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THE PERIODIC TABLE
Third edition, December 2005
ORIGIN OF THE PERIODIC TABLE
There are more than 100 elements. If you were to pick a few elements at random
and compare them, you would probably find great differences in their properties.
However, there are certain elements that are similar to one another in many ways.
During the nineteenth century, when most of the elements were being discovered, many
chemists tried to classify the elements according to their similarities.
Antoine Lavoisier, (1734-1794) was one of the first to attempt to organize and
classify the known elements and sadly, I have no information concerning his efforts. His
career was literally cut short. Remember, stupid French extremists beheaded Lavoisier in
1794 during the Reign of Terror because he was a member of the nobility and a tax
collector. Despite his abbreviated life, Lavoiseur’s contributions make him a strong
contender for the title, Father of Chemistry. Chemistry is a separate and distinctive
branch of science largely due to the early efforts of Lavoiseur to categorize, classify,
refine, and re-define chemical concepts and procedures.
Johann Dobereiner (der-buh-RINE-ner) (1780-1849), a German chemist
responsible for the idea triads unified by similar mathematical differences between their
atomic masses and similar chemical properties. Lithium, sodium, and potassium make a
triad of three chemically similar elements and if the difference between their atomic
masses is computed: 23-6.9 = 16.1 and 39.1 - 23 = 16.1, it appears that these three belong
together. However, the system falls apart quickly when more elements are added and his
system, introduced in 1848, lasted for only a few years at best.
The Englishman John Newlands (1837-1898) who organized elements in octaves
were among those who noted that the atomic masses of elements seemed to be related to
chemical properties.
For example: Li-1, Be-2, B -3, C -4, N -5, O -6, F-7 and
Na-8,
Mg-9, Al-10, Si-11, P-12, S-13, Cl-14 and
K-19,
Ca-20, Ga-31, Ge-32, As-33, Se-34, Br-35
(Note that the third row would not have been completed by Newlands because a
few of those elements were yet to be discovered such as gallium in 1875 and germanium
in 1886. Bromine was a recent discovery in 1826, lithium and selenium was discovered
in 1817, boron, magnesium, calcium , strontium, and barium all in 1808 and sodium and
potassium in 1807. Arsenic has been around since AD1250.
Notice that the properties repeat with every eighth element, thereby suggesting
octaves. However, this system falls apart when the transition elements are considered. It
was a clever idea but provided no room for expansion or for unknown elements. Some of
those were not discovered until later, scandium and samarium in 1879 for examples, but
many were already known and apparently ignored by Newlands.
Lothar Meyer (1835-1895) , a German chemist, developed this idea further.
Meyer stated that there was a pattern of repetition in the properties of the known elements
when they were arranged in order of their atomic masses. Meyer is often considered a
co-discoverer of the periodic table. He doesn't receive much attention so I always try to
include something about him here.
The Russian chemist Dimitri Mendeleev (duh-MEE-tree men-duh-LAY-ef)
(1834-1907) was the first to publish the classification of the elements that became the
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basis for the system used today. Incidentally, he was a remarkable fellow in many ways.
Born in Siberia, he was the youngest on seventeen children born to his parents.
Apparently he was recognized quickly as an intelligent little fellow because his parents
made the effort to move to St. Petersburg when he was still a young child so that he could
be cultivated in the more cosmopolitan atmosphere. He attended the University of St.
Petersburg, earned a master's degree in chemistry, and accepted a teaching position at the
university that he held until the 1890's. As a professor of chemistry, Mendeleev wrote a
classic organic textbook, observed a solar eclipse from a hot air balloon, engaged in
radical reform politics as a professor and eventually lost his job with the University of St.
Petersburg. In 1869, he developed a table of the elements that was so carefully planned
and contained such detail that he is credited as the major contributor to a successful
classification of the elements. He must certainly be considered one of the most gifted
scientists of his time if not of the modern scientific period. It has been suggested that his
portraits exhibit a mass of hair and a bushy beard because he would only get a haircut and
a shave once a year in the spring. Students have asked me if Dimitri was a married man.
Yes, Mendeleev married and I know that he had two sons. One of them became a noted
scientist in the area of geology, I believe but the other suffered from mental problems.
He may have suffered from schizophrenia. I have no idea what the fate was of his sixteen
older siblings.
