00:02
Atoms of other elements who do not have completed
outer shells would like to achieve this in
order to achieve a lower energetic state.
And they can do this either by swapping or
by sharing electrons. And this is the origins
for ionic or covalent bonding, as we will
see in subsequent lectures. For the first
18 elements, a filled valence shell usually
consists of eight electrons. That's obviously
because you can only complete—you can only
have—eight electrons in the outer shell
of the second principle quantum number. It's
only when you go up to the third shell—the
third principle quantum number—it's possible
to put 18 electrons in. Elements with higher
atomic numbers require 18. Elements in groups
1, 2, and 3 on the periodic table—so that
from the extract I showed you at the beginning—will
tend to lose electrons. Why? Well, because
they usually only have 1, 2, or 3 electrons
in the outer shell. It's easier to lose 1,
2, or 3 than to gain—in this case—7, 6,
or 5. Any elements which find themselves in
group 7 or group 6 will tend to want to gain
electrons in order to complete their outer
shell. Again, the reason for this is simple:
because it's easier for them to acquire two
electrons than it is for them to lose, in
the case of fluorine, 7, and in the case of
oxygen, 6. Now you may say, "Well, what do
you mean it is easier, or it's simpler?" But
the reality is, as you start stripping away
electrons, you, of course, are increasing
the size of your positive charge. So if, for
example, you move a single electron from an
outer shell of, let's say, lithium, and you
form an Li+, that's relatively straightforward.
But if you were to then… Let's say, for
the case of aluminum, you remove one electron;
you've created a disproportionality between
the charges. You now have fewer electrons
than you have protons. This increased nuclear
charge density technically pulls electron
density inwards and creates an ion, which
is small. Moving an electron away from that
shell again—so in other words, a second
electron—becomes more challenging, because
now you have a distribution of force, a distribution
of electrostatic attraction, which is greater
than before. So subsequent losses become more
and more difficult to achieve. In the case
of acquiring electrons—in the case of group
1 or group 2—of course, this means adding
greater electron negativity, a greater electron
density, onto those atoms, and electrons will
repel each other. This, of course, creates
a energetically unstable system.
So let's consider the formation of ions in
their simplest form. Let's look at lithium.
That's a group 1 or alkali metal on the periodic
table, and we can see, or we should note from
its electron configuration, possessing three
electrons, the first two go in the 1s—give
us 1s2— and the third one goes into the
2s, which gives us 2s1. The easiest way that
lithium can form a full outer shell is to
lose a single electron. And indeed, that's
how it exists. You tend to find that lithium
itself is very reactive, because it tends
to lose its electrons very quickly. In fact,
it is one of the most electropositive elements
in the periodic table. And so, in the case
of our ion, lithium, which has been formed,
it will now have the formal electron configuration
of 1s2, because now that electron from the
2s2… 2s1 orbital has been lost and now has
the same electron configuration as helium,
which, as we discussed earlier, is stable
by virtue of the fact it has a full outer
shell.
04:10
So if we look at group 17 elements or group
7 elements, such as fluorine, chlorine, bromine,
or iodine—halogens—we can see that they
have an electron configuration which is quite
the opposite to the lithium. They, indeed,
need to gain electrons in order to complete
their outer shell. If we take fluorine as
an example, can see we have two electrons
in the 2s2 subshell and five electrons in
the 2p subshell. By gaining a electron, it's
possible to complete the outer shell—the
second shell—by ensuring it now has eight
electrons in it. This, in the case of fluorine,
will generate the fluoride anion. Note the
difference: It's an anion, not a cation, because
it has gained an electron rather than lost
an electron. In this case, F– (the anion)
has the same electronic arrangement as neon.
05:08
That is to say it has a full second shell
(configuration being 1s2, 2s2, 2p6) and therefore
mimics that stable noble gas. So fundamentally,
it's important to understand that when it
comes to ionic chemistry, electrons are not
found on their own, and they must be attached
to an atom. And ions are also not found on
their own. It is not possible, for example,
to buy fluoride an ion. They must be countered
with a cation of some sort. Energetically,
this would be impossible.
