Metals, Molecules, and Ions

Chm 1311 Lecture for 8 June 2000 cont'd.

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1H 2He
3Li 4Be 5B 6C 7N 8O 9F 10Ne
11Na 12Mg 13Al 14Si 15P 16S 17Cl 18Ar
19K 20Ca 21Sc 22Ti 23V 24Cr 25Mn 26Fe 27Co 28Ni 29Cu 30Zn 31Ga 32Ge 33As 34Se 35Br 36Kr
37Rb 38Sr 39Y 40Zr 41Nb 42Mo 43Tc 44Ru 45Rh 46Pd 47Ag 48Cd 49In 50Sn 51Sb 52Te 53I 54Xe
55Cs 56Ba 57-
72Hf 73Ta 74W 75Re 76Os 77Ir 78Pt 79Au 80Hg 91Tl 82Pb 83Bi 84Po 85At 86Rn
87Fr 88Ra 89-
104Rf 105Db 106Sg 107Bh 108Hs 109Mt 110Uun 111Uuu 112Uub 113Uut
Lanthanides 57La 58Ce 59Pr 60Nd 61Pm 62Sm 63Eu 64Gd 65Tb 66Dy 67Ho 68Er 69Tm 70Yb 71Lu  
Actinides 89Ac 90Th 91Pa 92U 93Np 94Pu 95Am 96Cm 97Bk 98Cf 99Es 100Fm 101Md 102No 103Lr  

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.
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.

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.


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 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:

4 P(s) + 3 O2(g)  arrow right P4O6

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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Last modified 7 June 2000. Chris Parr