Mendeleev arranged the elements in order of their atomic masses. In an early
version of his work, he placed elements with similar properties in vertical columns. If no
known element had the expected properties to fit a particular place on the table, he left
that place empty. Mendeleev assumed that elements not yet discovered would fill the
empty places. He even predicted some of the properties of these undiscovered elements.
The value of Mendeleev's table was confirmed when elements were discovered that had
the expected properties.
Comparison of Mendeleev's "Ekasilicon" with Germanium
Property
Mendeleev's
Actual
prediction
value
for "ekasilicon"
for germanium
atomic mass
72
72.59
density, g/cm
5.5
5.32
formula of oxide
EsO2
GeO2
density of dioxide, g/cm3
4.7
4.228
formula of chloride
EsCl4
GeCl4
density of tetrachloride, g/cm
1.9
1.844
Mendeleev's table was called the periodic table. It was based on the original
periodic law that stated that the physical and chemical properties of elements were
periodic functions of their atomic masses. This meant that if elements were arranged in
order of increasing atomic masses, those with similar properties would show up at regular
intervals, of periods. These similar elements were called families. For example, the
halogen family contains the chemically elements fluorine chlorine, bromine, and iodine.
Lithium, sodium, potassium, rubidium, and cesium make up another family called the
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alkali metals.
Mendeleev realized that there were a few irregularities in his chart. For example,
iodine has a smaller atomic mass than tellurium. However, iodine is chemically similar to
chlorine and bromine. If placed in the column for those elements, iodine follows
tellurium, as all the elements in groups based on similar properties, disregarding the fact
that occasionally an element seemed to be out of order because its atomic mass was less
than, not greater than, the atomic mass of the element that preceded it. Mendeleev
believed further study would show that his placement of the elements according to their
properties was the correct ordering.
Mendeleev was able to use accurately determined atomic masses to determine
some of the trends on the periodic table due to the work of Stanislao Cannizzaro, an
Italian scientist (1826-1910). He was an interesting fellow. He participated in the
revolutions of 1848 as a major with artillery and fled to France when the insurgents got
their butts whipped by reactionaries. Marx wrote an interesting history of the revolutions
of 1848. Making his way to Paris he met M.E. Chevreul, a well-known chemist. He
stayed with this charming fellow and began contributing to the body of knowledge called
chemistry. In 1860 he presented a paper describing his method of accurately determining
the atomic masses of the elements that represented a great improvement over previous
methods.
Cannizzaro did return to Italy as a professor of Chemistry at the University of
Palermo (1861-71) and at the University of Rome (1871-1910). For his contributions in
chemistry, he was awarded the Copley Medal in 1891 by the Royal Society.
Moseley and the modern periodic law
Henry Moseley (1887-1915) was a brilliant English physicist. He performed
experiments that revealed the atomic numbers of several elements. In 1914, Moseley
suggested that the elements in the periodic table be arranged in order of increasing atomic
number instead of atomic mass. Today, scientists know that electron structure is the
major factor responsible for the chemical properties of an element. Electron structure is
related to atomic number, which does in fact determine the position of an element in the
periodic table.
The modern periodic law, as proposed by Moseley, states:The chemical and
physical properties of the element are periodic functions of their atomic numbers.
Put another way, the law states that when the elements are arranged in order of increasing
atomic number, there is a periodic repetition of their properties.
During World War I, 1914-1918, Moseley was killed in action while still in his
twenties. This was really a stupid waste on the part of Britain but then, WWI was a really
stupid war with no specific gains fought for. If western civilization is ever completely
destroyed, you may date its final decline from the 1913-1918 historical period. Partly
because of this loss, Britain belatedly adopted a policy of using scientists only as
noncombatants in times of war.
Reading the Periodic Table
In modern forms of the periodic table, the elements are arranged in horizontal rows in
order of increasing atomic number. Such a table shows the symbol for each element, its
atomic number, its atomic mass, and usually much additional information.
Each horizontal row in the table is called a period. There are seven periods,
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numbered 1 to 7. All seven periods begin at the left of the table with an active metal and
all but the seventh period end at the right of the table with a noble gas.
Each vertical column of elements is called a family. There are 18 groups,
numbered 1 to 18. The elements in a group have similar physical and chemical properties.
Every element is a member of both a period and a group.
