Contiguous
Conduction |
One of the handier features of the
Periodic Table
is the large
scale grouping of elements that are metals vs.
those that are not. Even the schizoid
semimetals fall in a cluster, albeit neither
rowwise nor columnar. This hints at underlying regularity in atomic structure
which we'll study soon. But at the moment, let's explore just what properties
these are that collect themselves so conveniently in a Periodic Table. |
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METALS |
Clearly the great majority of atoms are metals. Someone once remarked that God
must love the beetle above all other life since He made so many versions of it.
The Periodic Table leads us to believe that God's into heavy metal as well.
Metals share a number of properties we all recognize. They're usually hard,
lustrous (shiny), and
conduct electricity well.
Less well know are properties of ductility and malleability.
A metal is ductile if it can be drawn into a wire; the word itself
(think of air ducts) comes from the Latin meaning "to lead"
(as opposed to "follow"). So you can lead a metal to a wire
dye...and successfully stretch it out into long thin conductors. Clearly,
ductility has to do with the way the metal solid can be deformed without
tearing. A different deformation is involved in malleability; a metal is malleable
if it can be pounded really flat. In class, we mentioned that Rutherford's
alpha rays (2He2+ nuclei) were shot at an ultrathin
gold foil, less than 1000Å thick (100 nm)
because gold, 79Au, is so malleable.
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Metallic
Conductivity |
Conductivity and luster are, in fact, related. Metals conduct because their
electrons are easily shoved around throughout the solid by electric fields.
Metals are shiny because they reflect light well. Light is known to include
oscillating electric fields to which a metal's electrons dance uninhibitedly,
like Snoopy on the top of his dog house. But the very easy induction of that dance
means that the metals reradiate the light just as easily. Hence a shiny mirror
finish.
And the oddball of the lot is mercury, 80Hg, "HydrarGentum" in Latin
meaning "fluid silver", that is liquid at "standard conditions" (0°C &
1 atmosphere pressure). If standard conditions were instead lukewarm, both Cesium,
Cs, and Gallium, Ga, would also be liquids; their melting points are that close
to room temperature. But most metals have reasonably high melting points.
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In contrast, those elements in PINK on the Periodic
Table are nonmetals. Notice that hydrogen, 1H, seems to be
out-of-place, bedding down with the alkali metals. Well, if we really wanted to
press the point, we could argue that hydrogen can be made into a metal
at abysmally high pressures...like those near the center of the planet Jupiter!
It's a spinning metallic hydrogen core thought to be the source of Jupiter's
enormous magnetic field. But around here, hydrogen is quite, quite nonmetallic.
(That's why on the periodic table handout I gave you, I put H closer to the
middle of the table instead.)
One would expect nonmetallic properties to differ significantly from metallic
ones. Just so: no nonmetal is ductile or malleable in its solid state, but then
the majority of the non-metals aren't solids at standard conditions. They're
gases. Of the minority exceptions, only bromine,
35Br, is a liquid at standard conditions...and a horribly
corrosive one at that.
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And, of course, the nonmetals don't conduct electricity...except carbon
in its graphitic form. Remember we noted that graphite, the most stable form
of the element, was composed to planes of carbon atoms interconnected just like
chicken wire? Electrons can flow fairly easily in the spaces between those
planes. But not easily enough to render carbon lustrous; so one would never
mistake "pencil lead," 6C, for automobile battery lead,
82Pb. |
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(Semimetals) |
It's not hard to imagine that the frontier between the metals and the
nonmetals is occupied by
weird atoms
that are neither fish nor fowl. Their
most interesting property is that with small changes in electrical fields,
they could be persuaded to conduct or to insulate (not conduct). This on/off
character is perfect for local Dallas industry which fixates on
semiconductors for computer and communications applications.
The essence of computer memory is a reliable ON ("1") or OFF ("0") state,
and silicon, 14Si, doped with germanium, 32Ge, or
arsenic, 33As, fit the bill, especially since the relatively
expensive dopants are needed in such small quantity while the element of
the chip structure, silicon, is the basis of the most abundant mineral on
earth, sand.
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Compounds |
Although the elements of the Periodic Table are shown as atoms,
many of them don't appear atomic in Nature. For example, with the exception
of the noble gases (2He and down), the gaseous elements are diatomic,
H2, N2, O2, F2, etc. What may be
even more surprising is that when, on lowering their temperature, you get these
elements to condense into liquids or solids, the basic units of matter are
still the diatomic molecules; they don't give up their partners just
because of a change of state. That's loyalty.
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Even some of the normally solid nonmetals have essentially molecular solids;
sulfur is really collections of S8
while phosphorous exists as
P4. Now while chemists routinely use the diatomic forms of the gaseous
elements in writing chemical equations in which they're involved, traditionally,
we don't do that with the other molecular elements. So you'll see:
4 P(s) + 3 O2(g)
P4O6
even though those 4 phosphorous atoms in the P4O6 are
likely the same quartet that started in the phosphorous solid!
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But when elements combine with other elements into compounds,
the combinations and their properties can be stunningly numerous. Still, it is
possible to divide them all into two or three major categories such as
molecular solids and ionic solids. (A third category is a metal
bonding, but we'll investigate that much later.)
Just as O2 exists as diatomic molecules even when it freezes, so too do
molecules like CO (carbon monoxide). But this contrasts with molecules like LiF
which bond in a radically different way. We'll discover that metals are pretty
loose with their electrons, giving them up without much of a fight. On the other
extreme, the halogens will not only cling jealously to their electrons, they'll
try to steal them from other atoms! This is what makes halogens so corrosive.
So when Li encounters F, it puts up no resistance to losing an electron to
become a Li+ ion. The electron gets firmly planted on the fluorine
atom which becomes a F- ion. The Li+F- pair are
now held tightly by the natural electrostatic attraction of unlike charges.
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This difference between CO and LiF becomes critically important in their
solids since while the carbon and oxygen are firmly bonded to one another (in
what happens to be the strongest diatomic bond), they are only weakly attracted
to neighboring CO molecules. Not so in solid LiF. The Li+ finds itself
surrounded by 6 nearest neighbor F- ions, all equally enticing
from an electrical attraction point of view. The same fickle attitude is shared
by each of the F- ions! So there ceases to be molecules of
LiF in the solid but rather one gigantic molecule, every ion linked to
the next in all directions. Not surprisingly then, LiF (25.94 amu) is a high
melting solid while CO (28.01 amu) is a gas even though CO has the higher
molecular weight!! |
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Just to set the record straight, the Periodic Table shown
in these lectures isn't the one Mendele'ev proposed. His original one
looks like this.
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