Question

Carbon monoxide forms complexes with low-valent metals. For example, \(\mathrm{Ni}(\mathrm{CO})_{4}\) and \(\mathrm{Fe}(\mathrm{CO})_{5}\) are well known. CO also forms complexes with the iron(II) ion in hemoglobin, which prevents the hemoglobin from acting in its normal way. Is CO a Lewis acid or a Lewis base?

Short answer

CO is a Lewis base.

Step-by-step solution

Define Lewis Acid and Lewis Base

A Lewis acid is a substance that can accept a pair of electrons, while a Lewis base is a substance that can donate a pair of electrons. This concept is based on the ability to accept or donate electron pairs rather than protons or hydroxide ions, which distinguishes it from other acid-base theories.

Identify the Role of CO

Carbon monoxide (CO) has a lone pair of electrons on the carbon atom. In metal carbonyls like \(\mathrm{Ni}(\mathrm{CO})_{4}\) and \(\mathrm{Fe}(\mathrm{CO})_{5}\), CO coordinates to the metal by donating its lone pair to form a bond with the metal atom.

Analyze the Interaction with Metals

The formation of complexes such as \(\mathrm{Ni}(\mathrm{CO})_{4}\) involves the donation of electron pairs from CO to the metal center. This donation is characteristic of a Lewis base, which provides electrons to form covalent bonds.

Conclusion on CO's Role

Since CO donates electron pairs to form bonds with metal atoms, it acts as a Lewis base. The fact that CO can interfere with hemoglobin by forming stable complexes further supports its role as a Lewis base because it donates electrons to coordinate with iron.

Key concepts

Lewis Acid

A Lewis acid is defined as an entity that can accept an electron pair. This definition emphasizes the role it plays in forming a covalent bond through the acceptance and utilization of electron pairs. Lewis acids have vacant orbitals or are positively charged, enabling them to readily accept electrons into their structures.
Some common examples include metal ions, which often have empty orbitals, ready to engage in this type of bonding. For instance, the transition metals, due to their complex electron arrangements, often serve as effective Lewis acids in various reactions.
In simpler terms, envision a Lewis acid as an electron pair receptor. These receptors are opportunistic and will actively search for electron donors during chemical reactions.
The interaction typically results in a molecule that is more stable than its individual parts. Recognizing this helps identify a substance's potential role in various chemical pathways.

Metal Carbonyls

Metal carbonyls are complexes composed of transition metal centers surrounded by carbon monoxide ligands. A ligand is an ion or molecule that binds to a central metal atom to form a coordination complex. In metal carbonyl complexes, carbon monoxide acts as a ligand by donating its lone pair of electrons from the carbon atom to the metal.
This electron donation forms an intriguing bond between the CO and the metal atom, showcasing CO's characteristic as a Lewis base. Not only do they form the bond through electron donation, but they also stabilize it through a process known as back-bonding.
Back-bonding occurs when the metal donates some of its electron density back into empty orbitals on the CO ligand. This dual interaction makes metal carbonyl complexes fascinating subjects for study in organometallic chemistry.
Examples of metal carbonyls include \[\ce{Ni(CO)4}\] and \[\ce{Fe(CO)5}\], where metals like nickel and iron are perturbed by these electron-rich carbon monoxide molecules. These interactions explain the versatility and reactivity of metal carbonyls in industrial catalysis and other chemical processes.

Hemoglobin and Carbon Monoxide

Hemoglobin, a protein in red blood cells, is primarily responsible for transporting oxygen throughout the body. However, hemoglobin can also bind to carbon monoxide (CO). This interaction is particularly dangerous because CO has a higher affinity for hemoglobin than oxygen does.
When CO binds to the iron in hemoglobin, it prevents the normal binding and release of oxygen. This displacement occurs because CO donates a pair of electrons to the iron, forming a stable complex. Consequently, this blocks the sites that would usually carry oxygen, leading to decreased oxygen delivery in the body.
The preferential binding of CO to hemoglobin is part of what makes carbon monoxide poisoning so insidious. Victims can suffer from a lack of oxygen despite having hemoglobin in their bloodstream. This is a classic example of CO acting as a Lewis base, where it donates electrons to the metal ion in hemoglobin, interfering with its normal physiological function.
Understanding CO's interaction with hemoglobin highlights the importance of preventing CO exposure in enclosed environments, ensuring sources of carbon monoxide, like stoves and heaters, are adequately vented.

Electron Pair Donation

Electron pair donation is the key mechanism through which Lewis bases like carbon monoxide interact with other substances. In this process, a Lewis base offers its lone electron pairs, usually those present in its outer shell, to form covalent bonds with Lewis acids.
Such interactions are crucial in both natural and industrial chemical processes. For instance, in metal carbonyls, the electron pair donation from CO leads to stable metal-ligand complexes. Here the metal acts as a Lewis acid, accepting the electrons provided by CO, the Lewis base.
This concept of electron pair donation extends beyond just metal carbonyl systems. It is a vital component in nucleophilic reactions, catalysis, enzyme function, and more.
By understanding electron pair donation, students can predict and rationalize the behavior of complex molecules. Recognizing this electron-sharing relationship sheds light on why certain reactions occur, and how complex structures like those in organometallic chemistry are stabilized.