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# Colour in Coordination compounds- d-d transition, Charge Transfer Spectra, Practice Problems and FAQs

When I finished high school, I insisted on my father purchasing a Ruby ring for me. I wanted a Ruby ring since I was a child because the bright red colour of Ruby has always captivated me. My father's close friend, a Goldsmith congratulated me on my academic achievement when we visited his jewellery store. He posed me a question, “you are so intelligent, and so can you figure out why your loving Ruby has such a bright red colour?”

At that time, I was not knowing. Do You know what the answer is?

See, Ruby is composed of octahedral chromium(III) complexes incorporated into the alumina lattice; d–d transitions at these centres produce the red colour. Do you understand anything? You will at the end. Go ahead and find out.

Let us go over these d-d transitions in greater detail.

• Colour in Coordination compounds
• d-d transition
• Charge transfer spectra
• Practice Problems

## Colour in Coordination compounds

The Coordination compounds are complexes containing mostly a transition metal atom/ion attached by a coordinate bond to ligands.

The majority of transition metal complexes are coloured in solid or solution form.

As we all know, visible light (white light) is a complex mixture of radiations with wavelengths ranging from blue (about 400 nm) to red (about 700 nm). Transition metal ions have the ability to absorb certain radiations from the visible region of the spectrum, resulting in coloured transmitted or reflected light.

When some of these wavelengths from a white light beam are absorbed after passing through a sample, the transmitted light is no longer white.

The resulting colour of the complex can be explained by either

1. d-d transition or
2. Charge transfer spectra

According to CFT, approaching ligands removes the degeneracy of the five d-orbitals of the transition metal and rearrange the electrons present in them. Two forces come into the act- one is the energy required to force the pairing of electrons in the same d-orbital and the energy difference between the two sets of orbitals now produced.

## d-d transition

Let. the energy difference between the two sets of d-orbitals in transition metal complexes is minimal. When exposed to visible light, an electron from a lower energy d-orbital may get excited to a higher energy d-orbital.

Example: In the case of octahedral complexes,

d-d transition

Due to the absorption of light at a specific wavelength in the visible region, there is transmission or reflection of the light with the rest of the wavelengths. The complementary colour is generated from the wavelength left over.

For example, if red light (long wavelength) is absorbed from the white light, it appears green to us (short wavelength). When green light (with a short wavelength) is absorbed, it appears red to us.

Colours like red and green are known as complementary colours. Thus, the colour of a complex in a solid or solution is caused by light that is not absorbed but is transmitted (complementary colour).

A metal complex must contain an unpaired electron for a d-d transition to occur.

## Charge transfer spectra

The transition in which an electron is transferred from one atom or group to another one, this is known as charge transfer spectra. Charge transfer transitions are also considered as internal redox reactions.

Charge transfer transitions are mainly of three types

1. Ligand to metal charge transfer (LMCT)
2. Metal to metal charge transfer (MMCT)
3. Metal to ligand charge transfer (MLCT)

### 1. Ligand to metal charge transfer (LMCT)

Charge transfer occurs from the ligand's filled molecular orbitals to the metal’s empty or

partially filled d-orbitals. Generally, metals having a high oxidation state (+5,+6,+7) with

donor ligands show LMCT.

### Metal to metal charge transfer (MMCT)

The excitation and subsequent transfer of electrons from a lower oxidation state (O.S) cation to neighbouring cation of a higher O.S shows MMCT.

Example

Example

### Metal to Ligand charge transfer (MLCT)

When charge transfer occurs from the metal’s filled orbitals to the empty orbital of ligand, it is known as MLCT. Generally, metals having a low oxidation state (0,+1,+2) with acceptor ligands

show LMCT.

Example:

## Practice Problems

Q1. Which of the following complex ions should not absorb visible light?

