Electron Configurations and the Periodic Table

In this section, we will look at the relationship between electron configurations and the shape of the periodic table. We will start with a short review of electron configurations, take a closer look at the patterns of electron configurations of various groups in the periodic table, and look at some short cuts for determining electron configurations from position on the periodic table.

Electron Configurations

In Lesson 5 electron configurations were introduced, but you did not need to learn how to write electron configurations. In this section we will go over the basics of electron configurations.

An electron configuration is, basically, an "address" that shows where the electrons can be found in an atom. We use the wave-mechanical model when creating electron configurations. Remember that the wave-mechanical model has atomic energy levels that are subdivided into "sublevels", which are further divided into "orbitals".

The farther the energy level is from the nucleus, the higher the energy and more electrons can fit into the larger space. The first energy level, the level closest to the nucleus, has space for only one sublevel - called the 1s sublevel - which holds one orbital. Each orbital, no matter which sublevel or energy level it is in, can hold a maximum of two electrons. So the first energy level can hold two electrons.

If an atom has more than two electrons, then they will be found in higher energy levels. In the second energy level there is more space than in the first energy level and we can see that there are two sublevels, called the 2s and 2p sublevels. The 2s sublevel has just one orbital, which can hold two electrons, but the 2p sublevel has three orbitals, which can hold a total of six electrons (two electrons per orbital). This continues on up as shown below.

Let's revisit this diagram from the previous lesson and use it to try and write the electron configuration of chlorine.

Chlorine has 17 electrons. Working up from the bottom, we can put 2 in the 1s sublevel, 2 in the 2s, 6 in the 2p, and 2 in the 3s. That leaves us with 5, all of which can fit in the 3p.

So our configuration is 1s² 2s² 2p⁶ 3s² 3p⁵

For atoms with just a few electrons, the order in which the electrons fill up the energy levels is exactly the way we would predict. However, once we get to the third energy level, things get a little strange. It turns out that it takes a little less energy for an electron to go into a 4s orbital than into a 3d orbital, and so the 4s fills up before electrons fill the 3d. This can be seen in the diagram at right, showing how the 4s sublevel is lower in energy (and thus easier for electrons to fill) than the 3d.

Writing Electron Configurations

When you write an electron configuration you write down where each electron can be found. Remember that for a neutral atom, the number of electrons must be equal to the atomic number (the number of protons). Electron configurations show each sublevel (1s, 2s, 2p, etc) and the number of electrons in each sublevel as a superscript after the sublevel.

For now, it will be useful to use the diagram at the right to help you write electron configurations. Start with the top arrow and follow it to its tip. Then move to the tail of the next arrow, and so on, writing each sublevel as you go. The order will be 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, etc.

Remember that s, p, d, and f sublevels hold 2, 6, 10, and 14 electrons each, respectively. While the diagram has "g" "h" and "i" sublevels, they will never need to be used.

Let's try the electron configuration of helium. The atomic number of helium is 2, so a neutral atom of helium must have 2 electrons. The electron configuration would be 1s².

Why don't you try your hand at a few others? Remember to fill each sublevel completely before moving on to the next sublevel. (You can leave a sublevel partially full if you "run out" of electrons before you fill it up.)

Write out the electron configurations for: Be, O, Ar, and Ca.

Answers:

Be has 4 electrons, 1s² 2s²

O has 8 electrons, 1s² 2s² 2p⁴

Ar has 18 electrons, 1s² 2s² 2p⁶ 3s² 3p⁶

Ca has 20 electrons, 1s² 2s² 2p⁶ 3s² 3p⁶ 4s²

Abbreviating Configurations

As you can already see, these configurations can get very long. Calcium has only twenty electrons and is already quite laborious to write out. If you go a few rows down to radium (88 electrons) the configuration is very long indeed: 1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d¹⁰ 4p⁶ 5s² 4d¹⁰ 5p⁶ 6s² 4f¹⁴ 5d¹⁰ 6p⁶ 7s². Surely there must be a more compact way to do this, especially considering that the ordering of sublevels is the same for all elements, so each subsequent one contains the configurations of earlier ones in it.

