1:32 PM
@hBy2Py Ah, ignorance is bliss ;D
Now @Koolman, I've had a look at the question again (and referred some material) and I've arrived at @hBy2Py's conclusion as well: The question is incomplete...that, or it's very poorly worded. Now you'll have to make a few assumptions before you can start figuring out the answer...
Now, I'm using (what I understand of) the Crystal Field Theory
You'd probably already know how the color of (transition metal) coordination compounds are determined using the above-said theory.
Greater the crystal field splitting brought about by the ligands, greater the energy required to bump up electrons from lower d-orbitals (for example: t2g in octahedral complexes) to higher d-orbitals (eg in octahedral complexes). Which would mean the more energetic (lower wavelength) end of the electromagnetic spectrum would end up getting absorbed.
So, since wavelengths in one portion of the spectrum are absorbed, the rest pass on and give the complex the color we see...
Say, if the crystal field splitting were very large (as would be in the case of CN- ligands) light from the blue-violet region would be absorbed, and the complex would appear yellow.
Similarly, if the crystal field splitting were very small (as would be the case, when you're dealing with F- ligands) light from the red-magenta region would be absorbed, and the complex would appear green.
Now, assuming that it's just a part of the visible spectrum that gets absorbed, we'll move on to your options.
Your first option's wrong, since the color change (roughly) ought to be from bluish-green to a green-yellow; since Cl- is a much weaker ligand that H2O.
Option two seems about right, since a strong field ligand is replaced by a weaker one.