THE PERIODIC TABLE

Now that we've seen how electrons are bound to atoms, that they

occupy orbitals with strict rules about how many atoms occupy what kind of odd-

looking region around the nucleus, we're in a position to understand the layout of

the periodic table. That understanding will help us to use the table as a tool for

solving all kinds of problems, and for remembering facts that are easy to forget -

the table will make it easy.

First let's have a look at the periodic table. Click on the one below to print it. I'll

discuss some of its features below using smaller diagrams and pulled-out sections,

so you might want to refer to the full table often.

The first thing that probably strikes you about the table is its odd shape. You've

seen it before, but you've probably never understood why it's laid out like it is. It's

all about the electron configurations. Take a look at this table, in which each

element has been replaced by its electron configuration, to understand:

Look at the first (yellow) row in the table above. Hydrogen and Helium are

elements 1 and 2, respectively, and have those numbers of electron (we always

assume neutral atoms in the table). The electron configurations are H: 1s

and He:

. After that, the n=1 shell is full; it contains only an s-orbital and that can only

hold two electrons. I've placed He over next to H in this table just to show that it

lines up with all of the elements below it. We generally place it on the far right

(like the first table) for another reason (I'll get to that later).

Now the blue row: Here the n=2 row is being filled. First the s-orbital is filled with

two electrons (Li and Be), then across to the six p-

orbital electrons, ending with Neon, which has a full n=2 shell - no more electrons

will fit into the n=2 shell. Notice that He, Ne and all of the atoms below Ne have

full shells with increasing n. These atoms also happen to be the least

reactive atoms in the table - food for thought.

Now the n=3 shell (pink): First the 3s orbital is filled with two electrons (Na, Mg),

then we skip over to the 3p orbitals, filling them until we get to Argon (Ar). What

happens next is perfectly in keeping with the diagonal rule that we learned in

the last section. First the 4s orbital, because it is of lower energy than the 3d

orbitals, is filled (K and Ca), then we fill the 3d orbitals in the next row down.

The periodic table is arranged in order of orbital filling, according to the diagonal

rule. The first two columns fill s-orbitals. The rightmost six columns fill p-orbitals.

The middle group of ten fills the d-orbitals, and the Lanthanide and Actinide series

(block below the main table) fill the f-orbitals.

The periodic table is just another expression of the diagonal rule of electron-

orbital filling, and can be used to write the electron configuration of any atom, just

by reading left-to-right, top-to-bottom.

Now we can address the issue of the relative stability of atoms. It was already

noted that elements in the rightmost column of the table are especially non-

reactive. In fact, their electrons are tightly bound and difficult to remove. Note that

all of these atoms have eight electrons in their outermost (highest n) shell.

Sources: http://www.drcruzan.com/Chemistry_PeriodicTable.html