SEEKING PATTERNS
-Jack R. Holt
A SEMINAR SPEAKER
Thou cunning'st pattern of excelling nature…
-William Shakespeare (Othello act IV, sc 2)
Sometimes a teacher has the thrill of having
a student return as a teacher. I had such an
experience at the beginning of this month when
David Seaborn returned to give a seminar about
his research on algal ecology in Virginia. In the
introduction to his presentation, Dave said that
scientists seek patterns in nature. The statement
struck me, and I became lost in thought over it
for a short time.
At first I considered whether all areas of
science met that criterion. Maybe a search for
patterns was just characteristic of ecology, but
did it apply to most other areas of science? I
considered those ideas for a time and searched
my mind for examples. I returned from that
limbo when Dave changed to the next slide and
began to present patterns that he had discovered
in his research. Of course I had forgotten a pad
to take down notes and spied a loose piece of
paper under my desk. As I checked to see if it
would be an appropriate piece of scratch paper, I
turned it over and saw the characteristic rows
and columns of the Periodic Table of the
Elements (see Figure 1). There in my hand was
confirmation of the generality of Dave's
assertion. The patterns of boxes on that table
illustrated fundamental relationships between
elements and provided the framework to
understand and predict deeper, subtler patterns of
the ultimate constituents of matter.
A WATERSHED
Chemistry has reached a state of development
when…a meeting of a great number of
chemists…be held so that a unification of a few
important points shall be approached. -Carl
Welzein (1860; excerpt from invitation letter to
Karlsruhe International Chemical Congress)
The origin of the Periodic Table can be
traced back to Lavoisier and before. However,
several critical events occurred around 1860 that
make it a watershed year in the history of
chemistry. Recall that terms, symbols, atomic
weights, etc. had not been standardized.
Communication from one laboratory to the next
was a mess. Textbooks attempted to standardize
communication but general agreement about
terminology was hotly contested. Kekule (18261896; see Chance and the Prepared Mind)
suggested that a general conference of chemists
might come to some agreement. Carl Welzein
(1813-1870) sent out a call for an International
Chemical Congress to convene in Karlsruhe,
Germany. The meeting drew 127 of Europe's
leading chemists to discuss these issues. The
purpose of the meeting was to create a
"unification of a few important points" in organic
as well as inorganic chemistry. Most of all, the
group sought to standardize the way in which
atomic weights were calculated and expressed.
FIGURE 1.
A modern version of the Periodic Table of the Elements.
http://periodictable.com
Image from:
Recall that the concept of atomic weights
was as old as Dalton's Atomic Theory (see Old
Books and Atoms). The atoms of each element
seemed to have a number of unique and
consistent properties and atomic weight was
among them. Atoms were too small to measure,
but techniques of measuring the mass of one
element relative to another had become
established with Dalton and perfected with
Berzelius. Still, there were problems. Atomic
weights had to be determined by the combining
weights of simple compounds. For example,
water when separated by electrolysis, yielded
two volumes of hydrogen to one volume of
oxygen. The same ratio applied when combining
oxygen and hydrogen to make water. Thus,
water seemed to be one atom of oxygen and two
atoms of hydrogen. The oxygen per unit volume
weighed 16X that of each unit volume of
hydrogen.
Therefore, assuming that equal
volumes of gas contained equal numbers of
particles, an atom of oxygen was 16X more
massive than an atom of hydrogen.
The
combining weights of other elements relative to
known atomic weights gradually allowed
chemists to fill in the list of known elements, a
list that had grown to 54 by 1859. Particular
problems arose, however, when calculating some
relative weights because gasses like hydrogen
seemed somewhat capricious.
FIGURE 2. Two explanations for the formation
of HCl. The top equation follows prevailing pre1860's theory. The bottom equation follows
Avagadro's hypothesis and experimental
observation.
Consider the simple gasses, hydrogen and
chlorine. When combined, they make another
gas, hydrogen chloride (recall that we used HCl
dissolved in water to make out iron and zinc
salts). The odd thing is that one volume of
hydrogen combined with one volume of chlorine
makes two volumes of hydrogen chloride. How
is that possible if each equal volume of gas has
equal numbers of particles?
