Question

In which of the following is(are) the electron configuration(s) correct for the species indicated? a. Cu \([\mathrm{Ar}] 4 s^{2} 3 d^{9}\) b. \(\mathrm{Fe}^{3+} \quad[\mathrm{Ar}] 3 d^{5}\) c. Co \([\mathrm{Ar}] 4 s^{2} 3 d^{7}\) d. La \([\mathrm{Ar}] 6 s^{2} 4 f^{1}\) e. \(\mathrm{Pt}^{2+} \quad[\mathrm{Xe}] 4 f^{14} 5 d^{8}\)

Short answer

The correct electron configurations are found in options b (\(\mathrm{Fe}^{3+} \quad[\mathrm{Ar}] 3 d^{5}\)), c (Co \( [\mathrm{Ar}] 4 s^{2} 3 d^{7} \)), and e (\(\mathrm{Pt}^{2+} \quad[\mathrm{Xe}] 4 f^{14} 5 d^{8}\)).

Step-by-step solution

Identifying the atomic numbers

First, we need to identify the atomic numbers for each of the elements mentioned. Here are the atomic numbers for each element: - Cu (Copper): 29 - Fe (Iron): 26 - Co (Cobalt): 27 - La (Lanthanum): 57 - Pt (Platinum): 78 Now that we have the atomic numbers, we can determine the electronic configurations for each species.

Checking electron configurations

a. Cu \( [\mathrm{Ar}] 4 s^{2} 3 d^{9} \) Copper's atomic number is 29. Its correct electron configuration is \( [\mathrm{Ar}] 4 s^{1} 3 d^{10} \), since one electron from the 4s orbital is transferred to the 3d orbital as it provides extra stability. Therefore, this choice is incorrect. b. \(\mathrm{Fe}^{3+} \quad[\mathrm{Ar}] 3 d^{5} \) Iron's atomic number is 26. Its electron configuration is \( [\mathrm{Ar}] 4 s^{2} 3 d^{6} \). As Fe becomes Fe³⁺, it loses 3 electrons. The two 4s electrons are lost first, followed by one 3d electron, giving the final electron configuration \( [\mathrm{Ar}] 3 d^{5} \). Therefore, this choice is correct. c. Co \( [\mathrm{Ar}] 4 s^{2} 3 d^{7} \) Cobalt's atomic number is 27. Its correct electron configuration is \( [\mathrm{Ar}] 4 s^{2} 3 d^{7} \), the same as given. Therefore, this choice is correct. d. La \( [\mathrm{Ar}] 6 s^{2} 4 f^{1} \) Lanthanum's atomic number is 57. Its correct electron configuration is \( [\mathrm{Ar}] 5 d^{1} 6 s^{2} \). The given configuration is incorrect. e. \(\mathrm{Pt}^{2+} \quad[\mathrm{Xe}] 4 f^{14} 5 d^{8} \) Platinum's atomic number is 78. Its electron configuration is \( [\mathrm{Xe}] 4 f^{14} 5 d^{9} 6 s^{1} \). As Pt becomes Pt²⁺, it loses 2 electrons. The 6s electron is lost first, followed by one 5d electron, giving the final electron configuration \( [\mathrm{Xe}] 4 f^{14} 5 d^{8} \). Therefore, this choice is correct. Thus, the correct electron configurations are found in options b, c, and e.

Key concepts

Atomic Numbers

The atomic number of an element is a fundamental property that tells us the number of protons in an atom's nucleus. This number also defines the element's identity and position in the periodic table. For example, copper (Cu) has an atomic number of 29, which means every copper atom has 29 protons.

Knowing the atomic number is crucial because it helps determine the electron configuration of an atom, which is how electrons are distributed in an atom's electron shell. Because electrons and protons are equal in a neutral atom, the atomic number directly indicates the number of electrons present in the most stable form of an atom. If the atomic number changes, the identity of the element itself changes, highlighting its importance in chemical calculations and species identification.

Species Identification

Species identification in chemistry typically involves determining the identity of ions or neutral atoms within a compound. This can often be complicated by similar electron configurations or atomic numbers. However, electron configurations reveal subtle changes that indicate whether a species is an ion or a neutral atom.

For instance, when an electron configuration is analyzed during a question such as the one in the exercise, it indicates the species being discussed. In the step-by-step solution, when Fe is converted to Fe³⁺, three electrons are removed, which must be precisely accounted for to identify the species correctly. This process involves understanding the order of electron removal and how it affects the overall charge and identity of the species. Accurately interpreting these configurations ensures you are identifying and understanding the species in question properly, which is key in many chemical reactions and analyses.

Electron Stability

Electron stability is a concept that explains why electrons occupy certain energy levels and orbitals in a particular way. Electrons tend to distribute themselves to achieve the lowest energy state, which often means completely filling or half-filling their orbitals to gain extra stability.

A classic example from the exercise is copper (Cu), where we expect to see the electron configuration \( [\mathrm{Ar}] 4s^{1} 3d^{10} \) instead of \( [\mathrm{Ar}] 4s^{2} 3d^{9} \). This is due to the slightly greater stability found in a filled 3d subshell. Changes like these are rooted in quantum mechanics, where half-filled and fully filled shells provide symmetrical configurations that are energetically favorable. Understanding these changes is essential, especially when predicting the behavior of elements during chemical bonding or reactions.

Electron Removal Order

The electron removal order is a principle that refers to the sequence in which electrons are lost from an atom to form ions. Typically, electrons are removed starting from the outermost shell, but sometimes the specific sequence can vary based on orbital energies.

In the case of transition metals, like Fe and Pt in the exercise, electrons are generally removed from the 4s orbital before the 3d orbital, despite the 4s orbital being filled first under normal circumstances. For example, when forming the \( \mathrm{Fe}^{3+} \) ion, two electrons are removed from the 4s orbital first, followed by one from the 3d orbital, resulting in the configuration \( [\mathrm{Ar}] 3d^{5} \).

• Electrons in s-orbitals are usually the first to go when forming cations.
• Electrons in higher energy levels will be removed before those in lower energy levels.

Understanding the electron removal order provides insights into the formation and behavior of ions, crucial in reaction mechanisms and redox reactions within chemistry.