Periodic Trends

If you were to look carefully at many of the properties of the elements, you would notice something besides the similarity of the properties within the groups. You would notice that many of these properties change in a fairly regular fashion that is dependent on the position of the element in the periodic table. And of course, that is what you will do next. As you compare elements from left to right across the periodic table, you will notice a trend or regular change in a number of properties. The same thing happens if you go up and down on the periodic table and compare the properties of the elements. In this section, we will look at a number of basic properties of the elements and see how they change regularly from left to right and top to bottom of the periodic table. We will also relate these properties to patterns in how elements react.

Note that everything we will discuss in this section only applies to atoms in their pure elemental form.

Atomic Size

The first of these properties is the atomic size. You know that each atom has a nucleus inside and electrons zooming around outside the nucleus. It should seem reasonable that the size of an atom depends on how far away its outermost (valence) electrons are from the nucleus. If they are very close to the nucleus, the atom will be very small. If they are far away, the atom will be quite a bit larger. So the atomic size is determined by how much space the electrons take up.

Measuring the size of atoms is, in some ways, like measuring the size of cotton balls. The value you get depends on the conditions under which they are measured. A "free" cotton ball has a different size than when it is in the package surrounded by many other cotton balls. Different values for the sizes of atoms are obtained depending on both the method used and the conditions in which the atoms find itself. However, no matter what method we use, as long as we are consistent, very clear trends emerge. The table below shows the sizes of atoms found using the "covalent radius" approach, measured in picometers.

Let's make some comparisons in a family and in a period. As you go down a group, like from hydrogen to lithium to sodium on down, the atomic size increases. As you go across a period, as from sodium to argon, notice that the size decreases. You need to memorize those trends. (note: there are some hiccups in these trends, such as neon being larger than fluorine and some odd fluctuations in the transition metals. You do not need to memorize these anomalies)

Now let's talk about why that's the case and relate it back to the various factors presented earlier. Remember that the nuclear charge and the shielding electrons combine to make the effective nuclear charge. That is a very important factor when you are comparing elements in a period. As you go across a period, the nuclear charge increases and the number of energy levels stays the same. Consequently, the number of shielding electrons stays the same and the effective nuclear charge increases. As the effective nuclear charge increases, it pulls the electrons in closer and closer to the nucleus. So as you go across a period, the increase in the nuclear charge causes a decrease in the atomic size because the electrons in the valence energy level are pulled in closer and closer.

Now let's make comparison within a group such as hydrogen down to cesium (Cs). It is true that the nuclear charge is increasing, but so is the number of shielding electrons. The number of shielding electrons increases by the same amount that the nuclear charge increases. So the effective nuclear charge felt by the valence electrons stays the same. There is no increase in the effective nuclear charge but there is an increase in the number of energy levels that are being used. Consequently, the electrons in the valence energy level are farther and farther away from the nucleus because they are in higher energy levels. Consequently, the important factor in a vertical comparison on the periodic table is the number of energy levels that are being used because the increase in the number of shielding electrons cancels out the increase in the nuclear charge.

To summarize, as you go across a period, the increase in the nuclear charge is the most important factor because the number of energy levels stays the same. As you go down a group, the increase in shielding electrons more or less cancels out the increase in nuclear charge, leaving the increase in the number of energy levels as the most important factor. This is true not only for atomic size but for other properties as well.

Atomic radius increases as you go down the table and decreases as you go to the right. You should both know this fact and be able to explain the reasons behind it.

Practice with Comparing Atomic Size

Now try your hand at answering the following questions (also shown in exercise 9 in your workbook). Check your answers below and then continue with the lesson.

For each of the following sets of atoms, decide which is larger, which is smaller, and why.

a. Li, C, F

b. Li, Na, K

c. Ge, P, O

d. C, N, Si

e. Al, Cl, Br

Answers for Comparing Atomic Sizes

a. Li, C, F: All are in the same period and thus have the same number of energy levels. Therefore, the important factor is the nuclear charge. Li is the largest because it has the smallest nuclear charge and pulls the electrons toward the nucleus less than the others. F is the smallest because it has the largest nuclear charge and pulls the electrons toward the nucleus more than the others.

b. Li, Na, K: All are in the same group and thus have the same effective nuclear charge. Therefore, the important factor is the number of energy levels. Li is the smallest because it uses the smallest number of electron energy levels. K is the largest because it uses the largest number of electron energy levels.

c. Ge, P, O: All are in different groups and periods, therefore both factors must be taken into account. Fortunately both factors reinforce one another. Ge is the largest because it uses the largest number of energy levels and has the smallest effective nuclear charge. O is the smallest because it uses the smallest number of energy levels and has the largest effective nuclear charge.

d. C, N, Si: Not all are in the same group and period, so, again, both factors must be taken into account. C and N tie for using the smallest number of energy levels, but N has a higher effective nuclear charge. Therefore, N is the smallest. C and Si tie for having the lowest effective nuclear charge, but Si uses more energy levels. Therefore, Si is the largest.

e. Al, Cl, Br: Not all are in the same group and period, so, again, both factors must be taken into account. Cl is the smallest because it has higher effective nuclear charge than Al and uses fewer energy levels than Br. Which is largest is less straightforward. Al has a lower effective nuclear charge (by four), but Br uses more energy levels (by one). Because the difference in effective nuclear charge is larger, it should be the more important factor in this case, making Al the largest.