In 1811 an Italian chemist named Amadeo
Avagadro (1776-1856) proposed that the gaseous
elements combine as diatomic molecules. Thus,
hydrogen gas existed as H2, a diatomic molecule.
That would explain the HCl problem. More
importantly, if most atomic weights had been
calculated relative to a gas like hydrogen, then
relative weights would be half the size of atomic
weights.
Nevertheless, the influence of
Berzelius (see Old Books and Atoms) was
considerable, and, in his view, diatomics were
impossible.
Several of the attendees held the diatomic
theory as key to a true system of molecular
weights. None spoke as passionately or as
convincingly as fellow Italian Stanislao
Cannizzaro (1826-1910). He addressed the
congress and demonstrated with logic and data
how Avagadro's assumption simplified the
problem.
Although there was no general
agreement at Karlsruhe, Cannizzaro brought
copies of his papers and passed them out. As a
consequence, the Avagadro assumption of
diatomic molecules became generally accepted
with a concomitant standardization of atomic
weights.
OBSERVATION OF SPECTRA
Spectrum analysis…offers a wonderfully simple
means for discovering the smallest traces of
certain elements in terrestrial substances…and
even of the solar system.
-Kirchhoff & Bunsen (1860)
The year 1860 also saw the development of
the spectroscope as a tool for the investigation
and discovery of elements. Gustav Robert
Kirchhoff (1824-1887) and Robert Wilhelm
Bunsen (1811-1899) began to study how
different elements glow when heated in a flame.
Kirchhoff decided to spread the spectrum of light
with a prism. Then, he and Bunsen discovered
that bright lines in the spectra of glowing
elements seemed to be characteristic of each
element. So, they documented patterns of
spectral lines in sodium, lithium, potassium,
strontium, calcium, and barium. By 1863,
Bunsen, Kirchhoff and others identified cesium,
rubidium, thallium, and indium as new elements
by using the spectroscopic method.
The spectroscope required a relatively
invisible flame to work well. Bunsen designed
such a gas burner, and even now, the Bunsen
Burner has remained a common piece of
equipment in most laboratories. Through the
decade of the 1860's, standard tables of atomic
weights and new tools such as the spectroscope
led to the identification of more elements and the
creation of more confusion. Finally the chaos
began to clear as chemists started to order the
elements according to standard atomic weights
and look for patterns.
FIGURE 3. Illustration of a spectroscope from
the paper by Kirchhoff and Bunsen (1860).
THE PERIODIC LAW
If all the elements be arranged in order of their
atomic weights, a periodic repetition of
properties is obtained.
-Dmitri I. Mendeleev (1869)
Dmitri Ivanovich Mendeleev (1834-1907),
through the efforts of his mother, made his way
from Siberia to St. Petersburg where was taught
by
von
Liebig's
student,
Alexander
Woskressensky (1809-1880). Thus, Mendeleev
began his career in science following the
footsteps of von Liebig (see Saturday Scientist
vol 8, no. 3) with original work in organic
chemistry. Soon, however, he turned to a group
of similar elements: tungsten, osmium,
vanadium, and iridium.
On the recommendation of chemist and
composer Alexander Borodin, Mendeleev was
sent by the Tsar to Germany for advanced study.
While there, he attended the Karlsruhe Congress
and began to consider how elements might be
ordered according to their atomic weights. On
his return to Russia, Mendeleev wrote a textbook
of organic chemistry in only 7 months. During
that time he speculated about the relationship
between atomic weight and elemental properties.
FIGURE 4. Spectra for hydrogen, helium, sodium, and mercury.
By 1867 he began to write Osnovy Khimi
(Foundations of Chemistry). Then, he struggled
with ways to present the elements in related
groupings. After he published his first edition of
Foundations, he noticed that the pattern of
elemental properties repeated according to the
series of atomic weights.
FIGURE 5. Dmitry I. Mendeleev.
He was not the first to notice that some
elements seemed to occur in families. For
example, chlorine, bromine and iodine all tended
to ionize with a charge of -1. They formed
similar kinds of acids and similar salts. They
made up the halogen family. On the other hand,
lithium, sodium, and potassium comprised a
group that tended to make ions of +1 charge.
They made similar bases and similar salts. Other
families suggested themselves as well.
He published his first periodic table in 1869.
