|Metals||Semimetals||Nonmetals||Z, atomic #Symbol|
|Ionic molecules share many properties. A high melting point is only one. But all of
these shared properties are a result of the virtual capture of electrons from the
"electropositive" elements to the "electronegative" ones. Eventually, we'll see a
quantitative measure of what electropositive and electronegative mean, but at this point
it suffices to say the electropositive atom relinquishes the bonding electrons to the
electronegative one. The attraction between the two atoms is then predominately
|Melting requires breaking that electrostatic bond between some of the ions but then
re-establishing it with others as the ions migrate in the melt. Boiling, of course,
breaks all of those bonds, and boiling points for ionic molecules are therefore
|Of course, in the melt, the ionic molecules can do something that they couldn't do
as a solid. Since the ions are free to roam, these melts conduct electricity. In
fact, this property is of critical importance in the separation of some metals from their
ionic ores. The electrical current applied to the melt can cause the metal ions to migrate
to the negative electrode and be electrically neutralized (for these positive "cations" to
become neutral metals, the proper word is not neutralized but "reduced"). This is a topic
we'll take up under electrochemistry.
|The matrix of +-+-+- charges in these solids makes them vulnerable to cleavage, for a
blow which dislocates one part of an ionic solid by only one atom forces the attractive +-
pairs along the fault all to turn to ++ and -- repulsions. So perfect crystals of ionic
materials can be snapped if struck along the correct planes.
|Non-ionic bonding (called covalent) between molecules means that there is no radical
separation of charges as occurs with ionic bonding. This gives molecular solids quite
different properties. Although the covalent bonding between atoms in the molecules can be
quite strong, there's no electrostatic attraction between neighboring covalent molecules;
so it is easier to disrupt such solids by melting, and they have much lower melting points.
|Since molecular solids involve no ions, their melts do not conduct electricity at
ordinary voltages. In general, then, they make good electrical insulators.
There is an
important exception, but he only serves to highlight the rule. The exception is the
element CARBON. Since every C atom in the material has an equal affinity for its electrons,
all the C-C bonds must be covalent. Yet carbon in its solid form, graphite, bonds to
each of three neighbors forming a chicken-wire plane. Diamond, another form (allotrope) of
the element, goes graphite one dimension better and permits each atom to bond to FOUR
neighbors in an infinite, 3-d, tetrahedral lattice. (Now you know why diamond is so hard;
it's holding on in all directions at once.)
|So graphite and diamond aren't merely molecular solids, they are (infinite) covalent
solids and share some of the tough properties of the (infinite) ionic ones. Although
diamond doesn't conduct electricity, graphite does (sort of) by permitting the electrons
to run down the planes between the graphite sheets (see the left picture above). So no
ions are involved in graphite's conduction. Graphite is actually used to make
insulators as well since when the planes are oriented at random in graphite flakes, the
conduction drops off.
|So although it's C,
the other covalent molecules are not so lucky. For example, running down the Periodic Table
starting with nitrogen, the molecules formed are N2, P4, As4,
but no infinities. Likewise, it's O2, S8, Se8, but no
infinities. So such molecules are only vaguely interested in their neighbors in
ways we'll study in Chapter 11.
|Even more instructive is to run across the Periodic Table starting with Li and
looking at the "hydrides" (bonds formed with hydrogen) which show up as|
OK, H0Ne is a cheat; I really mean that there is no neon hydride because, of course, Ne belongs to the noble gases and sneers at virtually all bonding. And it isn't accepted practice to include the "1" in molecules like H1F, but I wanted to emphasize the periodicity of the hydrides: 1, 2, 3, 4, 3, 2, 1, 0. This was certainly not lost on Mendele'ev! It means that something fundamentally simple is going on.
|And the Periodic Table is periodic in ions as well (mostly) as seen in the common ions
of that same second row:|
So the charges these guys will tolerate are in an intelligible sequence too.
|Ionic Names||Even though it's really
we always refer to table salt as NaCl. The Na lies directly beneath Li on the Periodic
Table, so we're pretty confident that its preferred ion would be Na +. Likewise,
Cl lying below F should prefer Cl -, and it does. All positive ions are called
cations (they get made at cathodes in batteries), and all negative ions are
called anions (yes, there are anodes too). But the naming isn't as simple as
calling NaCl, "sodium chlorine." Instead, atomic anions get a new suffix to distinguish them
from their atoms; so while Cl is chlorine, Cl - is chloride. Logically, of
course, something should be done to distinguish cations from their atoms, but historically,
chemists have never bothered! So NaCl becomes sodium chloride.
|This all works as long as atoms can only ionize one way. For example, calcium can only
ionize as Ca 2+, and nitrogen only as N 3-, so their ionic compound
would be calcium nitride. But it surely wouldn't be CaN!
(I like that; it CaN't be CaN.) The charges don't balance out to make
a neutral molecule, and calcium nitride is, in fact, neutral. It shouldn't take you long to
realize that we have to take the ions in different proportions corresponding to their
different charges! So the simplest form we could propose would be Ca3N2
which is now electrically neutral: six pluses and half a dozen minuses.
|As long as the ions can support only one magnitude of charge, we can name ionic things in
the manner above. But when we hit the transition metals, their ions aren't unique. I love to
trot out my favor phrase and tell you "we'll learn why later," but it's probably not satisfying.
