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In both the compounds the oxidation state of Chromium is +6 so why is there a difference in the colours of their aqueous solutions.

$\ce{K2Cr2O7}$ is red-yellow and $\ce{K2CrO4}$ is bright yellow

I and leaning towards the fact that due to the increased amount of bond pairs in $\ce{K2Cr2O7}$ it becomes harder to remove the d orbital electron and that causes the light emitted to be of a higher wavelength(Lower frequency) and that's why there is colour difference.

NOTE: Even though both dichromate and chromate have the same oxidation state for Cr (i.e. +6) they show different colours. My question is how and why?

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  • $\begingroup$ I do know that the origin of color resides in the electron shells. The color of the photon that's emitted when an electron falls from its excited state back to its base state depends on the location within the electron shells of those two states, so it sounds like you're on the right track. $\endgroup$ – Cleve Weller May 11 '15 at 5:18
  • $\begingroup$ Have you made a Lewis dot diagram of the 2 molecules? I think you will see a difference in how the molecules are arranged. $\endgroup$ – LDC3 May 11 '15 at 5:53
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    $\begingroup$ I suggest you look into the phenomenon of Ligand to metal charge transfer. That'll give you your answer. $\endgroup$ – getafix May 11 '15 at 6:09
  • $\begingroup$ A sweeping observation is that the transition metals compounds have orbitals which have electron transitions that are in the visible range of the electromagnetic spectrum so they are colored. Most other inorganic compounds don't, and they are colorless. So color is all about bond orbitals not just the oxidation the state of the metal. $\endgroup$ – MaxW Oct 27 '15 at 18:33
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I'm really excited for this because I get to reference the almighty color wheel!! Fair warning, this answer is much more qualitative than quantitative, but that's more interesting sometimes anyways.

If you look at the structure of the chromate and dichromate ions next to each other (see here for structures: https://en.wikipedia.org/wiki/Chromate_and_dichromate), the only major difference between the two is that the Cr-O bond joining the two chromate ions (missing an oxygen) is now a single bond. This means that bond will vibrate at a lower frequency, and because frequency and wavelength are inversely related, that bond will absorb a longer wavelength of light.

Now to the color wheel! It is a general chemistry (often unexplained) fact that the color we see is the complementary color of the wavelength of a bond's vibration. Thus, in the case of the chromate ion, we see yellow, and across from yellow is the purple-ish region. That means, if one of the bonds in the chromate ion, and thus two of the bonds in the dichromate ion, were absorbing a longer wavelength like we said earlier, on average we would expect something just longer than purple-ish, like blue, to be absorbed.

The complementary color of blue is red slash orange, and that is in fact the color we see in the dichromate ion!

At the heart of all this is the principle that the colors we see are those wavelengths of light which on average are not absorbed by a large number (on the order of Avogadro's number) of molecules.

An approach like this will only be reliable for very similar molecules like the two we have here.

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  • $\begingroup$ I rather disagree with calling the bonds double bonds but otherwise a nice answer. (Making the bond longer = weaker works fine imho.) $\endgroup$ – Jan Oct 27 '15 at 18:04
  • $\begingroup$ I have to sort of suspend my beliefs a little to agree with your answers because bonds constitute colors to a compound when there exists a long chain conjugation. I have never read of bonds themselves contributing color. Are you sure that this color difference is not due to their lattice formation differences? $\endgroup$ – Agyey Arya Jan 30 '16 at 13:00
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    $\begingroup$ Well it can't be due to lattice structure as these are aqueous solutions. And sure single bonds can contribute color to molecules if they simply vibrate at the frequency of visible light. As an example how else can we explain colored gases like NO2? $\endgroup$ – jheindel Jan 30 '16 at 19:00
  • $\begingroup$ I really would cast doubt on your answer. To my knowledge, the vibrations of interatomic bonds give rise to absorption of infrared red radiation, not EM radiation in the UV/visible region. Your explanation is hence likely incorrect. $\endgroup$ – Tan Yong Boon Apr 8 at 13:20
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I will try to take a different approach to your question (in terms of acidity and basicity).

Chromate (yellow) and dichromate (orange) ion are at equilibrium in solution. The equilibrium equation can be represented by $$\ce{Cr2O7^2- + H2O <=> 2CrO4- + 2H+}$$

By Le Chatelier’s Principle, if certain conditions (concentration, temperature, pressure, volume, etc.) are changed, the amount of each ion present in solution is affected. If a shift in pH causes the solution to become more acidic (i.e. add $\ce{HCl}$), the equilibrium will shift to the left, and more dichromate ion will be present. The reaction tries to offset the increase in $\ce{H+}$ concentration and shifts to the left accordingly.

On the other hand, if a shift in pH causes the solution to be more basic (i.e. add $\ce{NaOH}$, which will cause some $\ce{H+}$ to react with $\ce{OH-}$ and produce water), the equilibrium will shift to the right, and more chromate ion will be present. This is a means by which the reaction tries to offset the loss of $\ce{H+}$.

The color of these ions is pH dependent, as indicated by the color changes when the above reactions take place.

Hope this helps a little.

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    $\begingroup$ While this is absolutely correct, I believe the question is why these compounds have different colours on a molecular basis and not why they change colour. $\endgroup$ – Martin - マーチン May 11 '15 at 13:06
  • $\begingroup$ My apologies, I thought OP was asking about the difference in colors of their aqueous solutions and thought of a different line of reasoning. Following this post now, because I would like to learn more... $\endgroup$ – imaginov May 11 '15 at 13:12
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The real answer after a very long research is that the change in colour is due to the charge transfer spectrum or specifically Ligand to Metal Charge Transfer Spectrum, read about it
here
https://en.wikipedia.org/wiki/Charge-transfer_complex
and here http://chemwiki.ucdavis.edu/Core/Physical_Chemistry/Spectroscopy/Electronic_Spectroscopy/Selection_Rules_for_Electronic_Spectra_of_Transition_Metal_Complexes/Metal_to_Ligand_and_Ligand_to_Metal_Charge_Transfer_Bands

Charge transfer spectrum says that there is a partial charge transfer between the ligand and the metal. This difference in charge changes the oxidation state momentarily and gives a different color than expected.

Due to the difference in bondings in both the given complexes there occurs a different amount of charge transfer leading to different colors despite of the same initial oxidation state.

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    $\begingroup$ Whilst this may theoretically answer the question, it would be preferable to include the essential parts of the answer here, and provide the link for reference. $\endgroup$ – bon Feb 10 '16 at 15:30
  • $\begingroup$ @bon I was assuming a previous knowledge of the matter involved and provided the links as way to refresh your ideas on the matter. $\endgroup$ – Agyey Arya Feb 10 '16 at 15:32
  • $\begingroup$ Actually this only explains why chromate and dichromate have a colour at all. To explain the difference between the two colours, you would have to explain why either chromium’s orbitals lost energy or oxygen’s gained it to reduce the energy difference between the two. Cf my answer on permanganate for the general principle. $\endgroup$ – Jan Feb 10 '16 at 20:24
  • $\begingroup$ I suggest you use the links provided to answer your questions, but in short the answer is that due to difference in bonding in both the complexes there exists a different amount of charge transfer leading to different OS in both the complexes. $\endgroup$ – Agyey Arya Feb 12 '16 at 2:29

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