The periodic table and electron structure
It was the study of similarities in physical and chemical properties of elements
that led to the development of the periodic table. Because electron structure determines
chemical properties, you can expect that the table will be related to electron
configurations. In particular, you will find that for every member of a group or family,
the electron arrangement of the valence shell is the same. Characteristic electron
configurations also appear in groups 3 to 12. In these groups, the s level of the outermost
energy level contains one or two electrons, and the d sublevel of the second-fromoutermost energy level is becoming filled.
Periods of Elements
In general, the number of the period in which an element is found is the same as
the number of the energy level of its valence electrons. It is also the same as the number
of occupied energy levels in atoms of the element.
Periods 1,2,and 3 are known as the short periods because they contain
elements with up to only 2, 8. and 8 electrons, respectively, in the first three levels.
Most of the elements in these periods are relatively common in the earth's crust, oceans,
and atmosphere. Atoms of these elements contain only s and p electrons. Period
1consists of hydrogen and helium. In helium (Z=2), two electrons complete the first
energy level. In the elements of Period 2, electrons are being added to the second energy
level. In neon (Z=10), the first two energy levels are filled. In the elements of Period 3,
electrons are being added to the third energy level. The 3s and 3p orbitals are filled in
argon (Z= 18).
Periods 4 and 5 are long periods, each containing 18 elements. Period 4 begins
with potassium (Z=19) and calcium (Z=20), atoms of which contain only s electrons in
the valence shell. The next nine elements, beginning with scandium (Z=30), have
properties unlike any of the elements can not be placed under any of the elements in the
earlier periods because only elements with similar properties go in the columns
established by the elements in those periods. Thus, these 10 elements, scandium through
zinc, become the first members in their own groups, or families, numbered 3 to 12. These
families appear in the d block of the periodic table.
Any element with an atom that has an incomplete d subshell of that gives rise to a
cation or cations with incomplete d subshell is known as at transition element. The nine
elements in Period 4 having atomic numbers 21 to 29 fit that definition. After zinc
(Z=30), the six elements that complete Period 4 are gallium (Z=31) to krypton (Z=36).
Each successive element has one more electron in the p sublevel than the element
preceding it. This pattern repeats that Periods 2 and 3.
Period 5 fills in much the same way as Period 4. Elements with Z=37 and Z=38 are in
the s block, followed by 10 elements in the d block and six elements in the p block.
Period 6 contains 32 elements, and Period 7 contains space for 32 elements, of
which 23 are known. Period 6 contains a series of transition elements that follow
lanthanum (Z=57), called the lanthanoids. Most of these elements fill at sublevel f.
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Period 7contains a series of transition elements that follow actinium (Z=89), called the
actinoids. As with the lanthanoids, most of the actinoids are characterized by the filling
of f sublevel orbitals are in the second-to-the-outermost energy level, the 5f sublevel.
The actinoids go on the f block below the lanthanoids, outside the main body of the table.
You will not find these elements in common household products such as oven cleaners or
air fresheners as they are radioactive and artificially produced.
Groups of Elements
Each group has characteristic properties that are directly related to electron
configuration. In going from top to bottom of any group, each element has one more
occupied energy level than the element above it. Otherwise, their electron structures are
quite similar. As a result, the elements in a vertical group have similar chemical
properties. The alkali metals are all the elements in Group 1 except hydrogen. The
alkaline earth metals are the elements in Group 2. Both of these groups of metals are so
active that they occur in nature only in compounds. One reason they are so active is that
their atoms readily lose electrons to form positive ions.
The nitrogen family of elements is found in Group 15 and the oxygen family,
(also called the chalcogen family) in Group 16. Both are groups of nonmetals. Nitrogen is
relatively inactive, while oxygen is highly active, able to form compounds with nearly
every other element.
The nonmetals in Group 17 are called the halogens. Like the active metals, these
nonmetals are so active that they do not occur free in nature. Fluorine is the most active
of all nonmetals.
The members of Group 18 are called the noble gases. Formerly, these elements
were called "inert gases" because they were both known to combine with other elements.
But because compounds containing xenon and krypton have been produced in the
laboratory, the name "inert" is no longer is used.
Periodicity in Properties
A study of the periodic table shows certain regularities in the properties of the
elements. Recall that the members of a vertical group of elements have similar properties.
