Problem 17 - Chemistry: The Central Science

Question

Write the electron configuration for each of the following ions, and determine which ones possess noble-gas configurations: (a) \(\mathrm{Sr}^{2+}\), (b) \(\mathrm{Ti}^{2+}\), (c) \(\mathrm{Se}^{2-}\), (d) \(\mathrm{Ni}^{2+}\), (e) \(\mathrm{Br}^{-}\), (f) \(\mathrm{Mn}^{3+}\).

Step-by-Step Solution

Answer

The electron configurations for the given ions are: (a) \(\mathrm{Sr}^{2+}\): \(1s^2 2s^2 2p^6 3s^2 3p^6 4s^2 3d^{10} 4p^6\) - Noble-gas configuration (b) \(\mathrm{Ti}^{2+}\): \(1s^2 2s^2 2p^6 3s^2 3p^6 3d^2\) - Not a noble-gas configuration (c) \(\mathrm{Se}^{2-}\): \(1s^2 2s^2 2p^6 3s^2 3p^6 4s^2 3d^{10} 4p^6\) - Noble-gas configuration (d) \(\mathrm{Ni}^{2+}\): \(1s^2 2s^2 2p^6 3s^2 3p^6 3d^8\) - Not a noble-gas configuration (e) \(\mathrm{Br}^{-}\): \(1s^2 2s^2 2p^6 3s^2 3p^6 4s^2 3d^{10} 4p^6\) - Noble-gas configuration (f) \(\mathrm{Mn}^{3+}\): \(1s^2 2s^2 2p^6 3s^2 3p^6 3d^5\) - Not a noble-gas configuration The ions \(\mathrm{Sr}^{2+}\), \(\mathrm{Se}^{2-}\), and \(\mathrm{Br}^{-}\) possess noble-gas configurations.

Step 1: Write electron configuration for neutral atoms

For each of the given elements, we can write their electron configurations based on their position in the periodic table and their atomic number. Remember to start with the element with the lowest atomic number and work upwards. (a) Strontium (Sr) - Atomic number 38: \(1s^2 2s^2 2p^6 3s^2 3p^6 4s^2 3d^{10} 4p^6 5s^2\) (b) Titanium (Ti) - Atomic number 22: \(1s^2 2s^2 2p^6 3s^2 3p^6 4s^2 3d^2\) (c) Selenium (Se) - Atomic number 34: \(1s^2 2s^2 2p^6 3s^2 3p^6 4s^2 3d^{10} 4p^4\) (d) Nickel (Ni) - Atomic number 28: \(1s^2 2s^2 2p^6 3s^2 3p^6 4s^2 3d^8\) (e) Bromine (Br) - Atomic number 35: \(1s^2 2s^2 2p^6 3s^2 3p^6 4s^2 3d^{10} 4p^5\) (f) Manganese (Mn) - Atomic number 25: \(1s^2 2s^2 2p^6 3s^2 3p^6 4s^2 3d^5\)

Step 2: Modify electron configuration for ions

Now, let's modify the electron configurations for each ion based on their charge. Adding electrons increases the electron configuration (for negatively charged ions), and removing electrons decreases the electron configuration (for positively charged ions). (a) \(\mathrm{Sr}^{2+}\): Remove 2 electrons: \(1s^2 2s^2 2p^6 3s^2 3p^6 4s^2 3d^{10} 4p^6\) (b) \(\mathrm{Ti}^{2+}\): Remove 2 electrons: \(1s^2 2s^2 2p^6 3s^2 3p^6 3d^2\) (c) \(\mathrm{Se}^{2-}\): Add 2 electrons: \(1s^2 2s^2 2p^6 3s^2 3p^6 4s^2 3d^{10} 4p^6\) (d) \(\mathrm{Ni}^{2+}\): Remove 2 electrons: \(1s^2 2s^2 2p^6 3s^2 3p^6 3d^8\) (e) \(\mathrm{Br}^{-}\): Add 1 electron: \(1s^2 2s^2 2p^6 3s^2 3p^6 4s^2 3d^{10} 4p^6\) (f) \(\mathrm{Mn}^{3+}\): Remove 3 electrons: \(1s^2 2s^2 2p^6 3s^2 3p^6 3d^5\)

