Periodic Trend: Ionization Energy: Videos & Practice Problems

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.

Ionization Energy

Ionization energy refers to the energy required to remove an electron from a gaseous atom or ion, measured in kilojoules. For example, when considering nitrogen in its gaseous state, the removal of an electron transforms it into a positively charged ion. This process can be represented by the equation:

N(g) + Energy → N⁺(g) + e^-

In this reaction, energy is absorbed, indicating that ionization energy is always a positive value. This characteristic aligns with the definition of an endothermic reaction, where energy is taken in to break a bond. In this case, the bond being broken is the connection between the electron and the nitrogen atom.

Furthermore, it is important to understand that ionization energy is directly related to the potential energy of the electron being removed. As we delve deeper into the concepts of potential energy and ionization energy, we will explore how these ideas interconnect. For now, it is essential to recognize that the process of removing an electron results in the formation of a more positively charged species, highlighting the significance of ionization energy in chemical reactions.

Ionization Energy Trends

Ionization energy refers to the energy required to remove an electron from an atom in its gaseous state. A key trend in the periodic table is that ionization energy increases as you move from left to right across a period and from bottom to top within a group. This means that elements located towards the upper right corner of the periodic table, such as helium, exhibit the highest ionization energies, while those in the lower left corner, like francium, have the lowest.

For instance, francium has the lowest ionization energy due to its position, while helium, being in the top right corner, has the highest. However, there are notable exceptions to this general trend. For example, nitrogen has a higher ionization energy than oxygen, despite being in the same group, and beryllium has a higher ionization energy than boron. These exceptions highlight the complexity of ionization energy trends, but the overarching pattern remains that ionization energy increases as one moves towards the upper right of the periodic table.

Ionization Energy Example

Ionization energy refers to the amount of energy required to remove an electron from an atom in its gaseous state. The periodic trend indicates that ionization energy generally increases as you move from the bottom left to the top right of the periodic table. This is due to the increasing nuclear charge and decreasing atomic radius, which make it more difficult to remove an electron.

In the context of the elements provided—phosphorus, fluorine, potassium, chromium, and bromine—potassium is located in the lower left corner of the periodic table. As such, it has the smallest ionization energy among these options. This is because elements that are further to the left and lower down in the periodic table tend to have larger atomic radii and lower effective nuclear charge experienced by the outermost electrons, making it easier to remove them.

Thus, potassium, being the furthest left and lowest in the group, exhibits the smallest ionization energy, confirming the general trend observed in the periodic table.

Group 6A vs. Group 5A Elements

Ionization energy refers to the energy required to remove an electron from an atom. An interesting exception to the general trend of increasing ionization energy across a period occurs when comparing elements in the same period, specifically between group 5A and group 6A elements. Typically, one might expect that ionization energy increases as you move from left to right across the periodic table. However, this is not the case when comparing nitrogen (group 5A) and oxygen (group 6A).

The key to understanding this exception lies in the stability of the p subshell orbitals. P orbitals achieve maximum stability when they are either half-filled or fully filled. In the case of nitrogen, which has the electron configuration of 1s² 2s² 2p³, its three p electrons are arranged in a half-filled state. This configuration is particularly stable. When one electron is removed from nitrogen, it transitions to the configuration of helium (1s²), resulting in a less stable arrangement of 2s² 2p².

On the other hand, oxygen, with the electron configuration of 1s² 2s² 2p⁴, has one more electron than nitrogen. This additional electron disrupts the half-filled stability of the p orbitals. When oxygen loses one electron, it achieves a half-filled configuration of 2s² 2p³, which is more stable. Consequently, because oxygen is more willing to lose this electron to reach a more stable state, it requires less energy to remove it compared to nitrogen. Thus, oxygen exhibits a lower ionization energy than nitrogen despite being further along in the periodic table.

In summary, the stability of p orbitals plays a crucial role in determining ionization energy, with half-filled and fully filled configurations being the most stable. This principle explains why oxygen has a lower ionization energy than nitrogen, despite its position in the periodic table.

Group 3A vs. Group 2A Elements

Ionization energy refers to the energy required to remove an electron from an atom. An interesting exception to the general trend of increasing ionization energy across a period occurs between group 2A and group 3A elements. Typically, one would expect that as you move from left to right across a period in the periodic table, the ionization energy increases. However, this is not the case when comparing Beryllium (Be) from group 2A and Boron (B) from group 3A.

Beryllium has the electron configuration of Helium (He) 2s², meaning its 2s orbital is completely filled. This full s subshell configuration is particularly stable, making it more difficult to remove an electron. When an electron is removed from Beryllium, it transitions to a less stable state of 2s¹.

In contrast, Boron has the electron configuration of Helium 2s² 2p¹. If Boron loses its single 2p electron, it achieves the stable configuration of Helium 2s². This transition results in a more stable electron arrangement, as the remaining 2s orbital is fully filled. Consequently, the ionization energy of Boron is lower than that of Beryllium, despite Boron being further to the right in the periodic table.

This phenomenon highlights the importance of electron configuration and stability in determining ionization energy, particularly the stability associated with fully filled s orbitals.

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

Answer: (D)

Which of the following has the highest ionization energy? 

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

Answer: (B)

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

Answer: (D)

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.