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The Organization of the Periodic Table
The periodic table contains all of the known elements that come together to form all of
the matter in our universe. These elements are the different atoms that exist—these form
the “building blocks” that, when put together, make up the molecules that make up our
world.
When Mendeleev organized the periodic table by mass, he noticed that the elements’
physical and chemical properties exhibited periodicity. That is, these properties repeated
at regular intervals, and certain elements behaved similarly. From these observations, it
was a short leap to conclude that there was some connection between these elements that
led to them exhibiting similar “behaviors” or traits.
The periodic table is organized with purpose. First, the table is divided roughly in half, with
the left side comprising the metals and the rightmost elements making up the non-metals.
Between the metals and the non-metals are the metalloids, which are sort of in-between
and exhibit properties that are similar to both metals and non-metals.
Next, the table is divided into groups, which make up the different columns on the periodic
table. You are responsible for knowing four of these groups. On the left, Group I is referred
to as the Alkali metals. Group II is the Alkaline Earth metals. On the right, Group VII (17)
is referred to as the Halogens, and Group VIII (18) is referred to as the Noble Gases. These
groups are often referred to as “families”, either by name or by group number. They exhibit
similar properties—for example, the Noble Gases are non-reactive and gaseous at room
temperature, whereas the Alkali metals are some of the most reactive elements on the
periodic table.
Why do the Nobel gases exhibit similar behaviors? The answer lies in the structure of the
atom—specifically in the configuration of the atoms’ electrons. Electrons, which have a
negative charge, are found outside the nucleus of the atom. Some are closer to the nucleus,
some further from it. The outermost electrons—those furthest from the nucleus—are
called the Valence Electrons, “valence” coming from the Latin valentia, or “bravery”
(perhaps because they are most likely to react with other atoms?). Elements in the same
group (or family) have the same number of electrons in their outer energy level. That is
why they have similar chemical properties.
To understand this at a deeper level, we must explore why atoms react. Atoms react with
one another in order to fill their outermost energy level, and thereby increase their
stability. The only way for an atom with less than a full valence shell (or outermost energy
level) to reach stability is to gain enough electrons that it fills its shell, or by losing its
outermost electrons to reveal the full energy level beneath it. This is difficult to understand
from simply reading, but try to look at the picture below and imagine what is happening.
Figure 1 - Sodium Atom
Take Sodium (Na) as an example. Sodium is an Alkali metal, in Group I, and
therefore has one valence electron in its outer energy level. To become
stable, Sodium must lose this lone electron so that its outermost energy level
is full. (It’s much easier to do this than to gain seven more electrons—sort of
like if you could walk one block to the store or seven blocks. You’ll always
take the route that’s only one block, right?) It then has one more electron
than protons. This makes it an ion, a negative ion actually since it has more
electrons (-) than protons (+).
On the opposite side of the Periodic Table, you have Chlorine (Cl). A member of Group VII
(17), also known as the Halogens, Chlorine has seven valence electrons in its outer energy
level. Now, to become stable, do you think it will be easier for Chlorine to gain one electron,
or to lose its seven? If you guessed “gain one”, good work. You’re right. It requires very little
energy for Chlorine, or any of the Halogen group, to gain one electron and fill their
outermost valence shell. They become positive ions because they lose electrons.
Can you see now why Sodium and Chlorine work together so well to form NaCl (Sodium
Chloride)—also known as table salt? We’ll learn more about this later.
This line divides the metals
and the non-metals.
Notice that at the top of the Periodic Table are numbers—one for each group. The first two
groups, 1 and 2 are the Alkali metals and Alkaline Earth Metals. These have one and two
valence electrons (e-) respectively. Everything in the first column has one, every element in
the second column has two valence e-. This trend continues, but you have to “jump” across
the transition elements—the block beginning with Sc and ending with Zn in the middle, to
Group 13, where Boron sits with three valence electrons. Notice that the second digit of
13, or the Roman numeral number below the group number (III) tells you how many
valence electrons all of the atoms in the family have.
Carbon sits in Group 14 and has four valence electrons (note that it does not have 14! A
full outer shell has eight valence electrons!). Nitrogen, with-- guess how many!—five
valence electrons tops Group 15. So it goes for Oxygen, Chlorine, and finally Neon.
If sodium is very reactive, neon is the precise opposite. Neon, and all of its family members
in Group 18—the Noble Gases—is not reactive at all. Neon and the rest of its family are not
reactive because they have full outer energy levels, with a complete set of eight valence
electrons.