This similarity occurs because the members of a group have the same number of valence
electrons. In the sections that follow, the periodicity of several properties of the elements
will be discussed. These properties are related to the presence of opposite charges within
the atom and the total number of electrons in the atom. This will be incredibly
interesting.
Some properties of the elements are directly related to the attraction of the
positive nucleus for the negative electrons. This attraction, which depends on both the
quantity of charge and the distance separating charges, is called coulombic attraction.
Two properties of elements that are closely related to the coulombic attraction between
nucleus and electrons are ionization energy and electronegativity. Two other properties of
the elements are more closely related to the number of electrons present in the atom.
These properties are atomic radius and ionic radius. The attraction of opposite charges
also plays a role in determining these properties. Each of these properties is discussed in
later sections of this chapter.
Ionization Energy and Periodicity
Energy and ionization. The ionization energy of an atom is the energy required to
remove the most loosely held electron from the outer energy level of that atom in the
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gas phase. Another name for ionization energy is ionization potential. The removal of a
single electron from an atom can be represented by the equation
M + energy ---> M + + e-.
In this equation, M(g) represents the gas phase of any element and the "energy" is
the ionization energy. M+(g) represents the gas phase of the positive ion that is formed
after ionization energy is applied. The electron that is separated from the atom is
represented by the final symbol,e-.
M + energy ---> M+ + eM+ + energy ---> M+2 + eM+2 + energy ---> M+3 + e-
(first ionization energy)
(second ionization energy)
(third ionization energy)
The energy in the lovely equation is the amount of energy required to remove the
highest energy electron from the valence shell. Because this is the removal of the first
electron from a neutral atom, the energy required is the first ionization energy. The
energy required to remove a second electron is the second ionization energy. The energy
required to remove a third electron is the third ionization energy. Probably all of you can
reasonably guess that the energy required to remove a fourth electron must be the fourth
ionization energy. Pretty cool, huh. Each successive ionization requires more energy
because each successive electron separates from a particle that has increasingly
greater net positive charge.
Periodic trends in ionization energy.
Ionization energy is a periodic property of the elements. To see the trend in ionization
energy, look at the part of the handy enclosed table of first ionization energies for
representative elements in Period 2 , 3, and 4.
In kJ/mole
Li
520
Na
490
K
420
Be
900
Mg
730
Ca
590
B
800
Al
580
Ga
580
C
1090
Si
780
Ge
780
N
1400
P
1060
As
960
O
1310
S
1000
Se
950
F
1680
Cl
1250
Br
1140
For all three periods, the general trend is toward an increase in ionization energy
along with an increase in atomic number, with some minor exceptions, the atomic
number increases. Within a group of elements, atomic number also increases, suggesting
the ionization energy would increase. However, as you move down the periodic table
toward increasing atomic numbers within a group, the ionization energy actually
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decreases.
One reason for this decrease is that each successive member of a group has one
additional energy level of electrons. Thus, there is a greater distance between the positive
nucleus and the negative electrons, resulting in a decrease in the force of attraction.
Another reason for the smaller force of attraction is that in larger atoms successive layers
or principle energy levels "shield" the outer electrons from the attractive force of the
nucleus . This shielding effect is caused by the repulsion between kernel electrons and
valence electrons.
Electronegativity and Periodicity
Another property of elements that is related to the forces of attraction between the
positive nucleus and negative electrons in atoms is electronegativity. Electronegativity is
a measure of the ability of an atom of an element that is chemically combined with
another element to attract electrons to itself. The forces that hold atoms of chemically
combined elements together are called chemical bonds. Put another way,
electronegativity is a measure of the force of attraction that exists between an atom
and a shared pair of electrons in a covalent bond. Covalent bonds are formed when
pairs of electrons are shared between atoms.
The principles that determine the electronegativity of an element are similar to
those that apply to ionization energy. The scale of electronegativity proposed by Linus
Pauling is used most often. Values range from a low of 0.7 for several metals in Group 1
to a high of 4.0 or 4.1 for fluorine in Group 17.
The fact that fluorine has the highest electronegativity value suggests why
fluorine will never form a positive ion. There is simply no other element with a greater
electronegativity capable of removing an electron. In other words, fluorine is the biggest
dude on the block and it will only take electrons and never give one up.
