Question

(b) Explain why a solution of CuCl, in aqueous HCl is deep greenish
blue whereas similar aqueous HCl solution of CuCl is colourless.
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Answer 1

Answer:

Explanation:

The reason is indeed d-to-d transitions. The d orbitals of both CuCl and CuCl2 have two possible energy levels (based on crystal field theory).

Cu(II) is a d9 metal, and CuCl2 exists as a (distorted) octahedral around the copper atom. Based on crystal field theory, the d orbitals of an octahedral metal center can be represented as shown below. In this case, because Cu(II) is d9, the three lower energy suborbitals are fully filled in the ground state, and but the top two suborbitals only contain a combined three electrons. Therefore, the energy required to excite that electron is Δo. When that electron returns to the ground state, it gives off a different wavelength of light than it absorbed (in this case, the light emitted is apparently not visible light).

(THE PICTURE BELOW WILL EXPLAIN IT)

Cu(I) , on the other hand, is a d10 metal ion, so each of the d suborbitals is fully filled. In order for an electron to be excited in this situation, it would have to leave the 3d orbital (and the n=3 shell altogether) and enter the n=4 shell. This is rather large energy jump compared to the intraorbital jump an electron in CuCl2 has to make. Therefore, CuCl does not readily absorb visible light as CuCl2 does, so it appears colorless. That is, until it see moist air.