Question

Write electron configurations for each of the following. a. \(\mathrm{Ti}, \mathrm{Ti}^{2+}, \mathrm{Ti}^{4+}\) b. \(\mathrm{Re}, \mathrm{Re}^{2+}, \mathrm{Re}^{3+}\) c. Ir, \(I r^{2+}, I r^{3+}\)

Short answer

a. Ti: \(1s^2 2s^2 2p^6 3s^2 3p^6 4s^2 3d^2\), \(\mathrm{Ti}^{2+}\): \(1s^2 2s^2 2p^6 3s^2 3p^6 3d^2\), \(\mathrm{Ti}^{4+}\): \(1s^2 2s^2 2p^6 3s^2 3p^6\) b. Re: \(1s^2 2s^2 2p^6 3s^2 3p^6 4s^2 3d^10 4p^6 5s^2 4d^10 5p^6 6s^2 4f^14 5d^5\), \(\mathrm{Re}^{2+}\): \(1s^2 2s^2 2p^6 3s^2 3p^6 4s^2 3d^10 4p^6 5s^2 4d^10 5p^6 4f^14 5d^5\), \(\mathrm{Re}^{3+}\): \(1s^2 2s^2 2p^6 3s^2 3p^6 4s^2 3d^10 4p^6 5s^2 4d^10 5p^6 4f^14 5d^4\) c. Ir: \(1s^2 2s^2 2p^6 3s^2 3p^6 4s^2 3d^10 4p^6 5s^2 4d^10 5p^6 6s^2 4f^14 5d^7\), \(I r^{2+}\): \(1s^2 2s^2 2p^6 3s^2 3p^6 4s^2 3d^10 4p^6 5s^2 4d^10 5p^6 4f^14 5d^7\), \(I r^{3+}\): \(1s^2 2s^2 2p^6 3s^2 3p^6 4s^2 3d^10 4p^6 5s^2 4d^10 5p^6 4f^14 5d^6\)

Step-by-step solution

Consult the periodic table to determine the atomic number of each element, which gives us the number of electrons in an atom. Titanium (Ti) has an atomic number of 22, Rhenium (Re) has an atomic number of 75, and Iridium (Ir) has an atomic number of 77. #Step 2: Write down the electron configurations for neutral atoms#

For each of the neutral atoms (Ti, Re, and Ir), write down the electron configurations following the Aufbau principle, which states that electrons fill atomic orbitals in the order of their increasing energy. Ti: \[1s^2 2s^2 2p^6 3s^2 3p^6 4s^2 3d^2 \] Re: \[1s^2 2s^2 2p^6 3s^2 3p^6 4s^2 3d^10 4p^6 5s^2 4d^10 5p^6 6s^2 4f^14 5d^5 \] Ir: \[1s^2 2s^2 2p^6 3s^2 3p^6 4s^2 3d^10 4p^6 5s^2 4d^10 5p^6 6s^2 4f^14 5d^7 \] #Step 3: Write electron configurations for the ions#

When ions are formed, it is the outermost electrons which are either added or removed. For Ti: Remove 2e for \(\mathrm{Ti}^{2+}\) and 4e for \(\mathrm{Ti}^{4+}\) For Re: Remove 2e for \(\mathrm{Re}^{2+}\) and 3e for \(\mathrm{Re}^{3+}\) For Ir: Remove 2e for \(I r^{2+}\) and 3e for \(I r^{3+}\) #Step 4: Present the final electron configurations for the ions#

