Periodic Trend: Ionization Energy - Online Tutor, Practice Problems & Exam Prep

Ionization energy is the energy required to remove an electron from a gaseous atom or ion, resulting in a positive charge and an increase in potential energy. This process is endothermic, meaning energy is absorbed. As we move across a period from left to right and up a group in the periodic table, ionization energy increases, with helium having the highest and francium the lowest. Exceptions to this trend include the stability of half-filled p orbitals and fully filled s orbitals. For instance, nitrogen has a higher ionization energy than oxygen because its half-filled p orbitals are more stable, making it less willing to lose an electron. Similarly, beryllium has a higher ionization energy than boron due to the stability of its fully filled s orbitals. These exceptions highlight the nuanced stability factors that influence ionization energies within the periodic trends.

Periodic Trend: Ionization Energy

Now ionization energy is the energy absorbed to remove an electron from a gaseous atom or ion based on position in kilojoules. Now we're going to say here we have nitrogen in its gaseous state. We need to remove its electron. When I remove its electron it becomes positively charged. And here we would illustrate the electron that's been moved by placing it also as a product.

We said energy is absorbed, so we're taking in energy, we're gaining energy. So ionization energy will always be a value that's greater than 0. It'll be a positive number. Now what we also need to realize is since we're absorbing energy, oftentimes you'll hear this being called an endothermic reaction, and these reactions energy is absorbed in order to break a bond.

Here we're breaking the connection of the electron to the element. Also, what we need to realize here is that ionization energy equals the potential energy of a given electron. Later on we'll try to connect this whole idea of potential energy to ionization energy, but for now, just realize we're removing an electron from a gaseous ion or atom to create a more positively charged species.

Periodic Trend: Ionization Energy

In terms of ionization energy, the periodic trend is ionization energy will increase as we're moving from left to right across a period and up a group. So as we're heading towards the upper right corner of the periodic table, ionization energy goes up.

Here, if we take a look, we can see that the one with the lowest ionization energy is francium, which makes the most sense 'cause it's on the exact opposite pattern of the periodic trend. And we can see here that the one with the highest is helium. It's to the most top right corner of the periodic table.

Now, of course, there's going to be a little bit of some exceptions to this trend. We'll go into further detail about them. You'll notice here that nitrogen has a lower a higher ionization energy than oxygen. And you also notice here that beryllium has a higher ionization energy than boron.

There are other little subtle exceptions, but these four elements are the most commonly discussed when it comes to exceptions in ionization energy. So just realize that overall, as we head to the top right corner of the periodic table, ionization energy goes up.

Periodic Trend: Ionization Energy 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 for 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 to have the smallest ionization energy.

Periodic Trend: Ionization Energy

Now when it comes to exceptions in ionization energy, one of them we can say is when in the same. Or row, group 6 elements have a lower ionization energy than group 5 element. The explanation is that P subshell orbitals are most stable when half filled or totally filled with electrons.

If we look here, we have nitrogen versus oxygen. Oxygen is more to the right of the periodic table, so you would assume it should have a higher ionization energy, but in fact it doesn't. The reason is here. Nitrogen is already in a stable state, its P orbitals are half filled, right? We have 3 electrons here. When I remove one electron from nitrogen it becomes helium 2S22P2 and now as a result my P orbitals are no longer half filled and therefore not as stable.

For oxygen, on the other hand, here it's almost half built for the 2P2 orbitals, but we have one too many electrons. Oxygen will be willing to give up its electron to become 2S22P3 and in that way it becomes half filled for its P orbitals and more stable. Since oxygen is freely willing to give up these, this electron less energy is required to remove it. That's why it's ionization energy would be lower than nitrogen.

So remember, P orbitals are most stable when they're either half filled or totally filled.

Periodic Trend: Ionization Energy

Now, another exception that exists with ionization energy has to do with the fact that when in the same row, group 3A elements have a lower ionization energy than group 2A elements. The reason here being that S subshell orbitals are most stable when totally filled. So the most common example for this is between beryllium and boron.

Boron is more to the right of the periodic table, so you would expect it to have a higher ionization energy, but in fact that's not true here. Beryllium is already in a stable state. Its S orbital electrons are completely filled in, so this orbital is totally filled in and it's very stable. That means if I come in and try to remove an electron, it becomes 2S1. We're no longer in as stable a state as we want. Remember, S orbital electrons are most stable when they're totally filled in.

Boron, on the other hand, is He22S22P1. If we could lose this one electron from 2P1, all we'd be left with is a 2S2 orbital, which is very stable, so that's what's going to happen. So it becomes He22S2 once I remove that electron, and now I have an S set of orbital electrons that are totally filled in and therefore it's more stable as a result. This means that boron has a lower ionization energy than beryllium.

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 has the highest ionization energy? 

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

The energy of an electron in a one-electron atom or ion equals (–2.18 x 10^–18 J) (Z²/n²). Estimate the ionization energy for the valence electron of the Li atom and compare it to its theoretical value of 520 kJ/mol.

• A. 491 kJ/mol
• B. -179 kJ/mol
• C. -1682 kJ/mol
• D. -2953 kJ/mol
• E. -5062 kJ/mol

Here’s what students ask on this topic:

What is ionization energy?

Ionization energy is the amount of energy required to remove an electron from a gaseous atom or ion in its ground state. It's a fundamental concept in chemistry and physics that provides insight into the reactivity and chemical behavior of elements. When an electron is removed from an atom, it results in the formation of a positively charged ion, known as a cation. The ionization energy is usually expressed in electronvolts (eV) or kilojoules per mole (kJ/mol).

The first ionization energy refers to the energy needed to remove the first electron from a neutral atom. Successive ionization energies are defined for the removal of additional electrons, and these energies increase as more electrons are removed because it becomes harder to remove an electron from a more positively charged ion. The ionization energies of elements are influenced by the atomic radius, nuclear charge, and electron configuration. Elements with low ionization energies tend to be more reactive and are more likely to form positive ions in chemical reactions.

Why does ionization energy increase across a period?

Ionization energy is the energy required to remove an electron from an atom in its gaseous state. As you move across a period in the periodic table, the atomic number increases, which means there are more protons in the nucleus of each successive element. This results in a stronger positive charge in the nucleus, which attracts electrons more strongly.

Simultaneously, as you move across a period, electrons are being added to the same principal energy level (or shell). Because the inner shells shield the outer electrons from the nucleus only slightly, the effective nuclear charge—the net positive charge felt by the outermost electrons—increases. This makes it harder to remove an electron since they are being held more tightly by the increased nuclear charge.

Moreover, the atomic radius decreases across a period because the increased nuclear charge pulls the electron cloud closer to the nucleus, further increasing the ionization energy. Therefore, it requires more energy to overcome this attraction and ionize the atom, hence the increase in ionization energy across a period.

Why does ionization energy decrease down a group?

Ionization energy is the energy required to remove an electron from an atom in its gaseous state. As you move down a group in the periodic table, the ionization energy decreases. This is because atoms get larger with an increasing number of electron shells as you go down a group. The outermost electrons are farther from the nucleus and are also shielded by the inner layers of electrons. This increased distance and electron shielding reduce the force of attraction between the nucleus and the outermost electrons. Consequently, it becomes easier to remove an electron since it is less tightly bound to the atom. Therefore, less energy is required to ionize the atom, leading to a decrease in ionization energy as you move down a group.