Question

What is the number of half-filled orbitals in an atom of chromium?

Answer 1

Hint: The chemical element chromium has the symbol Cr and the atomic number 24. It's the first element in the sixth group. It's a steely-grey transition metal that's glossy, hard, and brittle. In stainless steel, chromium is the most common addition, which gives it anti-corrosive characteristics. Chromium is also prized for its ability to be finely polished while remaining tarnishing-resistant. Polished chromium reflects over 70% of visible spectrum light and almost 90% of infrared light.

Complete answer:
The distribution of electrons in an element's atomic orbitals is described by its electron configuration. Atomic electron configurations follow a standard nomenclature in which all electron-containing atomic subshells are arranged in a sequence (with the number of electrons they possess indicated in superscript).
The electronic configuration of chromium must be written. Let's start with chromium's atomic number. Chromium has an atomic number of 24. The first two electrons in Chromium's electron configuration will be in the 1s orbital. Because the 1s orbital can only store two electrons, Chromium's next two electrons are assigned to the 2s orbital. The six electrons after that will be in the 2p orbital. Up to six electrons can be held in the p orbital. In the 3s, the next two electrons. The next six electrons will be added later. Now we'll move the remaining two electrons to the 4s orbital. The electrons that remain migrate into the 3d orbital.
Electronic configuration to be expected
As a result, the electron configuration for Chromium is predicted to be \[1{{s}^{2}}2{{s}^{2}}2{{p}^{6}}3{{s}^{2}}3{{p}^{4}}4{{s}^{2}}3{{d}^{9}}\].
Electronic configuration in action
Subshells that are half-filled or fully filled have more stability. As a result, one of the 4s2 electrons goes to the \[3{{d}^{5}}\], filling it halfway. This diagram depicts the (proper) chromium atom configuration: \[1{{s}^{2}}2{{s}^{2}}2{{p}^{6}}3{{s}^{2}}3{{p}^{4}}3{{d}^{5}}4{{s}^{1}}\]
Remember that chromium has a half-filled 3d subshell, much like molybdenum (but not tungsten): \[[Ar]3{{d}^{5}}4{{s}^{1}}\]
The short version is that increasing parallel electron spins reduces the energy of the electron configuration, therefore it's not \[3{{d}^{4}}4{{s}^{2}}\]. More information about that may be found by reading this answer here.
All of the core orbitals are occupied twice. As a result, they are excluded from the list of half-filled orbitals.
You should now notice that there are six unpaired electrons, indicating that there are six half-filled orbitals.

Note:
When chromium metal is exposed to air, it passivates, forming a thin, protective coating of oxide on the surface. This layer has a spinel structure that is a few atomic layers thick, dense, and prevents oxygen from diffusing into the underlying metal. Iron, on the other hand, creates a porous oxide that allows oxygen to pass through, resulting in continuous rusting. Short contact with oxidising acids like nitric acid can improve passivation.