What characterizes the electron configurations of transition metals such as silver and iron?
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Fe is in the 3d transition series and Ag is in the 4d series. Thus, the distinguishing electron for the 3d series enters the 3d orbital while those for the 4d series enters the 4d orbital.
To understand the electron configurations of transition metals like silver (Ag) and iron (Fe), we need to consider the building up (Aufbau) principle and the periodic table.
The electron configuration of an atom describes the arrangement of electrons in its orbitals. It follows a specific ordering based on the Aufbau principle, which states that electrons fill the lowest energy orbitals first before moving to higher energy levels.
For transition metals, the atoms have partially filled d orbitals. The d orbitals come after the s and p orbitals in the order of increasing energy. The transition metals have valence electrons that fill the d orbitals.
Let's take a look at the electron configurations for silver (Ag) and iron (Fe):
Silver (Ag):
The atomic number of silver is 47. To determine the electron configuration, we start by filling the orbitals in ascending order of energy, following the Aufbau principle.
The electron configuration for silver is: 1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d¹⁰ 4p⁶ 5s² 4d⁹
Here, we can see that the valence electrons for silver are in the 5s and 4d orbitals. The 5s² electrons are considered the outermost shell.
Iron (Fe):
The atomic number of iron is 26. Following the same process, we determine the electron configuration for iron.
The electron configuration for iron is: 1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d⁶
For iron, the valence electrons are in the 4s and 3d orbitals. The 4s² electrons are considered the outermost shell.
In summary, the electron configurations of transition metals like silver and iron are characterized by the filling of d orbitals. The valence electrons occupy both the s and d orbitals, with the d orbitals being partially filled.