I have reproduced a modification of that table as
Table 1. Nevertheless, it was his first attempt and
illustrated in a very powerful way Mendeleev's
Periodic Law. An important attribute to the table
was that blank spaces (Mendeleev indicated
them by question marks) appeared in the pattern.
Thus, Mendeleev assumed that they represented
undiscovered elements. However, because they
were in families of elements, he could predict
their properties and, thereby, systematize a
search for them. Mendeleev predicted properties
of some unknown elements that were uncanny in
their accuracy, and, when those elements were
finally discovered, helped to establish the
Periodic Law.
TABLE 1. A modification of Mendeleev's first
Periodic Table of the Elements. Note that the
orientation of the periodic table is on its side
relative to Figure 1. Mendeleev predicted
particular properties for unknown elements
indicated by ?.
H=1
Li=7
Na=23
K=39
Rb=85.4
Be=9.4
Mg=24
Ca=40
Sr=87.6
?=45
?
Ti=50
Zr=90
V=51
Nb=94
Cr=52
Mo=96
Mn=55
Rh=104.4
Fe=56
Ru=104.4
Ni=Co=59 Pd=106.6
Cu=63.4
Ag=108
Zn=65.2
Cd=112
B=11
Al=27.4 ?=68
Ur=116
C=12
Si=28
Sn=118
?=70
N=14
P=31
As=75
Sb=122
O=16
S=32
Se=79.4
Te=128?
F=19
Cl=35.5
Br=80
I=127
Lothar Meyer (1830-1895), a German and
another attendee of the Karlsruhe Congress,
began to write a general chemistry text in the late
1860's. Like Mendeleev, he attempted to order
the elements and discovered that they exhibited
patterns of periodicity in their properties. He
ordered them according to atomic sizes (this can
be estimated by dividing the atomic weight by
the density of the element). He would have
scooped Mendeleev if he had published right
away, but Meyer delayed the release of his book
by several years.
THE LAZY ELEMENTS
The unreactivity of the noble gas elements
belongs to the surest of experimental results.
-Friedrich Paneth (1924)
That the elements exhibited a pattern of
predictable properties was somewhat mysterious.
Why should they behave that way? What did the
spectral lines indicate? That elements showed
such patterns in their spectra suggested to many
that the atoms had simpler constituent parts and
that the properties of elements just reflected the
way that they were constructed. William Prout
(1785-1850) had proposed such a hypothesis
early in the 19th Century. He suggested that all
elements were made of multiples of hydrogen,
the simplest element. Berzelius ridiculed the
idea and coined the phrase "Prout's Hypothesis"
as a pejorative term. Still, the mystery of the
periods could be allayed if some underlying
building blocks could be discovered.
In the search for evidence to support Prout's
Hypothesis or a modification of it, Lord
Rayleigh (1842-1919) made ever more accurate
determinations of the atomic weights of
atmospheric gasses. By 1892, he had determined
that oxygen was 15.882 times heavier than
hydrogen.
However, when he measured
nitrogen, Rayleigh found that atmospheric
nitrogen was denser than nitrogen generated
from ammonia. This strange result suggested
that nitrogen might exist in forms other than
diatomics in the air. At this point, William
Ramsay (1852-1916) took up the investigation
and removed all nitrogen from the air. What was
left seemed to have an atomic weight of 40 and
had a spectroscopic signature of an unknown
element. This new element refused to combine
with anything and seemed to be monatomic. He
called the new element Argon, Greek for lazy. It
seemed to fall between Chlorine and PotassiumCalcium in the periodic table. He figured that it
represented a whole new class of elements.
Soon, Ramsay discovered helium (an element
whose spectroscopic signature had been
observed in the solar spectra). Mendeleev, at
first was skeptical of the unreactive or Noble
Gasses, but added them to his table in 1905, just
two years before his death.
The end of the 19th Century saw revolutions
in chemistry and physics.
Discoveries of
mysterious radioactive elements supported
fundamental new insights as to the view that
atoms had parts, a concept that Mendeleev never
accepted.
ENTER THE PHYSICISTS
Now Rutherford has proved that the most
important constituent of an atom is its central
positively charged nucleus…
-Henry Moseley (1913)
Joseph
John
Thomson
(1856-1940)
succeeded Rayleigh at the Cavendish
Laboratory, Cambridge. His investigations led
to the discovery of negatively charged particles
that emanated from an electrode in a vacuum.