The standard way to abbreviate configurations is to use the noble (inert) gases as waypoints. As you should recall, the noble gases are the rightmost column of the periodic table, so they form a kind of "line break." The way we use them is this:

Calcium has a full configuration of 1s² 2s² 2p⁶ 3s² 3p⁶ 4s²
Going back from calcium, the noble gas that precedes it is argon. Argon's configuration is 1s² 2s² 2p⁶ 3s² 3p⁶
To abbreviate a configuration, we represent the configuration of argon by writing [Xx]. That is, if you see "[Ar]" in a configuration, you can read it as "1s² 2s² 2p⁶ 3s² 3p⁶"
So the abbreviated configuration of calcium is [Ar] 4s²

Applying this to other elements we practiced above, Be becomes [He] 2s², and oxygen is [He] 2s² 2p⁴. These may not seem like very substantial abbreviations. But for something like radon, the amount of effort saved is considerable. That whole mess up above just becomes [Rn] 7s².

Configurations and Table Structure

Historically, the properties of elements and the way those elements combined with other elements were the basis for the development of the periodic table (particularly the short form in the latter half of the 19th century). When physicists developed an understanding of how atoms were constructed (particularly the configuration of the electrons), it was possible to relate that structure to the periodic table. I really want to help you appreciate the deep beauty underlying the periodic table. It is truly remarkable the way our observations of the elements (their shared properties) match up with the conclusions of orbital theory so elegantly.

To begin, I would like you to take a moment to write out the electron configuration for these three elements: hydrogen, lithium and sodium. Hydrogen has one electron. Lithium has three electrons. Sodium has eleven electrons. So take a moment to do that before continuing on. Make sure you can get to the configurations below.

H: 1s¹

Li: 1s² 2s¹

Na: 1s² 2s² 2p⁶ 3s¹

Notice that these three elements are all in group IA of the periodic table. Notice also that each of those electron configurations ends in s¹. It is a different s¹ for each element - it is 1s¹ for hydrogen, 2s¹ for lithium, and 3s¹ for sodium - but notice the similarity in that they all end in s¹.

In the remaining parts of this section, we will take a closer look at the electron configurations of various groups in the periodic table, look at some short cuts for determining electron configurations, and look at how atomic orbitals are related to the periodic table.

Electron Configurations of Groups

In this subsection we will look at the patterns of electron configurations that exist in the periodic table--first for the representative elements, then the transition metals, and then the overall patterns.

Representative Elements

H, Li and Na all have their final electron put into an "s" sublevel; their electron configurations all end with s¹. If you look at the rest of the elements in that group you will see that the similarity continues on down the periodic table. All of the elements in group IA end in s¹ for their electron configuration. What is different about them is which level of s¹ they end with.

If you look at group IIA, you will see the same thing. Each of those elements has one more electron than the element before it, so those configurations will always end in s².

The electrons we are focusing on here are "valence" electrons: those in the outermost energy level of the atom. These are the most important electrons in any atom, as they determine how atoms bond and react.

If we expand our view to the rightmost six columns of the periodic table, we can see very clear patterns emerging.

These elements all have configurations that finish in a p subshell. As we move from left to right, that p subshell becomes more and more full.
As we move down a column, the configuration of the valence electrons stays the same, with just the energy level changing. This mirrors what we saw with the example of the alkali metals above.

Because of these highly regular patterns in electron configuration, the elements in the leftmost two groups (IA and IIA) and the rightmost six (IIIA - VIIIA) have highly predictable, periodically varying behavior. For this reason, these eight groups are referred to as the representative elements. Now let us look at other portions of the periodic table.

Transition Metals

The groups labelled with a B (IIIB, IVB, etc.) are known as the transition metals. Their electron configurations often do not follow the normal rules, and their behavior is more erratic than the representative elements. Let's take a look.