Remember from previous information that the electronegativity differences
between atoms predict the type of bonding. Imagine elements A and B with values of 3.0
and 2.8. The difference between them is 0.2. That is a difference less than 0.4, therefore
the bond is probably nonpolar covalent. Imagine elements A and C with values of 3.0 and
1.9. The difference is now 1.1, indicating a polar covalent. The electrons are still being
shared but not equally. Element A is negative and C is positive. Now imagine elements
A and D with values of 3.0 and 0.6. The difference is 2.4 and the bond is ionic with no
pretense of sharing. A becomes a negative ion and D is a positive ion much like chloride,
Cl- and sodium, Na+.
Electronegativity is an important concept used to describe the characteristics of
types of bonds; covalent and ionic. The relative position of an element on the periodic
table suggests something about its relative electronegativity and the type of bonding
pattern possible.
Position of Electrons
Another way to think of the properties of ionization energy and electronegativity
is to picture a positive nucleus attracting electrons located in one of two possible
positions. For ionization energy, the attracted electron is in the outer shell of the atom it
is being removed from. In the case of electronegativity, the attracted electron is one of a
pair of electrons shared by, and located between, two atoms.
Ionization energy and electronegativity are similar because they both involve an
amount of charge and the distance that separates charged particles. Ionization energy is a
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quantity that can be measured in the laboratory. By contrast, electronegativity is not a
measurement of energy. Values for electronegativity cannot be measured directly.
Electronegativity values are determined mathematically, using equations developed by
Pauling from observations of bond energies.
Ionic Radius
In many chemical changes, atoms are converted into ions by taking on or losing
electrons. Atoms of metals are more likely to give off electrons to become positively
charged ions. Atoms of nonmetals tend to take on electrons to become negatively
charged ions. For a monatomic ion, which is formed from a single atom, the ionic radius
is the effective distance from the nucleus of the ion to its outer shell of electrons. For
many monatomic ions, this outer shell consists of eight electrons, the filled sublevels of
the highest principal energy level. "Effective" in the definition is a reminder that ions,
like atoms, have electrons in the form of a diffuse cloud of negative charge. Therefore an
ion has no identifiable edge.
Comparing atomic and ionic radii.
Note that the atomic radii are defined as the radius from the nucleus to the
outermost electrons in a neutral atom. This is before any electrons are lost or gained.
Ionic radii are defined as the distance from the nucleus to the outermost electrons after
electrons have been lost (to form a cation, or positive ion) or gained (to form an anion or
negative ion.)
When an atom takes on one or more electrons to become a negative ion
(anion), its radius increases. Most atoms of nonmetals take on enough electrons so that
there are eight electrons in the valence shell. These added electrons increase the forces of
repulsion between electrons, pushing them farther apart. As a result, the radius of the ion
is greater than the radius of its parent atom.
When one or more electrons are removed from an atom to form a positive
ion (cation), the atom's radius decreases. A typical metal forms a positive ion when it
loses some or all of its valence electrons. Usually, the entire outer energy level is lost,
causing a large decrease in the radius of the atom as it is converted to an ion. Within a
group of elements, as the atomic number increases, ionic radius increases. With each
successive member of the group, a new occupied energy level is found in the kernel,
placing the valence shell farther from the nucleus. This causes an increase in the radius
of the ion. Trends in ionic radius within a period do not follow a clear pattern. It is not
easy to compare the ions of metals with the ions of nonmetals within a period because
both nuclear charge and electron configuration change as the ions form from the atoms.
Isoelectronic Species
Those kinds of atoms or ions that have the same electron configuration are
referred to as isoelectronic species. Many of the ions that form when atoms take on or
give off electrons have the same electron configuration as the nearest noble gas. For
example, an atom of nitrogen gains three electrons to form the nitride ion, N-3. The
electron configuration of the nitride ion is the same as that of neon (1s22s22p6). An atom
of magnesium (1s22s22p63s2) gives off two electrons to form the magnesium ion
Mg+2(1s22s22p6). The electron configuration for Mg2+, like that for N-3 is also the same
as that of neon, Ne. These three particles , or species, are isoelectronic. A longer series of
isoelectronic species includes N-3, O-2, F-,Ne ,Na+1, Mg+2, and Al+3. Each particle in this
series contains 10 electrons. Within the series as the atomic number increases, the
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number of protons increases, so that the positive charge increases. The larger positive
charge exerts a greater attractive force on the electron cloud, causing it to become smaller
as the electrons are pulled closer to the nucleus. Thus, the radius of the particle becomes
smaller. In general, within a series of isoelectronic species, as the atomic number
increases, the radius of the particle decreases.
