Organizing the elements

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 periodic table is the key to unlocking the secrets of many chemical problems. Keep one close and use it.

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 Take a look at this table, in which each element has been replaced by its electron configuration, to understand: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 (green) 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: 1s1 and He: 1s2. 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 (magenta): 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.

Organization of the periodic table

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.


The octet rule

It is difficult in this section not to personify atoms, but it helps: Often it is said that atoms "want to have" eight electrons in their outer shell - the octet rule.

Nature always seeks the state of lowest energy. It's why a ball rolls downhill spontaneously, but never up. At right are the 1s, 2s and 2p orbital energies of a few atoms in the first and second rows of the table. Notice that as we add electrons the energies decrease (but remain in the same order). By the time we get to Neon, with 8 electrons in the n=2 shell, the atom so stable it is inert (doesn't react).

Another way of expressing the octet rule is to say that an atom tends to lose or gain electrons in order to have eight electrons in its outer shell, or full s- and p-orbitals. Note that the periodic table (and the diagonal rule) prescribe that the p-orbitals of any level are filled before its d-orbitals are filled.


The inert gases

The elements in the rightmost column of the table are called the inert gases. They all exist as gases at room temperature and they all have full outer-shell octets, which makes them unreactive (inert).

Columns in the periodic table are called groups, and rows are called periods. Elements in groups tend to have similar chemical properties (e.g. all gases, all inert).

Note that when the electron configuration gets cumbersome, it's OK to abbreviate the "core" configuration using another inert gas symbol, usually in [ ].

Now lets explore some of the other groups.


The halogens

The second group from the right contains the Halogen group. These atoms all have seven electrons in their outer shells, one short of a full octet and a much lower energy.

As a result, these atoms tend to "steal" electrons from other atoms; they have a high electron affinity ("like" for electrons). Thus, the halogen atoms tend to exist as -1 ions, F-, Cl-, Br-, I- and At-.

Now we're getting somewhere. We've used electron configuration and its manifestation in the periodic table to find a whole group of atoms that exist mainly as -1 ions.

On to some other interesting and important groups. ...


The chalcogens

The elements in the oxygen column are sometimes (not often) referred to as the chalcogens. They are usually just referred to as the oxygen group. They lack two electrons to have a full octet, so they tend to form -2 ions by picking up more weakly bound electrons from surrounding atoms.

These atoms tend to exist less as ions, however, and more as neutral atoms in molecules. There well be much more to say about this later.

Physically, the chalcogens are a mixed group: O is a gas, S is a nonmetallic solid, and Se - Po are metals. Polonium is radioactive, and found as the source of ionizing radiation in many smoke detectors.


The Alkali Metals

The first column of the periodic table (the first group) is the Alkali metal group. This group contains the simplest element, Hydrogen.

The hallmark of this group is that the electron configuration of each element is that of the inert gas just "behind" it in the table, plus one electron in the next higher s-orbital.

Because there is so much energetic stability to be gained by losing this electron (more properly, having it taken by an atom with more of a need for a spare electron), these atoms tend to lose one electron to become a +1 cation.

The Hydrogen cation, H+, plays a crucial role in many chemical systems. Note that Hydrogen contains only a proton in its nucleus (except for its isotopes 2H, or deuterium, and 3H, or tritium). That means the Hydrogen cation is just a bare proton.

When you see an atom from the first group, you can be reasonably sure that it's a +1 ion in almost any context.


Alkaline-Earth elements

The Alkaline Earth elements have two electrons more than the next lower inert gas, two more than the most stable octet available, so they tend to lose two electrons to become more stable. Thus Alkaline Earths tend to exist mainly as +2 cations.

In one sense, Helium belongs in this group. We classify it over with the inert gases because it has a full n=1 shell, and because it is quite inert. Helium is directly involved in only a few rare chemical reactions.

In the next few sections we'll take a look at periodic trends that can be gleaned from the table. Often we can use the table to tell whether one element expresses more of some property than another just from the relative positions of the two in the table.

Chemical similarity

Elements within a group (column) of the periodic table tend to have very similar chemical properties. The properties of an element arise mainly from the outer-shell electron configuration.

Other classifications

In the periodic tables below, elements have been color coded to help you learn about some of the classifications of atoms. We'll start with the so-called standard states of elements, the state (solid, liquid or gas) likely to be found at room temperature and standard atmospheric pressure.

Most elements are solids. The lightest and most stable are gases, like the inert gases. Only bromine and mercury (Hg) are liquids, although some of the solid elements melt at fairly low temperatures.


Metals, nonmetals, metalloids

The elements in this periodic table have been classified into six categories. Most elements have some kind of metallic character, which we'll define later. The metals in groups 1 and 2 tend to be soft and very reactive (some explosive when mixed with water!).

The transition metals are some of the best conductors of heat and electricity. The metalloids, which lie between metals and nonmetals, are interesting and very important because under the right conditions, they can be coaxed to have metallic or nonmetallic properties. This makes them particularly important in the electronics industry.


Biologically-relevant elements

A relative handful of elements are key components of biological systems on Earth. Carbon, Nitrogen, Oxygen, Phosphorus, Sulfur and Hydrogen are the backbone atoms of all proteins, nucleic acids, carbohydrates and lipids.

Ions like Sodium and Magnesium are crucial electrolytes in and between cells, and many metals are essential cofactors of enzymes, without which enzymes can't work properly.


From where do elements come?

Cosmological theory tells us that at the beginning of the universe, the Big Bang if you will, mostly H and He were created. Fluctuations in the density of the expanding gases caused some regions to coalesce and create more of an attractive force on the atoms around them - gravity. As these areas became increasingly dense, compressed and hot, stars were formed. In the intense heat and gravitational pressure of stars, smaller elements are fused to create larger ones. Our own sun fuses atoms up to Iron, but mostly fuses Hydrogen atoms into Helium.

Any atom in our solar system larger than Iron, and probably quite a bit of what's here that's smaller,

had to come from other, hotter stars over the eons. Certain large elements can probably only be formed in the hottest stars, the supernovae.

Scientist and great science communicator, Carl Sagan, used to say that "we are all made of starstuff." All of the elements that compose us and our world came from stars, some unimaginably far away.

Below is a graph of the distribution of relative abundances of the elements in our solar system. Note that it's on a log scale, so for example, there is roughly ten times (101) the amount of H as He.

solar system abundance

Source: Wikipedia Commons


In another section, we'll learn about other chemical trends that show up in the periodic table. We've only begun to get a feel for its usefulness. Keep yours with you and look at it once in a while. Keep it nearby when you do problems. Sleep with it. Talk to it. There are also a number of great periodic table apps available for your electronic devices.


How much of the periodic table should I memorize?

Your chemistry teacher might disagree with me, so be careful here, but my view is this:

I ask my students to memorize the first ten elements of the table: H, He, Li, Be, B, C, N, O, F and Ne. These are the elements at the top of each column (group), and they are generally representative of the other elements in the group. If you know these elements, then you've got a framework on which to build a memory of the other elements in the group as you study chemistry. Before long you'll assimilate a lot more elements into your long-term memory and what gets stored there will make some logical sense.

For example, all of the noble gas elements (under He) are much more similar than different. It's a more productive way, in my view, to build a functional memory of the elements.

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