|Metals||Semimetals||Nonmetals||Z, atomic #Symbol|
|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.|
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.
|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
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.
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.
|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.|
|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.|
|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.|
|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:
even though those 4 phosphorous atoms in the P4O6 are
likely the same quartet that started in the phosphorous solid!
|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.
|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!!||
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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