Question

Silver and rubidium both form +1 ions, but silver is far less reactive. Suggest an explanation, taking into account the ground-state electron configurations of these elements and atomic radii.

Short answer

The difference in reactivity between silver (Ag) and rubidium (Rb) ions can be explained by their ground-state electron configurations and atomic radii. Silver has an electron configuration of [Kr] 4d^10 5s^1, while rubidium has an electron configuration of [Kr] 5s^1. Due to its position further to the right in the periodic table, silver has a smaller atomic radius than rubidium. This results in a higher ionization energy for silver, as its valence electron is more tightly held by the nucleus. Consequently, silver is less likely to lose an electron and form a +1 ion, making it less reactive than rubidium.

Step-by-step solution

Determine ground-state electron configuration

To understand the atomic properties of silver (Ag) and rubidium (Rb), we will first find their ground-state electron configurations. We can do this using the periodic table to identify their atomic numbers and then filling out the electron orbitals accordingly. Silver has an atomic number of 47, and its electron configuration is [Kr] 4d^10 5s^1. Rubidium has an atomic number of 37, and its electron configuration is [Kr] 5s^1.

Analyze atomic radii

The atomic radii of silver and rubidium play a role in determining their ionization energies and reactivity. In general, as we move from left to right in a period of the periodic table, the atomic radii decrease due to the increasing nuclear charge attracting the electrons closer to the nucleus. Silver, being further to the right in the periodic table compared to rubidium, has a smaller atomic radius.

Understand ionization energy

Ionization energy is the energy required to remove an electron from an atom or ion. Generally, elements with high ionization energies are less reactive, as they are less likely to lose an electron and form a positive ion. Comparing silver and rubidium, we expect silver to have a higher ionization energy due to its smaller atomic radius and thus greater attraction of the valence electron to the nucleus.

Relate electron configuration and atomic radii to reactivity

The ground-state electron configurations of silver and rubidium suggest that both elements prefer to lose an electron to achieve a stable configuration. However, silver's smaller atomic radius means that its valence electron is more tightly held by the nucleus than rubidium's valence electron, resulting in a higher ionization energy. Consequently, silver is less likely to lose an electron to form a +1 ion, making it less reactive than rubidium.

Key concepts

Ionization Energy

Ionization energy is the energy required to remove an electron from an atom in its gaseous state. It's a crucial concept for understanding reactivity because atoms with high ionization energies tend not to lose electrons easily, and thus, they are less reactive. For instance, consider silver (Ag) and rubidium (Rb). Silver has a higher ionization energy due to the increased nuclear charge and the stability of its d-orbital filled with electrons. This higher energy barrier means it's harder to remove an electron from silver, making it less reactive compared to rubidium, which has a lower ionization energy.

Understanding the ionization energy helps explain the behavior of elements when forming ions or participating in reactions. It is influenced by factors such as atomic radii, nuclear charge, and the electron configuration of an atom. Therefore, when predicting the reaction patterns of elements, especially the tendency to form positive ions, ionization energy provides key insights.
For example, ionization energies generally increase across a period from left to right on the periodic table. This trend is because the addition of protons in the nucleus creates a stronger pull on the electrons, which in turn raises the ionization energy.

Atomic Radii

Atomic radii refer to the size of an atom, which is determined by the distance between the nucleus and the outermost electrons. This concept is intrinsically linked to reactivity because the size of an atom can influence how easily it can gain or lose electrons. In general, larger atomic radii make it easier to remove an electron since the valence electrons are further from the nucleus and less tightly bound by the nuclear charge.

In the context of our exercise comparing silver and rubidium, rubidium has a larger atomic radius and thus a lower ionization energy, contributing to its higher reactivity. The trend on the periodic table shows that atomic radii increase as you move down a group and decrease as you move across a period from left to right. This trend results from the balance between the increasing positive charge in the nucleus and the increasing number of electron shells as one moves down a group, compared to the increasing nuclear charge without additional electron shells as one moves across a period.

Ground-State Electron Configurations

The ground-state electron configuration of an element is the arrangement of electrons around the nucleus of an atom in the lowest energy state. This configuration impacts an element's chemical properties and reactivity. For example, silver has a ground-state electron configuration of [Kr] 4d^10 5s^1, meaning it has a completely filled d-subshell and a single electron in the s-orbital. Rubidium, on the other hand, has an electron configuration of [Kr] 5s^1, with only one electron in the s-orbital of the fifth shell.

A filled or nearly filled subshell gives an element increased stability, which results in higher ionization energy as observed in silver. The valence electron of silver is more tightly held because it is shielded by the fully filled d-subshell. Conversely, rubidium's lone valence electron is more easily removed as it's further away from the nucleus and not as shielded by inner electrons. This accounts for the differing reactivities of the two elements when they form +1 ions.

Periodic Table Trends

The periodic table organizes elements in such a way that trends in electron configuration, ionization energy, and atomic radius can be easily identified and correlated with element reactivity. As you move across a period from left to right, ionization energies generally increase, atomic radii decrease, and valence electrons become more tightly bound to the nucleus. This occurs due to the increasing number of protons in the nucleus, creating a stronger effective nuclear charge. Conversely, as you move down a group, the added electron shells increase the atomic radii and decrease ionization energy.

These trends are critical when comparing elements like silver and rubidium. Silver, being situated further right on the periodic table than rubidium, has a smaller atomic radius and higher ionization energy, reflecting a lower reactivity. By understanding periodic table trends, one can predict the relative reactivity of elements, which is pivotal in fields like chemistry and materials science for anticipating how elements will behave in various physical and chemical processes.