the idea of a periodicity governing all the elements, a periodic law.

Mendeleev's first small table had to be filled in and then extended in all directions, as if filling up a crossword puzzle. Moving between conscious calculation and hunch, between intuition and analysis, Mendeleev arrived within a few weeks at a tabulation of 30-odd elements in order of ascending atomic weight, a tabulation that suggested that there was a recapitulation of properties with every eighth element.

On the night of Feb. 16, 1869, it is said, Mendeleev had a dream in which he saw almost all of the 65 known elements arrayed in a grand table. The following morning, he committed this to paper.

This first table was to undergo considerable revision over the next few years, but by 1871 it had taken its new familiar form of a chunky rectangle with intersecting groups and periods.

It was this table that I saw in the Science Museum and that was to be found in every textbook, lecture room and museum for a century. One could read the table up and down, going from one group to another (each vertical group was a family of elements with similar reactivity and valency) -- this was what Dobereiner and the pre-1860 chemists would have done. But one could also read it horizontally, getting a feel for each period as it moved through the eight groups. One could see the way in which the properties of the elements changed with each increment of atomic weight, until suddenly the period came to an end and one found oneself on the next period, where all the elements echoed the properties of those above. It was this, above all, that gave one a feel for the mysterious periodicity of the table, the reality of the great law it enshrined.

already had a little lab of my own, where I had spent many hours, and I must have seen in books small versions of Mendeleev's table. But it was seeing the huge table in the museum, being enraptured, really assimilating it for the first time, that moved me from a random or encyclopedic approach -- collecting all the chemicals I could, doing all the experiments I could -- to a more systematic one, exploring the trends of the elements for myself.

| | |

|One simple (and slightly | |

|dangerous) experiment was putting| |

|lumps of the alkali metals into | |

|water. Potassium would catch | |

|fire; rubidium was still more | |

|reactive, and cesium exploded | |

|when it hit the water, shattering| |

|its glass container. One never | |

|forgot the properties of the | |

|alkali metals after this. | |

One simple, highly dramatic (and slightly dangerous) experiment was putting small lumps of the alkali metals into water and seeing how they increased in reactivity as their atomic weight increased. One had to do this gingerly, with tongs, and to equip oneself and one's guests with goggles: lithium would move about the surface of the water sedately, reacting with it, emitting hydrogen, until it was all gone; a lump of sodium would move around the surface with an angry buzz, but would not catch fire if a small lump was used; potassium, in contrast, would catch fire the instant it hit the water, burning with a pale mauve flame and shooting globules of itself everywhere; rubidium was still more reactive, spluttering violently with a reddish violet flame, and cesium, I found, exploded when it hit the water, shattering its glass container. One never forgot the properties of the alkali metals after this.

The periodic table did not actually tell one the properties of the elements, but like a family tree, it assigned them places. The fun, for me, was to work backward from this, to see how an element's properties corresponded with its place. Tungsten, for example, was a favorite -- no other metal had such a high melting point. This, I first thought, made it unique, but now I could compare it with its neighbors in the periodic table and see that the highest metallic melting points were all to be found in Period 6. Tungsten, lay at the intersection of two mountain ranges, an Everest among other peaks, but not an anomaly.

I could plot the physical and chemical properties of all the elements against their atomic weights and obtain the most tantalizing graphs. If one plotted atomic volume against atomic weight, for example, one would get a many-peaked curve, with summits for the light Group I metals. Every property, it seemed, varied periodically and was somehow linked with atomic weight. But why any of the elements should have the properties they had, and why such properties should recur in periodicity with atomic weight, were complete mysteries to me, as they had been to Mendeleev.

From 1869 to 1871, Mendeleev expanded the table, going so far as to reposition elements that did not fit, revising their accepted atomic weights to make them fit, an act that shocked some of his contemporaries. Further challenges were presented by two groups of elements, the transition elements (these included rare metals like vanadium and platinum, as well as common ones like iron and nickel) and the rare-earth elements. Neither of these seemed to fit in the neat octaves of the earlier periods. To accommodate them, Mendeleev and others experimented with new forms of the table -- helical forms, pyramidal forms, etc. -- that, in a sense, gave it extra dimensions.

In an act of supreme confidence, Mendeleev reserved several empty spaces in his table for elements "as yet unknown." He asserted that by extrapolating from the properties of the elements above and below (and also, to some extent, from those to either side), one might make a confident prediction as to what these unknown elements would be like. He did exactly this, predicting in great detail a new element that would follow aluminum in Group 3: it would be a silvery metal, he thought, with a density of 6.0 and an atomic weight of 68. Four years later, in 1875, just such an element was found: gallium. He also predicted with equal precision the existence of scandium and germanium, and these too were soon discovered. It was this ability to predict elements in such detail that stunned his fellow chemists and convinced many of them that Mendeleev's system was not just an arbitrary ordering of the elements but a profound expression of reality.

But Mendeleev was astonished, as everyone was, by the discovery in the 1890's of an entire new family of elements, the inert gases. He was at first skeptical of their existence. (He initially thought that argon, the first found, was just a heavier form of nitrogen.) But with the discovery of helium, neon, krypton, xenon and finally radon, it was clear that they formed a perfect periodic group. They were identical in their inability to form compounds; they had a valency, it seemed, of zero. So to the eight groups of the table, Mendeleev now added a final Group 0.

With the inert gases in place, the number of elements in each period stood out: 2 (hydrogen and helium) in the first period; 8 each in the second and third; 8 typical plus 10 transition elements, or 18 each, in the fourth and fifth periods; 8 plus 10 plus 14 rare-earth elements, or 32, in the sixth period. These were the magical numbers -- 2, 8, 8, 18, 18, 32. But what did they mean? And what, in broader terms, was the basis of chemical properties?