Another example of an isoelectric series is one with arsenic, selenium, bromine,
krypton, rubidium, and strontium. If you are bored either because you can’t get a date
this Friday night or if your date is boring, try working the notations out remembering to
make the first three anions and the last two cations.
Metals, nonmetals and semimetals
Most of the elements are metals. In the periodic table, you can see that most of the
elements in groups 1 to13, some of the elements in groups 14-16 and all of the
lanthanoids and actinoids are metals. Most properties of metals can be explained by the
presence of relatively loosely held valence electrons that causes them to form positive
ions.
There are relatively few nonmetals among the elements. They make up parts of
Groups 13-16 and all of groups 17-18 in the periodic table and are found in the upper
right hand corner of the handy-dandy periodic table. Most properties of nonmetals can be
explained by the fact that they have a nearly filled p sublevel in the valence shell. Most
nonmetals have one, two, or three vacancies in the p sublevel.
In the periodic table there is a step like line drawn to the left of boron and
continuing downward to the left of silicon, arsenic, tellurium, and astatine. Five of the
elements to the right of this line and two of the elements to the left are semimetals. The
properties of semimetals are neither distinctly metallic nor distinctly nonmetallic.
Germanium, for example is brittle and a poor conductor like a nonmetal, but it forms
positive ions like a metal. Another name for these elements is metalloids. Semimetals
have certain unusual properties that make them useful as semi conductors in transistors.
Good examples of these elements are boron, germanium, and silicon.
Trends in metallic and nonmetallic character. Like other properties of the
elements, metallic and nonmetallic characters exhibit trends within the periods and
groups of elements. For any period, the elements near the left are more metallic, while
those near the right are more nonmetallic. For any group, the elements near the top are
more nonmetallic, while those near the bottom are more metallic. Elements at the lower
left corner of the periodic table are the most metallic. Except for the noble gases,
elements at the upper right corner of the table are the most nonmetallic.
WHAT'S IN A NAME?
Unlike explorers of old, who could name newly found lands whatever they
wanted, the researchers who discover new elements have to submit their choices to a
committee before they can leave their mark on the periodic table, the map of the chemical
world. In September 1997 the International Union of Pure and Applied Chemistry
resolved the debate on six laboratory-made elements, assigning the names rutherfordium
(Rf), dubnium (Db), seaborgium (Sg), bohrium (Bh), hassium (Hs), and meitmerium (Mt)
to elements 104 through 109. The numbers refer to the total protons in each element's
nucleus.
The first sample of rutherfordium was produced about 30 years ago, but the
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naming process was drawn out by competing claims from laboratories in the United
States, Germany, and Russia. A joint commission of physicists and chemists sought a
compromise that would satisfy all. The Russians gained recognition for work done at a
laboratory in Dubna-to the dismay of the Americans, who dispute some of that lab's
claims. The Americans wanted to honor 85-year-old Glenn Seaborg of the University of
California at Berkeley, as well as Ernest Rutherford, the British physicist who discovered
the atomic nucleus. The Germans wanted to enshrine Lise Meitner, co discoverer of
atomic fission, as well as one of their labs in the state of Hesse. And both the Germans
and Russians pushed for Niels Bohr. "I suspect the feeling between the United States and
Russia is a relic of the cold war," says Jeffrey Leigh of the chemists' union. "In the end,
this is a compromise that doesn't satisfy anybody completely." At least Seaborg, the
Nobel laureate discoverer of ten elements, including plutonium, has reason to be happy.
"This is the greatest honor that I've ever received," he says. "One hundred years, a
thousand years-there it is, still in the periodical table." Not only is Seaborg the first
living scientist to have an element named after him, he's also the only person who could
receive mail addressed only in elements: Seaborguim, Lawrencium (for the Lawrence
Berkeley Laboratory where he still works), Berkelium, Californium, Americium.
-Jeffrey Winters
taken from the January 1998 issue of
Discover Magazine
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