To answer last post's question of why the atomic radii shrink as more protons are added left to right on the periodic table, let’s again focus on the opposing charges between protons and electrons within an atom. Since these particles of opposing charges attract each other, they pull on each other. This is more so the case of protons pulling on electrons since a proton’s mass is so much greater than an electron’s negligible mass.

In the case of the decreasing atomic radii for periods, as one more proton is added, a greater positive force per additional proton is exerted on the negligible negative force from the additional electron. This collective attractive force contributed by a proton particle of more substantial mass than its electron counterpart, is capable of pulling the electrons closer to its center, thereby contracting the overall dimension of the atom.

So, why do the atomic dimensions increase in groups? Significant increases in proton attractive force are occurring, so yes, collective proton charge is also increasing...but so too are the repulsive forces among the very crowded electrons. And take a look at the figure: there is an increasing DISTANCE from the nucleus to the farthest flung electrons, effectively diminishing that proton attractive force. The repulsive forces therefore take over, the electrons spread out as much as possible, and to accommodate this the atomic dimension also expands!

Now comes the FUN after these technical discussions! I guess fun is all relative....

One of the patterns resulting from this organizational structure is that you intuitively think that the atom of an element must also physically grow in size if it’s adding electrons, right? Well, let’s take a closer look.
When we look at several elements in group IA, Sodium has 11 electrons, Potassium has 19, and Rubidium has 37. Each element has a significantly greater number of electrons than its preceding group member, and thus, requires an entirely new shell in which to house these electrons. Therefore, it increases by one orbital shell. And the atomic dimension does increase as a result.
When we look at a period, where we KNOW that each successive element has one more proton and one more electron (not discounting the neutron’s contribution to mass and size, but let’s simplify for this example), what we think might happen is the atom itself gets physically bigger at very tiny increments, since we’re adding in tiny increments.
In reality this is not the case! From left to right (e.g. n=5 and n=6), the atomic radius of each element decreases. Why? Answer, next post!

Next, let’s get back to our story of the periodic table, and look at the columns, called groups. An element within a group has the same number of VALENCE electrons, or as we learned before, the same number of electrons in whatever its outermost shell is. For example, look at group 8A where all the elements have 8 valence electrons. The difference between a group and the period is that we are talking about number of outermost electrons for each element in a group, versus the number of orbitals that house ALL of the electrons for each element in a period. In other words, with Sodium, we have 11 electrons, so, we need to use a spot in the 3s orbital to house the 11th electron. If you look at Argon, it too needs to utilize that 3rd orbital to house its last 8 of 18 electrons. Sodium needs only one slot for that 11th electron, versus Argon which needs to use 8 slots in the 3rd level orbital. However, both need to utilize that 3rd orbital shell.

Potassium has one more electron than Argon, and that is enough to bump that electron to the 4th level orbital shell. Potassium now differs from the previous period because it uses a 4th orbital shell. Yet it also has something in common with its group members...its valence electron number!

Remember, valence electron number is NOT the same as the total number of electrons for that element. For example, Sodium has a total of 11 electrons, but only 1 valence electron. Likewise, Cesium has only 1 valence electron, but in total has 55 electrons! Their commonality is having only 1 valence electron, and hence they react the same way chemically. And THAT is why they are group buddies.

We begin to see the brilliance of all the groundbreaking work of all the scientists who came before. It may look like a simple grid filled with letters, but it’s so much more than that.

The main patterns run across the table in left-to-right rows and up-to-down in columns. The rows are called periods, and stress the periodicity of the table, in that periodically, an element suddenly exhibits similar chemical properties to that of a previous element. For instance, Sodium is in the same period as Magnesium, Aluminum, Silicon, through to Argon. Then we get to Potassium, and boom! Two things happen: 1) Now Potassium has four atomic orbitals instead of three, the common factor in the previous period of elements. But even with this major physical change that forces us to start a new period and naturally forming a column with Sodium, chemically it happens to BEHAVE like Sodium! Therefore, each column, or group, has a pattern as well.

Each period is thus defined by the number of atomic orbitals that those elements possess. The red number to the left of each period indicates its “n” or orbital shell level number. So, if n=3, then each element in that period utilizes up to the 3rd level of atomic orbital shells. The shell itself may not be full, but the number of shells with occupying electrons is the same.

Mendeleev and his predecessors based their chart organization according to atomic weight and chemical properties, which were proven to have a strong correlation. However, as more elements were discovered, these two properties did not always accurately predict where the next element should fall; thus that strong relationship between atomic weight and chemical properties didn't always match.

Antonius van den Broeck was the first to hypothesize that there must be another factor that was more influential on the element’s identity other than atomic weight. He had no data however, to support his hypothesis.

So Henry Moseley devised an experiment based on the result of high-energy, negatively-charged electrons that are shot at elemental substances with each successive chemical element on the table. The X-ray emitted incrementally increased at a consistent rate. So if you look at the graph on the right side, this means that when he shot these electrons at Hydrogen, a certain X-ray emission occurred. Then when he shot it at the next known element, Helium, an equally unique X-ray emission resulted that was at a greater but consistent increase from Hydrogen. When he went on to Lithium, that X-ray emission increased again. When the square root was taken of these individual x-ray emissions, there was a beautiful linear correlation.

This graph is a simplified version explanation, as the emission quantity was NOT one-to-one as I’ve explained it, but the general concept is the same. What he did deduce from this is that another factor within the atom must be giving definition to the element, and as the quantity of that factor changed, so too did the identity of the element. This factor also correlated much more strongly to the organization of chemical property than atomic mass.

What was this factor? Atomic number, or, the number of protons, the positively-charged particle within the atom!