Why Do Fe(II) Complexes Exhibit Unexpected Colors?

In summary: Those are the diagrams that scientists use to understand the energetics of d6 electron systems. In summary, the d-d transitions are dipole forbidden and therefore have only weak intensity. Mostly charge transfer transitions are more intense and dominate the color impressio of the complex.
  • #1
lee403
16
1
There is something that I am not understanding about the visible colors of coordination complexes. I did a lab where I prepared [Fe(NH2trz)3]Br2 from FeBr2.
Fe(NH2trz)3 was a violet color and FeBr2 was red. My understanding of colors in coordination complexes is that reds are high energy because they absorb the complementary color most strongly which would be purplish. Purple is high energy light because it has a shorter wavelength. So if FeBr2 is red that means there is a large d-d orbital splitting or a large energy difference between the t2g and eg levels of the d-d orbitals. However, Br is a weak field ligand (for another reason [poor metal-ligand overlap] that I don't quite understand) and NH2trz since its bonding through the N is an intermediate ligand. I thought that since weak field ligands have small d-d orbital splitting they would absorb low energy (long wavelength) light, which would be red and would therefore appear purple. But the complexes I have are the opposite of how I have interpreted color so far. What part of my reasoning is incorrect?
 
  • Like
Likes Pratyaya Ghoshal Das
Chemistry news on Phys.org
  • #2
lee403 said:
There is something that I am not understanding about the visible colors of coordination complexes. I did a lab where I prepared [Fe(NH2trz)3]Br2 from FeBr2.
Fe(NH2trz)3 was a violet color and FeBr2 was red. My understanding of colors in coordination complexes is that reds are high energy because they absorb the complementary color most strongly which would be purplish. Purple is high energy light because it has a shorter wavelength. So if FeBr2 is red that means there is a large d-d orbital splitting or a large energy difference between the t2g and eg levels of the d-d orbitals. However, Br is a weak field ligand (for another reason [poor metal-ligand overlap] that I don't quite understand) and NH2trz since its bonding through the N is an intermediate ligand. I thought that since weak field ligands have small d-d orbital splitting they would absorb low energy (long wavelength) light, which would be red and would therefore appear purple. But the complexes I have are the opposite of how I have interpreted color so far. What part of my reasoning is incorrect?
d-d transitions are dipole forbidden and therefore have only weak intensity. Mostly charge transfer transitions are more intense and dominate the colour impressio of the complex.
 
  • #3
lee403 said:
There is something that I am not understanding about the visible colors of coordination complexes. I did a lab where I prepared [Fe(NH2trz)3]Br2 from FeBr2.
Fe(NH2trz)3 was a violet color and FeBr2 was red. My understanding of colors in coordination complexes is that reds are high energy because they absorb the complementary color most strongly which would be purplish. Purple is high energy light because it has a shorter wavelength. So if FeBr2 is red that means there is a large d-d orbital splitting or a large energy difference between the t2g and eg levels of the d-d orbitals. However, Br is a weak field ligand (for another reason [poor metal-ligand overlap] that I don't quite understand) and NH2trz since its bonding through the N is an intermediate ligand. I thought that since weak field ligands have small d-d orbital splitting they would absorb low energy (long wavelength) light, which would be red and would therefore appear purple. But the complexes I have are the opposite of how I have interpreted color so far. What part of my reasoning is incorrect?
You might want to be careful with Fe2+ compounds (d6 electron) and crystal field. Ever heard of "Tanabe-Sugano diagram"?
 

1. What is the color of Fe(II) complexes?

The color of Fe(II) complexes can vary depending on the ligands attached to the iron ion. Generally, Fe(II) complexes are pale green or yellow in color, but can also appear blue, purple, or pink.

2. Why do Fe(II) complexes have different colors?

The color of Fe(II) complexes is determined by the arrangement of electrons in the d-orbitals of the iron ion. Different ligands can cause a change in the energy levels of these orbitals, leading to a different color for the complex.

3. How does the color of Fe(II) complexes change with pH?

The color of Fe(II) complexes can change with pH due to the protonation or deprotonation of ligands. This can lead to a change in the arrangement of electrons in the d-orbitals, resulting in a different color for the complex.

4. Can the color of Fe(II) complexes be used for analytical purposes?

Yes, the color of Fe(II) complexes can be used for analytical purposes such as determining the concentration of the complex in a solution. This is known as colorimetric analysis, where the intensity of the color is directly proportional to the concentration of the complex.

5. How can the color of Fe(II) complexes be altered?

The color of Fe(II) complexes can be altered by changing the ligands attached to the iron ion. Different ligands can cause a change in the energy levels of the d-orbitals, resulting in a different color for the complex. Temperature and pH can also affect the color of Fe(II) complexes.

Similar threads

  • Chemistry
Replies
1
Views
4K
Replies
2
Views
2K
Replies
25
Views
1K
  • Biology and Chemistry Homework Help
Replies
1
Views
3K
Replies
4
Views
2K
Replies
14
Views
3K
  • Sci-Fi Writing and World Building
Replies
15
Views
3K
  • Sci-Fi Writing and World Building
Replies
21
Views
961
  • Programming and Computer Science
Replies
1
Views
1K
  • Math Proof Training and Practice
6
Replies
175
Views
20K
Back
Top