The reason for different stable ions of the same atom lie in atomic structure and we have to know
that first; so for now (shudder) we're just going to have to memorize the important ones.
|Worse still, there are two naming conventions in current use. The most rational is the
most modern one where the cation is given the atom's name with a (literal) parenthetical comment
on its charge. Thus, we have the iron(II) and iron(III) ions which are, of course, Fe2+
and Fe3+, respectively. So if we've memorized which charges atoms prefer, we're home
free in this notation. So your car battery uses both PbSO4 or lead(II)
sulfate and PbO2 or lead(IV) oxide.
(We'll get to the compound anions in a minute; so trust me on sulfate
having two minus charges.)
|The Dark Side of the cation names are the Ancient Ones. Under that
(what passed for a) system, the ions took the Latin name of the element
(now we have to know that Pb stands for the Latin name for lead, plumbus,
from which we take our name for plumbing at which the Romans were rather good!)
but they gain suffixes denoting the charge. Sort of. The good news is that there are only
two common suffixes, -ous and -ic. The bad news is that those don't stand for any
particular charge! Instead, -ous is given to the lowest common ion charge
while -ic is given to the higher. That makes
PbSO4 plumbous sulfate and PbO2 plumbic sulfate
for compounds stemming from Pb2+ and Pb4+, respectively.
|But while FeCl2 is certainly ferrous chloride, iron's higher charged
ion is Fe+3, so FeCl3 becomes ferric chloride. So both
lead(IV) and iron(III) salts get the -ic suffix because they're the highest those ions want
|One word of warning about mercury; it prefers mercury(I) and mercury(II) ions, but while the
latter are Hg2+, the former are weird. Mercury(I) ions run around as a pair!
Don't ask why or you'll embarrass me. So they are always written as Hg22+
rather than Hg+. Hence the chlorides, for example, are Hg2Cl2 and
HgCl2 for the mercurous and mercuric chlorides, respectively. The only such exception
we'll see in this course.
|Now we're ready for the table, no? Here below are common cations that come in more than one
|Fortunately, many metal atoms have cations of only one kind of charge. The best examples
are the alkali metals; all of them M+ ions; that goes for H+ too.
Ditto the alkaline earths are all
M2+ ions. And although you'll need to memorize the rest of the common
(non-schizophrenic) cations, at least you'll have the comfort of knowing that there'll be no
-ous or -ic or even parenthetical Roman numerals to have to deal with. Since we all
know these ions don't come in any other form, we call SrO strontium oxide and are
done with it.
|So here are the rest of the cations that are comfortable with who they are and don't need a second (or third!) personality:
Ag+, Al3+, Au3+ Bi3+,
Ni2+. There's some rhyme or reason to Zn2+ and Cd2+ since
they are in the same column in the Periodic Table, but periodicity is of limited help here.
|There's one other non-schizophrenic cation of importance, but it isn't atomic!
Oddly, while there are enormously large numbers of common compound anions, the only
really important compound cation is ammonium, NH4+. And the
reason it's important is that all its salts are water soluble. Handy to ensure they
go into solution and can be diluted or concentrated at will.
|Ignoring the Noble Gases, the last two columns in the Periodic Table give rise to pretty
simple atomic anions. All the elements below fluorine give rise to F - like ions
while those below oxygen all look like O2-.
(Check that: Po is a metalloid and won't bear an anion form.) Before we
run into other metalloids we can have confidence that both N3- and P3-
will bear their anticipated triply negative anionic charge. The only word of caution comes if
you think that these all show up in water; O2-, for example, is so ravenous that
it will steal a hydrogen ion from water itself by|
to make the hydroxide ion instead. So O2- doesn't survive in water.
|What can I say? There are a ton of interesting compound anions. The only things more
numerous that you're going to have to know (eventually) are the organic compounds (for reasons
we'll see in a few weeks). So let's trot them out:
|Did you catch the trend? Most of the names are -ite or -ate with the latter
having more oxygens? With the exception of fluorine, all the halogens show
roughly the same series that
is exemplified by chlorine above. But two suffixes are insufficient; so prefixes get added.
Great! Is there any logic to it at all?
|But you gotta know Greek to catch it.
|So, for example, a hypodermic needle goes under the dermus (your skin). So
hypo means below and hypoiodite IO- has one fewer oxygen
than iodite ion, IO2-. But a hyperactive child is one with
higher than normal activity; so perbromate BrO4- has one
more oxygen than bromate ion, BrO3-. Clever?
|Although metals react well with non-metals, non-metals can bond to one another as well.
The modern way to name such compounds is to prefix each non-metal atomic name with the
Greek number corresponding to its count in the molecule. (Yes, we had to know Latin
numerals for the metals, and now Greek number names for the non-metals.) So, for example,
dinitrogen pentoxide is N2O5 even though the Greek for "5" is
penta. You might think it should then be a "pentaoxide," but that puts two vowels
together and is harder to pronounce; so it's "pentoxide" instead. The other funny rule is
that the Greek "mono" (meaning one) gets left off the first atom if there's only one of him;
so CO isn't monocarbon monoxide but just carbon monoxide.
|So the table of Greek numbers can't be far behind, right?
|So carbon disulfide is CS2, and BrF3 is bromine trifluoride, and dichlorine heptoxide is Cl2O7, and so forth ad nauseum.|