Mendeleev constantly returned to these questions. He yearned for a new "chemical mechanics," comparable to the classical mechanics of Newton. And yet one wonders what he might have thought of the actual form of the revolution that took place after his death, a revolution wholly unimaginable in terms of classical mechanics.

The new insight into the internal constitution of atoms came in 1911, four years after Mendeleev's death, when Ernest Rutherford (bombarding gold foil with alpha particles and finding that, very occasionally, one was deflected back) inferred that the atom must have a structure like a miniature solar system, with almost all of its mass concentrated in a minute, very dense, positively charged nucleus surrounded at great distances by relatively weightless electrons. But the very essence of atoms was their absolute stability. And such an atom as Rutherford's, if ruled by the laws of classical mechanics, would not be stable; its electrons would lose energy as they orbited, eventually diving into the nucleus.

Niels Bohr, working with Rutherford in 1912, was intensely aware of this, and of the need for a radically new approach. This he found in quantum theory, which postulated that electromagnetic energy -- light, radiation -- was not continuous but emitted or absorbed in discrete packets, or "quanta." Bohr, by an astounding leap, connected these concepts with the Rutherford model and with the well-known but previously inexplicable nature of optical spectra -- that these were not only characteristic for each element but consisted of a multitude of discrete lines or frequencies.

All of these considerations came together in the Bohr atom, where electrons were conceived to occupy a series of orbits, or "shells," of differing energies about the nucleus. Unlike classical orbits, which decay, these quantum orbits had a stability that allowed them to maintain themselves, potentially, forever. (But if the atom was excited, some of its electrons might leap to higher energy orbits for a while and in returning to their ground state emit a quantum of energy of a certain frequency; it was this that caused the characteristic lines in their spectra.)

Bohr brought out his model of the atom in the spring of 1913. A few months later, Henry Moseley found a most intimate relationship between the order of the elements and their X-ray spectra. These spectra could be correlated, Moseley thought, with the number of positive charges in the nucleus, and for this the term "atomic number" was used. With atomic numbers, there were no gaps or fractions or irregularities, as with atomic weights. It was atomic number, not atomic mass, that determined the order of the elements. And Moseley could now say with absolute confidence that there were only 92 elements between hydrogen and uranium, including half a dozen as yet undiscovered. (Three of these had been predicted, though vaguely, by Mendeleev.)

Bohr's model suggested that every element's chemical properties, its position in the periodic table, depended on the number of its electrons and how these were organized in successive shells. Valency and chemical reactivity, the definers of Mendeleev's groups, were correlated with the number of valence electrons in the outer shells: with the maximum of eight electrons, an atom was chemically inert; with more, or less, than the maximum, it would tend to be more reactive. Thus the halogens, only one electron short in their outermost shells, were avid to pick up an eighth electron, whereas the alkali metals, with only a single electron in their outer shells, were avid to get rid of it, to become stable in their own way.

To this basic eightness, extra shells were added in the later periods: 10-electron shells for the transition elements and 14-electron shells for the rare-earth elements.

Bohr and Moseley provided a spectacular confirmation of the periodic table, grounding it, as Mendeleev had hoped, in "the invisible world of chemical atoms." The periodicity of the elements, it was now clear, emerged from their electronic structure. And the mysterious numbers that governed the periodic table -- 2, 8, 8, 18, 18, 32 -- could now be understood as the number of electrons added in each period.

Such an electronic periodic table is basically identical with Mendeleev's table, posited nearly half a century earlier on purely chemical grounds. Moseley and Bohr worked from the inside, with the invisible world of chemical atoms, and Mendeleev and his contemporaries worked from the outside, with the visible and manifest properties of the elements -- and yet they arrived at the same point. This is the beauty of the periodic table, indeed, that it looks both ways, uniting classical chemistry and quantum physics in a magical synthesis.

Given Bohr's orbits of different energy levels, one can, in principle, build up the whole periodic table by adding electrons one at a time, climbing the rungs of an atomic ladder from helium to uranium. And it is by such a building-up that we have been able to create new elements absent in nature, like the 20 elements (93-112) that now follow uranium in the periodic table, heavier atoms that do not depart from the regularities of the periodic law. In principle, one can work out the periodic table to element 200 and beyond and predict some of the properties of such elements. (These predictions are largely theoretical, because the highly radioactive transuranic elements tend to get more and more unstable. One may only be able to produce an atom at a time, and this may be gone in a few millionths of a second.)* But the idea of periodicity, it seems, has no discernible limits, and this, like all the confirmations of this century, would have delighted Mendeleev.

It is more than 50 years since I first saw the periodic table, and my delight in it has never faded. It is still the icon of chemistry, as it has been for 130 years; it continues to guide chemical research, to suggest new syntheses, to allow predictions of the properties of never-before-seen materials. It is a marvelous map to the whole geography of the elements.

My kitchen is papered with periodic tables of every size and sort -- oblongs, spirals, pyramids, weather vanes -- and on the kitchen table, a very favorite one, a round periodic table made of wood that I can spin like a prayer wheel. I carry two tiny periodic tables in my wallet -- a classical Mendeleevian one with antique lettering and a more modern one, a beautiful colored spiral that shows the elements, their atomic numbers, like a great nebula, whirling out beyond uranium to who knows what infinity.

Oliver Sacks was born in 1933 in London into a family of doctors and scientists. He earned his medical degree at Oxford University. Since 1965, he has lived in New York, where he is a practicing neurologist. Since July 2007 he has been Professor of Clinical Neurology and Clinical Psychiatry at Columbia University.

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