Ionization Energy

Now on to another property. It's called ionization energy. It can be defined as being 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.

In your lab work on atomic spectra you observed that a gas would conduct electricity and emit light when it was subjected to a high voltage. When there is little or no voltage applied to the gas in the tubes, no light is emitted and the gas does not conduct electricity. One method for measuring the ionization energy of a gas is to slowly increase the voltage applied to it until it does conduct electricity and emit light. The voltage at which that occurs can be used to calculate the ionization energy.

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.

This table shows the measured values for the ionization energies of the first twenty elements. If you take a close look at what happens to the ionization energy as you go from left to right across the periodic table, you will find that there is not really a steady increase in ionization energy as I had indicated. You could describe the pattern you see there as being a few steps forward then one step back, repeating itself as you move across. Progress is made, but it is not steady.
Ionization Energies (v)
H
13.6
  He
24.6
Li
5.4
Be
9.3
B
8.3
C
11.3
N
14.5
O
13.6
F
17.4
Ne
21.6
Na
5.1
Mg
7.6
Al
6.0
Si
8.2
P
10.5
S
10.4
Cl
13.0
Ar
15.8

 

The periodic nature of ionization energy is emphasized in this diagram. With each new period the ionization energy starts with a low value. Within each period you will notice that the pattern is really kind of a zigzag pattern progressing up as you go across the periodic table. 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. Nevertheless, as a general trend, from left to right across the periodic table, ionization energy does increase. Also as you go down the periodic table, the ionization energy does decrease for the reasons given. Graph of Ionization Energies of the Frist Twenty Elements

 

Note that 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.

 

Practice with Comparing Ionization Energies

Please take some time now to do the following exercises (also shown in example 7 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

Here are answers to the exercises above.

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.