Periodic Trend: Ionization Energy (Simplified): Videos & Practice Problems

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

Ionization energy (IE) is defined as the energy needed to remove an electron from a gaseous atom or ion, measured in kilojoules. For instance, when considering nitrogen in its gaseous state, the process of ionization involves inputting energy to extract an electron. This can be represented as:

Energy + N(g) → N⁺¹(g) + e^-

In this reaction, nitrogen (N) starts as a neutral gaseous atom, and upon losing one electron, it becomes a positively charged ion (N⁺¹), while the removed electron (e^-) is a product of the reaction. It is crucial to note that ionization energy is treated as a reactant since energy must be supplied to facilitate the removal of the electron.

The periodic trend for ionization energy indicates that it generally increases as one moves from left to right across a period and from bottom to top within a group. This trend can be attributed to the increasing nuclear charge and decreasing atomic radius, which enhance the attraction between the nucleus and the electrons.

Low ionization energy signifies that an electron can be removed with relative ease, indicating that less energy is required for the process. Conversely, high ionization energy implies that a significant amount of energy is necessary to remove an electron, making it less likely for the electron to be lost. Noble gases exemplify this concept, as they possess stable electron configurations in their outer shells, resulting in very high ionization energies. Thus, it requires substantial energy to remove an electron from noble gases, reinforcing their stability.

Understanding these fundamental principles of ionization energy is essential for grasping the behavior of elements in chemical reactions and their placement in the periodic table.

Periodic Trend: Ionization Energy (Simplified) Concept 2

Ionization energy is a crucial concept in understanding the behavior of elements in the periodic table. As you move from left to right across a period and from bottom to top within a group, the ionization energy generally increases. This means that elements like helium, which has the highest ionization energy at 2,372 kilojoules per mole, require significantly more energy to remove an electron compared to elements like francium, which has one of the lowest ionization energies at 393 kilojoules per mole. The relationship here is clear: a higher ionization energy indicates that it is more difficult to remove an electron, while a lower ionization energy suggests that an electron can be removed more easily.

It is important to note that while this trend holds true for most elements, there are exceptions that may arise due to various factors, but these are generally not the focus of introductory discussions. Additionally, some elements, particularly heavy ones, are often grayed out in periodic tables due to their instability. These elements have large atomic masses and numbers, making them difficult to study as they exist only for brief moments in laboratory conditions. Consequently, their ionization energies are not typically discussed.

In summary, the overarching trend in ionization energy is that it increases as you approach the top right corner of the periodic table, reflecting the increasing difficulty of removing electrons from atoms in that region.

Periodic Trend: Ionization Energy (Simplified) Example 1

Ionization energy refers to the amount of energy required to remove an electron from an atom in its gaseous state. The periodic trend for ionization energy indicates that it generally increases as you move from left to right across a period and from bottom to top within a group on the periodic table. This is due to the increasing nuclear charge and decreasing atomic radius, which enhance the attraction between the nucleus and the electrons.

In the context of the elements provided—phosphorus, fluorine, potassium, chromium, and bromine—potassium is identified as having the smallest ionization energy. This is because potassium is located on the left side of the periodic table and is lower down in the group, which aligns with the trend of decreasing ionization energy as you move left and down. Therefore, potassium's position makes it the element with the least ionization energy among the options given.

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 elements would lose an electron the easiest?

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

Answer: (B)

Which element from Group 7A has lowest ionization energy.

• A. I
• B. Br
• C. Cl
• D. F
• E. Ne

Answer: (A)

Which of the following has the highest ionization energy? 

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

Answer: (B)

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.