Question

Write electron configurations for the most stable ion formed by each of the elements Rb, Ba, Se, and I (when in stable ionic compounds).

Short answer

The electron configurations of the most stable ions formed by the elements Rb, Ba, Se, and I are: Rb+: \(1s^2 2s^2 2p^6 3s^2 3p^6 4s^2 3d^{10} 4p^6 5s^1\) Ba2+: \(1s^2 2s^2 2p^6 3s^2 3p^6 4s^2 3d^{10} 4p^6 5s^2 4d^{10} 5p^6\) Se2-: \(1s^2 2s^2 2p^6 3s^2 3p^6 4s^2 3d^{10} 4p^6\) I-: \(1s^2 2s^2 2p^6 3s^2 3p^6 4s^2 3d^{10} 4p^6 5s^2 4d^{10} 5p^6\)

Step-by-step solution

Determine the charge of the ion formed by each element

For each element, we need to find the charge for the most stable ion formed when they are in stable ionic compounds. For Rb (Rubidium) and Ba (Barium): They belong to group 1 and group 2 respectively, so they are both alkali metals. Rb will lose its single valence electron, and Ba will lose its two valence electrons. This gives Rb a +1 charge and Ba a +2 charge. For Se (Selenium) and I (Iodine): They belong to group 16 and group 17 respectively, so they are both nonmetals. Se will gain two electrons, while I will gain one electron. This gives Se a -2 charge and I a -1 charge.

Write the electron configurations for the most stable ions

Rb: Its atomic number is 37, so the neutral atom has 37 electrons. When it loses one electron (to form Rb+), it will have 36 electrons. So the electron configuration of Rb+ is: \[1s^2 2s^2 2p^6 3s^2 3p^6 4s^2 3d^{10} 4p^6 5s^1\] Ba: Its atomic number is 56, so the neutral atom has 56 electrons. When it loses two electrons (to form Ba2+), it will have 54 electrons. So the electron configuration of Ba2+ is: \[1s^2 2s^2 2p^6 3s^2 3p^6 4s^2 3d^{10} 4p^6 5s^2 4d^{10} 5p^6\] Se: Its atomic number is 34, so the neutral atom has 34 electrons. When it gains two electrons (to form Se2-), it will have 36 electrons. So the electron configuration of Se2- is: \[1s^2 2s^2 2p^6 3s^2 3p^6 4s^2 3d^{10} 4p^6\] I: Its atomic number is 53, so the neutral atom has 53 electrons. When it gains one electron (to form I-), it will have 54 electrons. So the electron configuration of I- is: \[1s^2 2s^2 2p^6 3s^2 3p^6 4s^2 3d^{10} 4p^6 5s^2 4d^{10} 5p^6\] To summarize, the electron configurations of the most stable ions formed by the elements Rb, Ba, Se, and I are: Rb+: \(1s^2 2s^2 2p^6 3s^2 3p^6 4s^2 3d^{10} 4p^6 5s^1\) Ba2+: \(1s^2 2s^2 2p^6 3s^2 3p^6 4s^2 3d^{10} 4p^6 5s^2 4d^{10} 5p^6\) Se2-: \(1s^2 2s^2 2p^6 3s^2 3p^6 4s^2 3d^{10} 4p^6\) I-: \(1s^2 2s^2 2p^6 3s^2 3p^6 4s^2 3d^{10} 4p^6 5s^2 4d^{10} 5p^6\)

Key concepts

Ions

In chemistry, ions are atoms or molecules that have gained or lost one or more of their electrons. This means ions have a net charge, either positive or negative. Atoms become ions in order to achieve a stable electron configuration, similar to that of the nearest noble gas.
When an atom loses electrons, it becomes a positively charged ion, also known as a cation. Conversely, when it gains electrons, it becomes negatively charged, known as an anion. The main goal for these atoms is to achieve stability by fulfilling the octet rule, which states that atoms tend to have eight electrons in their valence shell.

• Cation: Positive charge due to electron loss (e.g., Na⁺).
• Anion: Negative charge due to electron gain (e.g., Cl^-).

Alkali Metals

Alkali metals are located in Group 1 of the periodic table. These include elements like lithium (Li), sodium (Na), and rubidium (Rb).
These metals have a single valence electron which makes them highly reactive. The inner shells are full, and only the outermost shell needs one more electron to complete the octet.
Due to their desire to achieve a stable electronic configuration with a full outer shell, alkali metals readily lose their single valence electron when they form compounds. This process leads them to form +1 charged ions.
Some noteworthy properties of alkali metals include:

• High reactivity with water and oxygen.
• Soft texture and low densities.
• High tendency to form cations by losing one electron.

Electron Gain

Electron gain refers to the process where an atom captures extra electrons to form a negative ion or anion. This typically occurs with nonmetals, which strive to complete their valence electron shell, achieving stability through the octet rule.
In the periodic table, elements in Group 16 and Group 17, such as selenium (Se) and iodine (I), are known for gaining electrons.
Selenium gains two electrons to form a Se^2- anion, while iodine gains one to form an I^- anion. They aim to have an electron configuration akin to the noble gases nearest to them.

• Nonmetals: Often gain electrons to achieve full valence shells.
• Examples: Oxygen (O) forming O^2-, Iodine (I) forming I^-.

Electron Loss

Electron loss is the process through which an atom gives up electrons to form positively charged ions, known as cations. This is a common trait seen in metals, especially those in Group 1 and Group 2 of the periodic table.
In the case of Rubidium (Rb) and Barium (Ba), they lose electrons to form ions. Rubidium loses one electron becoming Rb⁺, while Barium gives up two electrons to form Ba²⁺.
This results in a stable electron configuration similar to the noble gases.

• Metals: Tend to give up electrons due to low ionization energy.
• Examples: Magnesium (Mg) forming Mg²⁺ after losing two electrons.