So therefore, if we look at our simple example
of lithium and fluorine in the formation of
an ion, the lithium wishes to lose an electron
in order to achieve its low-energy state.
The fluorine wishes to gain an electron in
order to achieve a relatively low-energy state.
And so, therefore, an electron in the case
of a reaction of lithium and fluorine (not
that I recommend you do this; it would be
very exothermic) would be the loss of an electron
from the lithium outer shell and the gain
by the fluorine of that electron into the
fluorine outer shell. And this formal movement
of charge of electrons from one atom to the
other is the basis of ionic bond formation,
where there is the formal transfer of an electron
from one atom to the other. And the results
of this, as you can see here in this simple
example, is the formation of the lithium fluoride
salt. Salts typically exist in solid form
in crystal lattices. This is because (as we
may catch up with a little later on) the nondirectional
nature of this permanent electrostatic charge.
07:01
So in other words, the Li+ doesn't just attract
a single fluoride anion; it attracts any other
fluoride anions that happen to be within the
vicinity in a nondirectional fashion. Salts
are usually water-soluble, and this is by
virtue of the possibility of the ions to interact
by dipole bonding with water molecules, which,
as we'll see from a biological perspective
in some of the later lectures, is particularly
important when trying to come up with the
best agonist or antagonist for a given receptor
or enzyme.
07:40
So if we look at magnesium, for example, which
I alluded to earlier, it has two electrons
in its outer shell. When it loses these two
electrons, in order to give the complete octet
(the 2s2, 2p6), it affords us the cation Mg2+.
This is stable by virtue of the neon being
the resulting electron configuration, which
is a stable configuration. The same applies
with aluminum. Here, we're losing three electrons,
the 3s2 and 3p1, to afford us the aluminum
cation, yeah? So this also is stable as a
consequence of having a complete outer shell
in the form of the neon electron configuration.
However, if we can pair these fixed possible
ionizations to the transition metal or the
d-block region of the periodic table, we see
that there is a discrepancy. So as I said,
magnesium and aluminum either exist as their
elements or as 2+ and 3+ cations respectively.
However, in the case of the transition metal
or d-block elements, it is possible for them
to have multiple oxidation states or multiple
ionizations. If we look, for example, at copper
in its elemental form, it exists as Cu. In
its ionized form, it can either exist as Cu+
or Cu2+. Therefore, unlike magnesium, it can
either shed one of its outer shell electrons
or both of them. The same applies with tin.
09:20
It can either shed two of its outer shell
electrons to give Sn2+ or four to give Sn4+.
09:28
And it's this diversity which affords transition
metals a wide diversity and variety of ionic
reactions, which we will touch upon in later
lectures.
09:39
So salts made up from monoatomic cations and
anions are named in the following way (so
this is the nomenclature part): NaCl, sodium
chloride; LiF, lithium fluoride; and MgBr2,
magnesium bromide. So let's just touch upon
these in isolation. If we look at sodium chloride,
what we're saying here is that the cation
actually retains its original elemental name,
and it's only the group 7 (in this case) element
which has become an anion, which adopts the
suffix -ide, okay? So the anion takes i-d-e
at the end of its name. However, the cation
remains the same. Note: In the case of sodium
chloride, it is a one-to-one formula unit.
10:30
There is one positive charge on the sodium
and one negative charge on the chlorine. So
the ratio between the two is 1:1, otherwise
known as the stoichiometry. The same applies
with lithium fluoride. However, if we look
at magnesium bromide, we can see that there
are actually two negatively charged cat…
anions, Br–, in order to offset the positive
2+ charge of the magnesium. But we do not
call it magnesium dibromide. It is assumed
that you understand that magnesium can only
bind to two things which are monovalent, like
bromine. And remember, at all times, salts
must be charge-neutral, which is the reason
where you have the two anions, Br–; there
must be two of those to offset the 2+ charge
of the magnesium. And if we were to, for example,
look at aluminum chloride, it would have to
be AlCl3, where we have a aluminum ion with
a formal charge of 3+ and three chlorine anions,
each with a charge of 1–.