Question

Which element has the larger ionization energy? a. Li,N \(\quad\) b. Kr, Ne \(\quad\) c. \(\mathrm{Cs},\) Li

Short answer

N has a larger ionization energy than Li, Ne has a larger ionization energy than Kr, and Li has a larger ionization energy than Cs.

Step-by-step solution

Comparing Ionization Energies for Li and N

In this case, both elements are in the same period (period 2), so we can compare their atomic numbers. Li has an atomic number of 3, and N has an atomic number of 7. Since ionization energy increases across a period, N will have a larger ionization energy than Li.

Comparing Ionization Energies for Kr and Ne

In this case, both elements are in the same group (Group 18 - noble gases), so we can compare their positions down the group. Kr is below Ne in the group, which means it has a larger atomic number. Since ionization energy generally decreases down a group, Ne will have a larger ionization energy than Kr.

Comparing Ionization Energies for Cs and Li

In this case, both elements are in the same group (Group 1 - alkali metals), so we can compare their positions down the group. Cs is below Li in the group, which means it has a larger atomic number. Since ionization energy generally decreases down a group, Li will have a larger ionization energy than Cs. So the elements with larger ionization energies are: a. N b. Ne c. Li

Key concepts

Periodic Trends

Understanding periodic trends is essential when analyzing elements' properties such as ionization energy. These trends tell us how the properties of elements change across the periodic table. One key trend is that ionization energy generally increases as we move from left to right across a period. This is because atoms have more protons, thus a higher nuclear charge, which pulls the electrons closer and increases the energy required to remove an electron.

Conversely, ionization energy tends to decrease as we move down a group. When new electron shells are added, the outer electrons are further from the nucleus and are shielded by inner-shell electrons, reducing the nuclear charge's effect on the outermost electron and making it easier to remove.

Atomic Number

The atomic number is a fundamental characteristic of an element and has a direct effect on its ionization energy. It represents the number of protons in the nucleus of an atom. As the atomic number increases, so does the positive charge in the nucleus, pulling the electrons in closer.

This closer attraction means that more energy is required to remove an electron; hence, the ionization energy tends to be higher. However, it’s important to consider that the atomic number also tells us which period and group the element is placed in the periodic table, which allows us to apply the periodic trends to predict variations in ionization energy.

Alkali Metals

Alkali metals, found in Group 1 of the periodic table, are known for having low ionization energies compared to most other elements. This is because they have a single electron in their outermost shell.

As you descend the group from lithium (Li) to cesium (Cs), the ionization energy decreases. This is due to increased atomic radius, additional electron shells, and electron shielding. Despite their increasing atomic numbers, these factors make it easier to remove the outer electron, thus lowering the ionization energy.

Noble Gases

Noble gases are located in Group 18 of the periodic table and are unique because they have full electron shells, which makes them very stable and unreactive. They have the highest ionization energies within their respective periods because there is a strong attraction between the outer electrons and the nucleus.

However, when comparing noble gases like neon (Ne) and krypton (Kr), the ionization energy will decrease with an increase in atomic number and electron shell number. Despite this, the decrease in ionization energy as you move down the group is less pronounced with noble gases than with other elements, due to their full valence shells.