He measured the mass of the particles and
estimated that they were 1000 times smaller than
a hydrogen atom. He supposed that atoms could
be made of such component parts. Later, the
cathode ray particle was renamed the electron.
Ernest Rutherford (1871-1937) replaced
Thomson at the Cavendish and continued
research into the constituents of atoms. He
discovered, through a series of brilliant
experiments, that the atom's positive charge (and
most of its mass) was concentrated into a very
tiny volume in the center of the atom. Thus, he
conceived of an atom as a small planetary system
with a small massive positive nucleus at the
center and the electrons in orbit around it.
A student of Rutherford, Henry Moseley
(1887-1915) examined X-ray signatures of the
elements. He noticed that the frequencies of
emitted X-rays increased according to the
placements of elements in the periodic table. He
supposed that his results confirmed the stepwise
increase in positive charge of the nucleus. Thus,
he reckoned that atomic weight simply was a
function of the positive charges.
As a
consequence, he created the concept of atomic
number.
FIGURE 6. Henry Moseley.
For example, hydrogen has an atomic
number of 1. It has one positively charged
particle or proton in its nucleus. Therefore, a
neutral hydrogen atom has one proton and one
electron. The electrons can vary, but the number
of protons determines the nature of each element.
Moseley examined the X-ray signatures of
most of the known elements before he joined the
armed forces during World War I. Unfortunately
for chemistry Moseley's life was cut short by a
sniper's bullet at Gallipoli. However, during his
brief existence, he fundamentally changed the
periodic table and its meaning.
Other physicists like Niels Bohr (18851962) considered the problem of spectroscopic
lines and the planetary model of the Rutherford
nucleus. Bohr was troubled by the problems
posed by such a model. Particularly, what kept
the electrons from crashing into the nucleus and
staying there? He supposed that electrons could
occupy only certain allowed orbits, and, as they
gained energy, electrons occupied only orbits of
particular higher energies farther from the
nucleus and then dropped back down, thereby
releasing that energy as light. The absorbed and
released energies manifest themselves as colored
lines in the spectrum of a heated element. Thus,
he saw that this theory could explain the
mysterious patterns of spectral lines and the
stability of atomic structure.
Gilbert Newton Lewis (1875-1946) saw that
the pattern of the periodic table itself reflected
how electrons filled allowed shells around the
nucleus. The innermost shell held up to 2
electrons. Thus, helium with atomic number 2
was neutral with a filled shell of electrons.
Because atoms tended to give or accept electrons
until they filled an outer shell, their reactive
properties depended upon the relative number of
electrons in the outer shell. Helium was full; so,
it was unreactive.
Examine Figure 1 or Figure 8. The elements
of the first row have just a single shell that can
take 2 electrons. The next row can have a shell
outside of that with up to 8 more electrons (10
total in neon). Row 3 can take an additional 8
electrons outside of the 10 (18 total in argon),
etc. Elements such as sodium (Na) with atomic
number 11 have the shells characteristic of neon
but a single electron in the next shell. Sodium
(and all other elements in its family) give up the
lone electron and ionize with a valence of +1
(that is, the ion has one more proton than
electron). Chlorine (atomic number 17) has an
almost full outer shell and accepts an extra
electron to fill it. Thus chlorine tends to become
charged as -1.
The problem of the inequality of atomic
weights and atomic numbers was solved by
James Chadwick (1891-1974), a colleague of
Rutherford, in 1932 with his discovery of the
neutron, an uncharged particle in the nucleus.
Also, because the neutron was about as massive
as the proton, it helped to explain why there
could be variation in atomic weight within an
element. This was the concept of isotopes and
was the last piece necessary to understand the
constituent parts of the atom and how they
related to an understanding of the patterns of the
periodic table.
The last major change to the periodic table
came with a consideration of the so-called rare
earth elements and where they might squeeze in
the overall pattern. Glen Seaborg (1912-1999)
solved that problem by pulling out the lanthanide
and actinide series and placing them as two rows
at the bottom of the table. This has become
known as the long form of the periodic table (see
Figure 1). Also, Seaborg created transuranic
(beyond uranium) elements like plutonium and
showed that they, too, obeyed the periodic law
with 15 elements in each of the rare earth rows.