Step 3: Identify noble-gas configurations

A noble-gas configuration is when an ion has a full outer electron shell, matching the electron configuration of a noble gas. Comparing the ions' electron configurations to those of noble gases, we can determine which ions possess noble-gas configurations. (a) \(\mathrm{Sr}^{2+}\) has the electron configuration of Kr (krypton): Noble-gas configuration (b) \(\mathrm{Ti}^{2+}\) does not match any noble gas: Not a noble-gas configuration (c) \(\mathrm{Se}^{2-}\) has the electron configuration of Kr (krypton): Noble-gas configuration (d) \(\mathrm{Ni}^{2+}\) does not match any noble gas: Not a noble-gas configuration (e) \(\mathrm{Br}^{-}\) has the electron configuration of Kr (krypton): Noble-gas configuration (f) \(\mathrm{Mn}^{3+}\) does not match any noble gas: Not a noble-gas configuration Thus, the ions \(\mathrm{Sr}^{2+}\), \(\mathrm{Se}^{2-}\), and \(\mathrm{Br}^{-}\) possess noble-gas configurations.

Key Concepts

Noble-Gas Configurations

In chemistry, noble-gas configurations are a core concept because they represent a stable state for atoms and ions. Noble gases, such as helium (He), neon (Ne), and argon (Ar), possess completely filled outer electron shells. This full electron shell configuration contributes to their stability and inertness, as they rarely participate in chemical reactions.

When an ion has the same electron configuration as a noble gas, it is said to have a noble-gas configuration or is "isoelectronic" with that noble gas. For example:

• Strontium ion \( \(\mathrm{Sr}^{2+}\), \) achieves the configuration of krypton (Kr) by losing two electrons.
• Selenium ion \(\mathrm{Se}^{2-}\) gains two electrons to reach the configuration of krypton as well.

Many elements naturally tend toward forming ions that resemble noble gases. This tendency explains why some ions, like \(\mathrm{Br}^{-}\), gain or lose electrons to achieve this "ideal stability".

Ion Electron Configurations

Ion electron configurations involve the adjustment of electron counts due to the gain or loss of electrons, resulting in a positive or negative ion. Let's break it down:

• A positive ion (cation), such as \(\mathrm{Sr}^{2+}\), loses electrons. Usually, electrons are removed starting from the highest principal quantum number, often the outermost shell.
• Conversely, a negative ion (anion), like \(\mathrm{Se}^{2-}\), adds electrons to its electron seashell.

To modify the electron configuration for an ion:- Determine the charge and then adjust electron numbers, - For cations, subtract the necessary number of electrons from the neutral atom's configuration,- For anions, add electrons to the neutral configuration.The goal is frequently to reach a noble-gas configuration, achieving greater stability. Understanding this concept is crucial for predicting the properties and reactivity of ions in chemical reactions.

Atomic Number

The atomic number of an element is fundamental to understanding its position on the periodic table and its electron configuration. The atomic number represents the number of protons in an atom's nucleus, and for a neutral atom, it also equals the number of electrons.

• This number dictates the electron subshell filling order: starting at \(1s\) and moving to \(4p\) or higher.
• Knowing the atomic number allows for determining the element's basic properties and helps compute their electron configuration.

For example:- Strontium (Sr), with an atomic number of 38, fills its electron orbitals up to \(5s^2\) in the neutral state (before losing electrons).- Titanium (Ti), atomic number 22, fills up through \(3d\) in its natural state before forming ions.

The concept of atomic numbers bridges the gap between the theoretical electron configuration and the practical structures observed across the periodic table.

Periodic Table

The periodic table serves as an essential tool for chemists, organizing elements by atomic number and recurring chemical properties. Its layout offers insights into electron configurations and the tendencies of elements to form ions:

• Groups or columns often share similar valence shell electron configurations, leading to shared chemical behaviors.
• The position of an element in a period or row relates to the outermost electron shell being filled.

Elements such as bromine (Br) and selenium (Se), close in the table, can exhibit related traits in gaining electrons. The periodic table guides users to predict ion charge and electronic structures easily. Each block or section (s, p, d, f) provides hints on which electron subshells are being filled, ultimately forming the basis for electron configuration determination. From identifying noble-gas configuration possibilities to anticipating ion formation trends, the periodic table is the chemist's roadmap.