Question

Oxyhemoglobin, with an ${\mathrm{O}}_2$ bound to iron, is a low-spin $\mathrm{Fe}$ (II) complex; deoxyhemoglobin, without the ${\mathrm{O}}_2$ molecule, is a high-spin complex. (a) Assuming that the coordination environment about the metal is octahedral, how many unpaired electrons are centered on the metal ion in each case? (b) What ligand is coordinated to the iron in place of ${\mathrm{O}}_2$ in deoxyhemoglobin? (c) Explain in a general way why the two forms of hemoglobin have different colors (hemoglobin is red, whereas deoxyhemoglobin has a bluish cast). (d) A 15 -minute exposure to air containing $400\mathrm{ppm}$ of CO causes about $10\mathrm{\%}$ of the hemoglobin in the blood to be converted into the carbon monoxide complex, called carboxyhemoglobin. What does this suggest about the relative equilibrium constants for binding of carbon monoxide and ${\mathrm{O}}_2$ to hemoglobin? (e) $\mathrm{CO}$ is a strong-field ligand. What color might you expect carboxyhemoglobin to be?

Short answer

(a) Oxyhemoglobin has 0 unpaired electrons (low-spin); deoxyhemoglobin has 4 unpaired electrons (high-spin). (b) An imidazole group from the histidine residue in hemoglobin is coordinated to the iron in deoxyhemoglobin. (c) The different colors are due to different electronic configurations and ligands, which change the energy differences between electronic transitions. (d) The larger equilibrium constant for CO binding shows it binds more strongly than O₂ to hemoglobin. (e) Carboxyhemoglobin has a cherry red color due to the strong field nature of CO on the electronic transitions in the complex.

Step-by-step solution

(a) Unpaired electrons of oxyhemoglobin and deoxyhemoglobin

To find the number of unpaired electrons on the metal ion (Fe2+) in oxyhemoglobin and deoxyhemoglobin, we need to consider the electron configurations corresponding to low-spin and high-spin complexes. For Fe2+, the electron configuration is $3d^6$. - In a low-spin complex, the electrons fill the lower-energy 3d orbitals before any high-energy 3d orbitals are occupied, making the complex diamagnetic with no unpaired electrons. - In a high-spin complex, the electrons are distributed across all the 3d orbitals before any pairing occurs, resulting in four unpaired electrons. Thus, oxyhemoglobin (low-spin) has 0 unpaired electrons, while deoxyhemoglobin (high-spin) has 4 unpaired electrons.

(b) Ligand in deoxyhemoglobin

In deoxyhemoglobin, the iron (Fe2+) is not coordinated to an O₂ molecule, but it still has 6 ligands maintaining its octahedral geometry. One of these ligands is an imidazole group from the histidine residue of hemoglobin (protein component). This histidine residue remains coordinated to the iron even when O₂ is not bound.

(c) Different colors of oxyhemoglobin and deoxyhemoglobin

The different colors of oxyhemoglobin (red) and deoxyhemoglobin (bluish) can be explained by the different electronic configurations and the ligands attached to the iron ion. The presence or absence of O₂-bound iron changes the strength of the ligand field and electronic transitions. The energy differences between these transitions determine the color absorbed, and therefore, the color we perceive.

(d) Relative equilibrium constants of CO and O₂ binding

The fact that a 15-minute exposure to 400 ppm CO causes about 10% of hemoglobin in the blood to convert into carboxyhemoglobin suggests that the equilibrium constant for binding of carbon monoxide (CO) to hemoglobin is much larger than the equilibrium constant for binding of O₂. This means that CO binds more strongly than O₂ to the iron ion in hemoglobin, making it difficult for O₂ to compete for the binding site.

(e) Color of carboxyhemoglobin

Since CO is a strong-field ligand, its binding to the iron in hemoglobin will result in a low-spin complex. Low-spin complexes typically exhibit color changes due to the increased ligand field strength and their associated charge transfer transitions. Carboxyhemoglobin is known to have a cherry red color, which is distinct from the red and bluish colors of oxyhemoglobin and deoxyhemoglobin, respectively. This change in color is due to the strong field nature of CO and its effect on the energy differences between the electronic transitions in the complex.

Key concepts

Low-Spin and High-Spin Complexes

Understanding low-spin and high-spin complexes is essential in chemistry, especially when examining transition metal ions like iron in hemoglobin. Iron can exist in two different types of complexes based on the arrangement of electrons in its d orbitals.

In a low-spin complex, such as oxyhemoglobin, electrons pair up in the lower energy orbitals first. This results in fewer unpaired electrons and a lower magnetic moment. Generally speaking, low-spin complexes like this arise from the presence of strong-field ligands, such as O₂, which results in a larger energy gap that electrons must overcome to reach higher orbitals, leading to a more stable and 'quieter' arrangement.

In contrast, a high-spin complex has more unpaired electrons because electrons will occupy all the d orbitals with parallel spins before pairing. This can be observed in deoxyhemoglobin, where weak-field ligands provide less stabilization, and the energy gap is smaller, allowing electrons to remain unpaired. These high-spin complexes are magnetically more active due to the greater number of unpaired electrons.

Ligand Coordination in Hemoglobin

Ligand coordination in hemoglobin plays a pivotal role in its function as an oxygen transporter. In hemoglobin, the iron ion is situated within a porphyrin ring and associates with several ligands, creating an octahedral arrangement.

Oxyhemoglobin is the form where oxygen acts as a ligand, while deoxyhemoglobin lacks the oxygen and instead maintains coordination through other groups. Notably, the nitrogen from an imidazole group of a histidine residue in the protein structure acts as a ligand. This maintains the octahedral coordination, ensuring that the iron ion remains in the correct position for oxygen binding when it becomes available again. This coordination is central to hemoglobin's ability to pick up and release oxygen molecules efficiently.

Color Differences in Hemoglobin Forms

The color differences in hemoglobin forms arise due to electronic transitions within the iron's d orbitals and the ligands coordinated to it.

Oxyhemoglobin appears red because the oxygen ligand creates a strong-field environment that influences the d-d electron transitions. Deoxyhemoglobin, on the other hand, appears to have a bluish tint, attributable to the weaker field of the histidine ligand and the resulting high-spin state with different electron transition energies.

These electronic transitions absorb light at different wavelengths, so the reflected or transmitted light gives off different colors. Think of it like a selective light filter, where each form of hemoglobin absorbs some colors and reflects others, resulting in varied apparent colors.

Hemoglobin Binding Affinity

The concept of hemoglobin binding affinity is crucial for its biological function. Hemoglobin's ability to bind oxygen and other gases involves a balance between affinity and release, so oxygen can be transported to where it is needed and released efficiently.

The binding affinity of hemoglobin varies depending on the ligand. For instance, the affinity for carbon monoxide (CO) is significantly higher than for oxygen. This is problematic in carbon monoxide poisoning, as CO competes with oxygen for the same binding site on the hemoglobin molecule, resulting in carboxyhemoglobin even at low concentrations of CO in the air.

This aspect of binding affinity is also why interventions to treat CO poisoning are necessary; since CO binds more tightly to hemoglobin than oxygen, it's challenging for the body to replace carboxyhemoglobin with oxyhemoglobin without assistance.