The rare earths return to the main body of
the table with element number 104. In theory,
any number of elements can be created.
However, most transuranics are so unstable that
they can exist for only seconds or fractions of
seconds before they break apart into smaller
elements.
SEABORN AND SEABORG
The periodic classification of the elements is one
of the most valuable generalizations in science.
-William H. Brock (1992)
In 1998 I stood in line at a meeting in Las
Vegas to meet Glen Seaborg. He was there to
pass out signed copies of a periodic table of the
elements that featured his modification.
I
remember that his hands were gnarled and he
looked ill. I shook his crooked hand as I
accepted a copy of the periodic table and
engaged him briefly in conversation. I knew
enough then to ask about his actinide
modification. His eyes lit up, and he seemed
younger as he talked about the periodic table.
Then, he said pointing to the piece of paper that
he handed me, "These aren't just colored boxes.
They represent work by many people for many
years, and what I did was small compared to
them."
Indeed, many contributed to a discovery and
understanding of the patterns contained in the
periodic table. But patterns in nature often
reflect underlying structure and elucidation of
the smallest part of that structure is difficult and
important. Now it seems that the periodic law is
a consequence of the deeper structure of quarks
and other subatomic particles.
The story of the Periodic Table of the
Elements illustrates very well David Seaborn's
assertion that scientists seek patterns and attempt
to explain them. Indeed, it is the need to explain
the patterns that sets science apart from other
human activities such as art. However, scientists
have to be careful that they do not confuse the
pattern for the explanation. At first, Mendeleev's
table had no explanatory power, but it provided
direction in the search for new elements. It
provided a framework against which the patterns
of spectroscopic analysis and other elemental
properties had to make sense. Thus, the periodic
law guided questions and helped to frame the
answers about atoms, their constituent parts, and
properties of elements.
FIGURE 7. Meeting Glen Seaborg in 1998.
-April 2001
Sources that I used to write the essay:
Brock, William H. 1992. The Norton History of
Chemistry. W. W. Norton & Co. New York.
Cobb, Cathy & Harold Goldwhite. 1995.
Creations of Fire. Plenum Press. New York.
Holt, Jack R. & Patricia A. Nelson. 2001. Paths
of Science, Explorations for Science
Students and Educators.
Kendall/Hunt
Publishing Company. Dubuque, Iowa.
Hudson, John. 1992. The History of Chemistry.
Chapman & Hall. New York.
Jaffe, Bernard. 1931. Crucibles, The Lives and
Achievements of the Great Chemists.
Jarrolds Publishers. London.
Kirchhoff, Gustav & Robert Bunsen. 1860.
Chemical Analysis by Observation of
Spectra. Annalen der Physik der Chemie.
110: 161-189.
Leicester, Henry M. 1956. The Historical
Background of Chemistry. John Wiley &
Sons, Inc. New York.
Moseley, H.G.J. 1913. The High Frequency
Spectra of the Elements. Phil. Mag. (1913):
1024.
Simmons, John. 1996. The Scientific 100, A
Ranking of the Most Influential Scientists,
Past and Present. Citadel Press. New York.
Trefil, James and Robert Hazen. 2000. The
Sciences, An Integrated Approach. 2nd ed.
John Wiley & Sons, Inc. New York.
FIGURE 8. The Periodic Table of the Elements that Glen Seaborg gave to me in 1998.
QUESTIONS TO THINK ABOUT
1. What was the importance of the Chemical Congress at Karlsruhe?
2. How did the formation of HCl show problems with prevailing views? How were
the problems solved?
3. What is a spectroscope? How did its use help in the discovery of atoms?
4. Ultimately, what did the spectroscopic signature mean?
5. What is the periodic law? How did it come to Mendeleev?
6. How did the periodic table help lead to the discovery of new elements?
7. Who, besides Mendeleev, discovered the periodic law?
8. What was Prout’s Hypothesis?
9. What did I mean by the “lazy elements”?
10. How did Moseley change the meaning of the periodic table?
11. How did Lewis increase the explanatory ability of the periodic table?
12. Who was Seaborg and what was his modification to the periodic table?