• [Co(CN)6]4-
• [Fe(CN)6]3-
• [Mn(H2O)6]2+
• [Mn(CN)6]4-

Solution: In option A, B and D the configuration of Metal ion is Co2+ (d7) (low spin), Fe3+ (d5) (low spin) and Mn2+ (d4) (low spin) is there respectively. So, there is a possibility of having d-d transition as unpaired electrons are present in each case.

Complexes with Mn (d5 configuration) are centrosymmetric (have a centre of symmetry), and thus d-d transition is not permitted. The colour of the complex due to the d-d transition is thus absent in Mn. As a result, they are nearly colourless.

Hence, the correct answer is an option (C).

Q2. [Ti(H2O)6]3+ absorbs light of wavelength 498 nm during a d-d transition. The octahedral splitting energy for the above complex is _____ × 1019 J (round off to nearest integer)

h = 6.626 × 10-34Js;c = 3 × 108m/s

• × 10-8J
• 4 × 10-18J
• 4 × 10-15J
• 4 × 10-17J

Solution: Given: Wavelength = 498 nm

h = 6.626 × 10-34Js;c = 3 × 108m/s

The octahedral splitting energy =

So, the correct answer is an option (D).

Q3. The correct order of the wavelength of absorption of the visible region is

• [Cr(CN)6]3+ < [Cr(H2O)6]3+ < [Cr(NH3)6]3+
• [Cr(H2O)6]3+ < [Cr(NH3)6]3+ <[Cr(CN)6]3+
• [Cr(CN)6]3+ < [Cr(NH3)6]3+ < [Cr(H2O)6]3+
• [Cr(H2O)6]3+ <[Cr(CN)6]3+ < [Cr(NH3)6]3+

Solution: The strength of ligands as per spectrochemical series is H2O < NH< CN-. As more is the strength of ligand, more is the Crystal field splitting energy, and lesser is the wavelength of absorption. So, the correct order of the wavelength of absorption should be

[Cr(CN)6]3+ < [Cr(NH3)6]3+ < [Cr(H2O)6]3+.

So, the correct answer is an option (C).

Q4. The colour of KMnO4 is due to

• LMCT (ligand to Metal charge transfer)
• MLCT (Metal to ligand charge transfer)
• MMCT (Metal to Metal charge transfer)
• d-d transition

Solution: The metal ions in KMnO4 contain d electrons, charge transfer occurs from O2- to Mn2+. The transition of a nonbonding 2p oxygen electron to the unoccupied molecular orbital level of the doing tetrahedral compound results in the lowest energy LM charge transfer. Hence, the reason for colour is LMCT (ligand to metal charge transfer). So, the correct answer is an option (A).

Question 1. Why do coordination compounds appear in different colours?
Because the difference between different sets of orbitals in octahedral complexes varies depending on the metal ion and the nature of the ligands, different complexes absorb different amounts of energy from the visible region and exhibit different colours.

Question 2. Why is the famous diamond Emerald green in colour?
The presence of transition metal ions causes the colours of many well-known gemstones, such as Emerald. Colours are produced in these by electronic transitions within d-orbitals of a transition metal ion (called d-d transitions).

Question 3. Why are anhydrous copper sulphate (CuSO4)white and pentahydrate copper sulphate (CuSO4.5H2O)blue?
Answer. The water molecules surrounding the Central Metal (Cu) act as ligands in hydrated CuSO4, resulting in a d-d transition and thus emitting blue colour in the visible region, making hydr ted CuSO4 appear blue. Because anhydrousCuSO4 contains no water during crystallization, it retains its white colour.

Question 4. What effect do ligands have on the colour of transition metal ions?
Answer. When the ligands bond with the transition metal ion, the electrons in the ligands repel the electrons in the metal ion's d orbitals. This increases the energy of the d orbitals and causes d orbital splitting. If no ligand is present, d orbitals are degenerate, and the transition metal ion becomes colourless.

Related Topics

 Oxidation number of elements in coordination compounds Organometallic Compounds EAN Rule Ligands Bonding in coordination compounds

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