Write down the final electron configurations for each ion: \(\mathrm{Ti}^{2+}\): \[1s^2 2s^2 2p^6 3s^2 3p^6 3d^2 \] \(\mathrm{Ti}^{4+}\): \[1s^2 2s^2 2p^6 3s^2 3p^6 \] \(\mathrm{Re}^{2+}\): \[1s^2 2s^2 2p^6 3s^2 3p^6 4s^2 3d^10 4p^6 5s^2 4d^10 5p^6 4f^14 5d^5 \] \(\mathrm{Re}^{3+}\): \[1s^2 2s^2 2p^6 3s^2 3p^6 4s^2 3d^10 4p^6 5s^2 4d^10 5p^6 4f^14 5d^4 \] \(I r^{2+}\): \[1s^2 2s^2 2p^6 3s^2 3p^6 4s^2 3d^10 4p^6 5s^2 4d^10 5p^6 4f^14 5d^7 \] \(I r^{3+}\): \[1s^2 2s^2 2p^6 3s^2 3p^6 4s^2 3d^10 4p^6 5s^2 4d^10 5p^6 4f^14 5d^6 \]

Key concepts

The Aufbau Principle

Understanding the Aufbau principle is crucial for students learning to write electron configurations. 'Aufbau' is a German word that means 'building up,' and the principle guides us on how electrons populate the atomic orbitals of an element.

Imagine each electron as a tenant looking for an 'energy home.' The principle suggests that electrons will fill the lowest energy orbitals available before moving to higher ones. This creates a predictable pattern that begins with the 1s orbital, progresses through the 2s and 2p, and continues in this fashion. To accurately determine the configuration of an atom, one should always begin with the Aufbau principle.

Let's apply this concept to a common exam question: 'Write the electron configuration for a Titanium (Ti) atom.' Consulting the periodic table, we see that Titanium has 22 electrons. According to the Aufbau principle, these electrons will fill up the orbitals in this order: 1s, 2s, 2p, 3s, 3p, and so on, until all electrons are accounted for, leading to the electron configuration for Titanium.

Periodic Table as a Guiding Map

The periodic table isn't just a display of elements; it's a map that reveals the electron structure of atoms. By understanding the table's layout, students can grasp the properties of elements and their ions.

Each row, known as a period, indicates a new energy level for electrons. For instance, elements in the second row (period) have their outermost electrons in the 2s or 2p subshells. Additionally, the columns, or groups, classify elements with the same number of electrons in their outer shell, which imparts similar chemical characteristics.

When the exercise at hand asks for transition metals like Titanium (Ti) or precious metals like Iridium (Ir), the periodic table shows us they belong to the d-block. This influences the order in which their d-orbitals are filled, and also guides us when we are removing electrons to form ions. Always return to the periodic table when deducing electron configurations, as it is the cornerstone of understanding elemental behavior.

Atomic Orbitals and Electron Housing

Atomic orbitals can be thought of as the rooms in an electron's 'house.' These are not just spaces but are defined by complex shapes and energy levels. Each orbital can hold a maximum of two electrons, which must have opposite spins due to the Pauli exclusion principle.

To explain further, let's take a look at the d-orbitals that hold up to ten electrons, making them pertinent in the discussion of transition metals. In our example with Titanium (Ti), after the first 18 electrons fill the 1s to 4s orbitals, the remaining four go into the 3d orbital. Understanding the shape and capacity of these orbitals greatly assists students in visualizing electron configurations and is a fundamental concept in chemistry and physics.

Ion Electron Configuration and Its Variations

Ion electron configurations differ from those of their neutral atoms. When atoms gain or lose electrons, they become ions and their electron configurations change. For positive ions (cations), electrons are removed starting from the orbital with the highest energy level down to the lowest.

For example, when a student is asked to write the electron configuration for \(\mathrm{Ti}^{4+}\), they must recognize that Titanium has lost four of its electrons. These are subtracted from the orbital with the highest energy level, which, for a neutral Titanium atom, is the 4s subshell, followed by the 3d subshell if there are still electrons to remove. The final electron configuration for \(\mathrm{Ti}^{4+}\) is thus \[1s^2 2s^2 2p^6 3s^2 3p^6\]. The electron removal process for cations can alter chemical properties, and students must be comfortable predicting these changes to excel in chemistry.