Al and Br can also be compared to one another indirectly by comparing both to Cl. Both Al and Br are larger than Cl. Al is larger than Cl because it has lower effective nuclear charge (by four). Br is larger than Cl because it uses more energy levels (by one). Because Al is larger than Cl by four "steps" and Br is larger than Cl by only one "step", Al is likely the largest of the three.

Ionization Energy

Now on to another property. It's called ionization energy. It can be defined as the energy required to remove the outermost electron from a gaseous atom. A "gaseous atom" means an atom that is all by itself, not hooked up to others in a solid or a liquid. When enough energy is added to an atom the outermost electron can use that energy to pull away from the nucleus completely (or be pulled, if you want to put it that way), leaving behind a positively charged ion. That is why it's called ionization; one of the things formed in the process is an ion. The ionization energy is the exact quantity of energy that it takes to remove the outermost electron from the atom.

If the ionization energy is high, that means it takes a lot of energy to remove the outermost electron. If the ionization energy is low, that means it takes only a small amount of energy to remove the outermost electron.

Let’s use your understanding of atomic structure to make some predictions. Think for a minute about how ionization energy would be affected by three of the factors we were talking about earlier: (1) nuclear charge, (2) number of energy levels, and (3) shielding.

As the nuclear charge increases, the attraction between the nucleus and the electrons increases and it requires more energy to remove the outermost electron and that means there is a higher ionization energy. As you go across the periodic table, nuclear charge is the most important consideration. So, going across the periodic table, there should be an increase in ionization energy because of the increasing nuclear charge.

Going down the table, the effect of increased nuclear charge is balanced by the effect of increased shielding, and the number of energy levels becomes the predominant factor. With more energy levels, the outermost electrons (the valence electrons) are further from the nucleus and are not so strongly attracted to the nucleus. Thus the ionization energy of the elements decreases as you go down the periodic table because it is easier to remove the electrons. Another way of looking at that is that if you are trying to take something from the first energy level, you have to take it past the second, the third, the fourth and so on, on the way out. But if something is already in the third or fourth energy level, it doesn't have to be taken as far to get away from the nucleus. It is already part way removed from the nucleus.

The zigs and zags on that graph correspond to the sublevels in the energy levels. So far in this lesson we have presumed that all the electrons in the second energy level are pretty much the same. Two factors make that not completely true. One factor is that because s and p orbitals have different shapes, the electrons in p orbitals have more energy and are further from the nucleus. The other factor is that when electrons are paired up in an orbital, they repel one another somewhat. Those two factors account for the zigzag nature of the increase in ionization energy.

The trends in the periodic properties of atomic size and ionization energy are related. Going across the periodic table from left to right, the electrons are more tightly held by the nucleus, causing the atoms to be smaller and the ionization energy to be higher. As you go down the periodic table, the electrons are further from the nucleus, causing the atoms to be larger and the ionization energies to be lower.

Ionization energy decreases as you go down the table and increases as you go to the right. You should both know this fact and be able to explain the reasons behind it.

Practice with Comparing Ionization Energies

Please take some time now to do the following exercises (also shown in example 10 in your workbook). When you have done that, check your answers below and then continue.

For each of the following sets of atoms, decide which has the highest and lowest ionization energies and why.

a. Mg, Si, S

b. Mg, Ca, Ba

c. F, Cl, Br

d. Ba, Cu, Ne

e. Si, P, N

Answers to Comparing Ionization Energies

a. Mg, Si, S: All are in the same period and use the same number of energy levels. Mg has the lowest I.E. because it has the lowest effective nuclear charge. S has the highest I.E. because it has the highest effective nuclear charge.

b. Mg, Ca, Ba: All are in the same group and have the same effective nuclear charge. Mg has the highest I.E. because it uses the smallest number of energy levels. Ba has the lowest I.E. because it uses the largest number of energy levels.

c. F, Cl, Br: All are in the same group and have the same effective nuclear charge. F has the highest I.E. because it uses the smallest number of energy levels. Br has the lowest I.E. because it uses the largest number of energy levels.

d. Ba, Cu, Ne: All are in different groups and periods, so both factors must be considered. Fortunately both factors reinforce one another. Ba has the lowest I.E. because it has the lowest effective nuclear charge and uses the highest number of energy levels. Ne has the highest I.E. because it has the highest effective nuclear charge and uses the lowest number of energy levels.

e. Si, P, N: Si has the lowest I.E. because it has the lowest effective nuclear charge and is tied (with P) for using the most energy levels. N has the highest I.E. because it uses the fewest energy levels and is tied (with P) for having the highest effective nuclear charge.