Periodic Trend: Ionization Energy (Simplified): Study with Video Lessons, Practice Problems & Examples

Ionization energy (IE) is the energy needed to remove an electron from a gaseous atom or ion, measured in kilojoules. It generally increases from left to right across a period and up a group in the periodic table. For instance, helium has a high IE of 2,372 kJ, while francium has a low IE of 393 kJ, indicating that francium's electrons are easier to remove. Noble gases exhibit high ionization energies due to their stable electron configurations, making them less likely to lose electrons.

Periodic Trend: Ionization Energy (Simplified) Concept 1

Now ionization energy, which is abbreviated as IE, represents the energy required to remove an electron from a gaseous atom or ion based on position in units of kilojoules. Now for example, let's take a look here. We have nitrogen in its gaseous state, so we're dealing with a gaseous atom. Ionization energy will be a reactant because we have to input energy into the gaseous atom in order to extract an electron. So the energy I just put in in terms of ionization energy helps me to remove an electron. Removing an electron from nitrogen, nitrogens are negatively charged, losing 1 means now nitrogen will be plus 1, still in its gaseous state, and the electron that I just removed will be a product next to it. So they've been separated from one another. So, again, when it comes to ionization energy, it's important to realize that ionization energy will be written as a reactant because we're inputting energy, and the electron we remove will be a product. Now what else can we say? Well, we can say here that the basic periodic trend is ionization energy will increase as we're moving from left to right across a period and going up a group. And it's important to realize that what type of ionization energy means that the electron is easily lost. Well, ionization energy is the energy to remove the electron. So if the energy required is very small, then it's going to be pretty easy to remove that electron. So we're going to say here that low ionization energies mean the electron can be removed very easily. On the flip side of that, if our ionization energy is very high, that means a lot of energy is required to remove an electron. That means that the electron will not be easily removed or easily lost. Now realize here that noble gases are perfect because they have the great electron arrangement or configuration, in their outer shells, so it would require a lot of energy to remove their electron. Because of this, noble gases will possess very high ionization energies. Alright. So just remember these fundamental principles when it comes to the idea of ion ionization energy.

Periodic Trend: Ionization Energy (Simplified) Concept 2

So remember, the general trend is as we move from left to right of a period and as we go up a group, your ionization energy will increase. This means that helium will possess the highest ionization energy at 2,372 kilojoules, whereas francium, which is on the exact opposite end, would have one of the lowest at 393 kilojoules. Remember, the higher the ionization energy is, the harder it is to remove that electron, and we can see that here. And we're going to see here that the smaller your ionization energy is, the easier it is to remove an electron. So Francium would be much easier to remove an electron than helium. Now, of course, there are exceptions in terms of this trend. In terms of those exceptions, you don't have to delve too much into it. We won't talk about, for the most part, those exceptions. Just realize that the general trend is as we're heading towards the top right corner of the periodic table, we expect our ionization energy to increase. Now you'll also see that some of these boxes are grayed out. That's because those are our heavy elements. Their atomic masses are so large, their atomic numbers are so large that they're pretty unstable. A vast majority of them have been synthesized within labs, and they only last mere moments in terms of existing. And because of that, we really don't talk about their ionization energies. So again, just remember, the general trend is as we head towards the top right corner, our ionization energy should be increasing.

Periodic Trend: Ionization Energy (Simplified) Example 1

Here it says, which of the following atoms has the smallest ionization energy? Remember, the general trend is as we're heading towards the top right corner, your ionization energy should be increasing. So if we take a look, we have phosphorus, we have fluorine, potassium, chromium, and bromine. We want the smallest ionization energy, so we're looking for something that's closer to the left side and lower down on the periodic table. Out of the choices presented, the one that fits that definition the best is potassium. It's the one most to the left and lowest down in terms of the periodic table, and therefore, has the smallest ionization energy.

Rank the following elements in order of increasing ionization energy:Br, F, Ga, K and Se.

• A. Ga < K < Br < Se < F
• B. Ga < Br < K < Se < F
• C. K < Se < Ga < Br < F
• D. K < Ga < Se < Br < F
• E. K < Ga < Se < F <  Br

Which of the following elements would lose an electron the easiest?

• A. Ar
• B. Li
• C. Zn
• D. Br
• E. B

Which element from Group 7A has lowest ionization energy.

• A. I
• B. Br
• C. Cl
• D. F
• E. Ne

Which of the following has the highest ionization energy? 

• A. Ar
• B. Na⁺
• C. Na
• D. Mg
• E. Kr

Here’s what students ask on this topic:

What is ionization energy and how is it measured?

Ionization energy (IE) is the energy required to remove an electron from a gaseous atom or ion. It is measured in kilojoules per mole (kJ/mol). For example, the ionization energy of nitrogen in its gaseous state involves inputting energy to remove an electron, resulting in a positively charged nitrogen ion and a free electron. The process can be represented as:

$N_{\left(g\right)}+\; IE\; \to N_{\left(g\right)}++\; e-$

Here, IE is the ionization energy required to remove the electron.

How does ionization energy change across a period and down a group in the periodic table?

Ionization energy generally increases as you move from left to right across a period and as you move up a group in the periodic table. This means that elements on the right side of a period, such as noble gases, have higher ionization energies compared to those on the left side. Similarly, elements at the top of a group have higher ionization energies than those at the bottom. For example, helium has a high ionization energy of 2,372 kJ/mol, while francium has a low ionization energy of 393 kJ/mol, indicating that francium's electrons are easier to remove.

Why do noble gases have high ionization energies?

Noble gases have high ionization energies because they possess a stable electron configuration with a full outer shell of electrons. This stability makes it energetically unfavorable to remove an electron, requiring a significant amount of energy. For instance, helium, a noble gas, has an ionization energy of 2,372 kJ/mol, one of the highest among all elements. This high ionization energy reflects the difficulty in removing an electron from a stable, fully occupied electron shell.

What is the significance of low ionization energy in chemical reactions?

Low ionization energy indicates that an electron can be easily removed from an atom, making the atom more reactive. Elements with low ionization energies, such as alkali metals, tend to lose electrons readily and form positive ions. This property makes them highly reactive, especially with nonmetals that have high electron affinities. For example, francium has a low ionization energy of 393 kJ/mol, making it very reactive and easily losing its outermost electron in chemical reactions.

What are some exceptions to the general trend of ionization energy in the periodic table?

While the general trend is that ionization energy increases from left to right across a period and up a group, there are exceptions. For instance, the ionization energy of oxygen is slightly lower than that of nitrogen, despite being to the right of nitrogen in the same period. This is due to electron-electron repulsion in the p-orbital of oxygen, making it easier to remove an electron. Similarly, the ionization energy of boron is lower than that of beryllium, even though boron is to the right of beryllium in the same period. These exceptions are due to the specific electron configurations and subshell arrangements of the elements.