There are patterns here but the patterns are not as reliable. Let's start with element number 21, scandium. From calcium to scandium, the additional electron goes into the 3d¹. So scandium has 4s²3d¹, or if you prefer, 3d¹4s², as shown in example 2 in your workbook. You can write it in either direction, but it is 4s² and 3d¹. Then the next electron (the 22nd one) also goes in the 3d sublevel . Thus titanium, Ti, has 3d² as part of its electron configuration. Vanadium, then, is 3d³ and so on across except that you will notice that chromium does not have 3d⁴ like you might expect. We continue with manganese at 3d⁵ and 4s² and continue across with 3d⁶, 3d⁷, 3d⁸ but copper is not 3d⁹. You will not be expected to know what the exact electron configuration is for the transition elements when they alter that configuration a little bit. So when you are asked to figure out the electron configuration for a transition element, it probably will be one that follows the pattern, rather than one that doesn't.

The s, p, d, and f Blocks

The point here is to emphasize some of the patterns that exist in the relationship between the electron configurations of the elements and the location of the elements on the periodic table and even the shape of the periodic table. Looking at this example note how the periodic table can be broken up into s, p, d and f blocks.

The first two columns of the periodic table are in the s block because the elements in these two groups have their outermost electrons in s orbitals. The s block has two groups because atoms can put two electrons in an s sublevel.

The p block consists of groups IIIA through the inert gases. There are six groups in this block because atoms can put up to six electrons into a p sublevel.

The transition elements comprise the d block. The d block has 10 columns because up to 10 electrons will fit into a d sublevel.

The f block at the bottom of the periodic table has 14 columns because up to 14 electrons can fit into an f sublevel.

Remember the pattern we worked with that resulted in the electron configurations. 1s, then 2s, 2p, 3s, then 3p, 4s, then 3d, 4p, 5s. That resulted in electron configurations that looked like this. 1s², 2s², 2p⁶, 3s², 3p⁶, 4s², 3d¹⁰ and so on. The periodic table shows that same arrangement in a different way. The first period has 1s¹ then ². The second period has 2s¹ then 2s². Moving to the p block, we see 2p¹, 2p², up through 2p⁶. Similarly, the third period has 3s¹ then 3s², then 3p^1-6. The fourth period with the transition elements in the middle has 4s^1-2, followed by 3d^1-10 (though there are some oddballs) followed by 4p^1-6. The periodic table continues on showing the pattern dictated by the electron configurations.

This pattern can be used to write electron configurations using the periodic table as a guide (rather than the arrow method up above). This is useful because, while you will not be allowed the arrow diagram on quizzes and tests, you will be allowed a periodic table.

Let's see how this works in practice. For now, we will right unabbreviated electron configurations.

If I want to write the configuration for nitrogen, I start at the top of the periodic table. As I am at hydrogen, I am in the s block, and I am in the first row. This means I first fill the 1s subshell, which can hold two electrons, so my configuration starts with "1s² ..." and will continue until I get up to seven electrons (the number nitrogen has).

What comes next? From the 1s, I jump to the next block of the table. I am now at lithium, placing me in the s block, second row. So next comes the 2s subshell, with a maximum of two electrons, so my configuration now reads "1s² 2s²..." with three electrons still to be placed.

The next block of the table puts me at boron, now in the p block. I'm still in the second row, so I will put electrons in the 2p subshell next. A p subshell holds up to six electrons, but I only have three, so I finish my configuration as "1s² 2s² 2p³".

The main trip-up here is the d and f blocks. As the diagram above shows, they are offset from the s and b blocks. This means that I don't follow the 4s with 4d, but with 3d. And I don't follow 6s with 6f, but with 4f. The d block is offset down by one row, and the f block is offset down by two.

You should take some time now and practice writing full electron configurations for elements using just the periodic table to guide you. You can start out by using the one directly above with the blocks and rows labeled, but you must eventually get to a point where you can use the normal periodic table for the class.

I suggest working on elements further down the table, as they are trickier and will give you more practice, and also more opportunity to see the patterns in configurations. Once you get the order down, writing any configuration becomes pretty easy.

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