With hydrogen covered, it’s now time to start looking at the main group elements, one group at
a time. Let’s first mention the fact that there are actually two ways to number these groups.
The older method uses Roman numerals and skips over the transition metals, which gives us groups
IA, IIA, then jumping over here for IIIA, all the way up to VIIIA for the noble gases.
However, the newer method simply uses regular numbers and includes the transition metals,
so we have groups 1 through 18 labeled like this, from left to right. We will be using
this newer method in these tutorials, but it is worth at least mentioning the other method,
so as to save yourself any confusion should you run across this in your studies.
With this straightened out, let’s take a look at Group 1, also known as the alkali metals.
Since hydrogen is not a metal and we covered it in the previous tutorial,
that leaves us with lithium, sodium, potassium, rubidium, cesium, and francium.
We should of course be quite familiar with the patterns inherent in the periodic table already,
so we know that what these elements have in common is that their electron configurations
end in ns1, meaning that their outermost shells are populated by one lone
s electron. Because they all have only one valence electron they share similar properties,
namely that they have the lowest Z effective in their respective periods. They have the largest
atomic radii in their respective periods. They are the most easily ionized elements in their periods,
readily forming M+ ions, meaning that they have very low ionization energies,
and this value gets lower moving down the group, as the atomic radius gets bigger due to adding
additional shells, and the valence electron gets farther and farther from the nucleus.
Lithium is the only element in this group that is capable of sometimes forming covalent bonds,
and this is due to its small size, which results in a strong hold of its electrons.
The rest tend to only make ionic bonds and metallic bonds.
So where can we find these elements, what are their natural sources? Lithium is derived
from ores like spodumene and petalite, which have the formulas given. Commercial deposits
are rare, and found mainly in Bolivia. Sodium is obtained from sodium chloride in seawater.
The seawater evaporates and produces halite deposits. Potassium is obtained in a similar
manner, from minerals in seawater that has evaporated. These minerals include sylvite, which
is potassium chloride, as well as carnallite and kainite, which are hydrates with these structures.
Rubidium is found in the ore lepidolite, with this complicated structure,
where a small amount of potassium is replaced by rubidium. This is found primarily in Brazil,
Russia, the United States, Canada, and Madagascar. Cesium is found in the ore
pollucite, which is found primarily in the Manitoba province of Canada.
Francium is exceptionally rare, formed as the result of radioactive decay of actinium,
which occurs in uranium minerals. All of Earth’s crust probably contains less than 20 to 30 grams
of francium at any given time. Francium can also be made artificially in particle accelerators.
Now a bit on the properties of these elements. Their melting points decrease
going down the group, though they are quite low in general as far as metals are concerned,
given the low lattice binding energies that result from having only one valence electron.
The same trend can be applied to their heats of sublimation and hydration energies.
Lithium, sodium, potassium, and rubidium are all soft and silver, while cesium has a gold sheen.
In terms of preparation, lithium and sodium can be prepared via electrolysis of salts. Molten sodium
chloride will undergo reduction to form liquid sodium and chloride ions. Potassium, rubidium,
and cesium can be prepared by treating their respective metal chlorides with sodium gas. These
are then purified by distillation. Additionally, these metal oxides can undergo hydrolysis to yield
metal cations, hydroxide ions, and in some cases hydrogen peroxide and oxygen gas.
Due to their high reactivities, these metals tend not to exist in pure elemental form in nature,
which is what makes their preparation necessary. Now moving on to reactivity. At modest
temperatures, lithium reacts uniquely with nitrogen gas to form lithium nitride, a ruby
red solid. All the metals react with water to form metal cations, hydroxide ions, and hydrogen gas.
Lithium does so slowly at room temperature. Sodium reacts vigorously. Potassium,
even more vigorously, while rubidium and cesium react explosively. Again,
this is due to the decreasing ionization energy moving down the group. These metals also react
with alcohols to produce metal alkoxides. A variety of other salts are also possible.
Lithium salts have anomalous properties due to the very small size of the lithium cation.
Lithium hydroxide decomposes to lithium oxide for better ion packing, while other alkali metal
hydroxides sublime. Lithium forms salts with halides that are soluble in organic solvents.
It can form salts with the conjugate bases of Bronsted-Lowry acids that are water soluble
to varying degrees. If derived from a strong acid, the salt is very soluble,
due to the mismatch of ion size. If derived from a weak acid, the salt is less soluble because of the
strong bonding in the crystal lattice, which outweighs the energy released in solvation.
Lithium hydride is covalent, as are alkyllithium reagents, which we learned about in organic
chemistry. Methyl lithium is actually a discrete tetramer when in the solid form.
Let’s now examine some applications. Lithium has tremendous utility in batteries. It is very light,
and in fact at 0.5 grams per cubic centimeter has the lowest density of any metal, but it is also
easily oxidized, so it is an ideal substance. Lithium is present in certain disposable
batteries, as well as rechargeable lithium ion batteries, some of which utilize lithium ions
that intercalate, or penetrate, a lattice of cobalt oxide and graphite. Lithium stearate is
used in lithium grease for lubrication in the automotive industry and heavy machinery. And
lithium cations stabilize mood swings in humans. Moving on to sodium, many sodium salts are widely
produced. Elemental sodium is a powerful reducing agent. Sodium vapor lamps are
efficient at converting electricity to light. Sodium hydroxide is utilized in drain cleaners.
And of course sodium chloride, or table salt, is indispensable in virtually every cuisine.
Potassium salts are vital to fertilizers. Both sodium and potassium ions are vital
to nerve transmission, as we learned about in the biopsychology series,
so potassium deficiency results in progressive paralysis. Also, potassium is a source of
radioactivity, as 0.01% of potassium is in the form of the radioactive 40K nuclide.
Rubidium provides the purple color in some fireworks.
Rubidium atomic clocks utilize a hyperfine transition in its ground state when interacting
with microwaves, and these are highly accurate. Cesium is also known for atomic clocks.
These utilize the same principle as the rubidium variety but are even more precise,
though they are more expensive. Additionally,
cesium salts are dense enough to be used to float away rock chips in undersea oil drilling,
and cesium layers on metal strips can be used to remove all traces of oxygen in sealed tubes.
Lastly, francium, as we mentioned is highly radioactive, and with a 22 minute half-life
for its most stable isotope, it has no practical applications. And with that we
wrap up our discussion of the first group on the periodic table, the alkali metals.