Carbon Monoxide (CO) is a slightly polar molecule, possessing a small, permanent net dipole moment. This classification often confuses people because the molecule is linear and composed of only two atoms, which typically suggests a simple, nonpolar structure. However, the unique electronic distribution between the carbon and oxygen atoms prevents the complete cancellation of charge expected in a perfectly balanced molecule.
Understanding Molecular Polarity
Molecular polarity describes the distribution of electrical charge across an entire molecule, which influences how that molecule interacts with others. This property is determined by two separate, yet interconnected, chemical concepts: bond polarity and molecular geometry. Bond polarity arises from the difference in electronegativity between two bonded atoms, where electronegativity is an atom’s inherent ability to attract electrons toward itself in a chemical bond. If this difference is zero, the bond is nonpolar.
If the electronegativity difference is non-zero, the bond becomes polar, with the shared electrons spending more time around the more electronegative atom, creating a partial negative charge (\(\delta^-\)) and a partial positive charge (\(\delta^+\)). However, the overall molecular polarity, or the presence of a net dipole moment, depends on the molecule’s three-dimensional shape.
A molecule can contain polar bonds but still be nonpolar overall if its geometry is perfectly symmetrical, causing the individual bond dipoles to cancel each other out. Carbon dioxide (\(\text{CO}_2\)), for example, has two polar carbon-oxygen bonds, but its linear shape means the dipoles pull equally in opposite directions, resulting in a net dipole moment of zero.
The Unique Structure of Carbon Monoxide
Carbon monoxide is a diatomic molecule composed of one carbon (C) and one oxygen (O) atom. These two atoms are connected by a triple covalent bond, which involves the sharing of six electrons to satisfy the octet rule for both atoms. Oxygen has an electronegativity value of approximately 3.44, while carbon’s value is 2.55, resulting in a difference of 0.89. This difference classifies the carbon-oxygen triple bond as a polar covalent bond, with electron density expected to be pulled toward the more electronegative oxygen atom.
This electron sharing is complicated by the concept of formal charges in the stable Lewis structure. Calculating the formal charge for this arrangement reveals a charge of \(-1\) on the carbon atom and a charge of \(+1\) on the oxygen atom. This charge distribution, which places a negative formal charge on the less electronegative carbon, is a unique feature of the carbon monoxide molecule.
The presence of the formal charges suggests an internal pull of electron density towards the carbon, which directly competes with the external pull toward the oxygen due to its higher electronegativity. This structural complexity means the overall charge distribution is more nuanced than a simple comparison of electronegativity values would suggest.
Determining Carbon Monoxide’s Net Dipole Moment
The net dipole moment of a molecule is a vector quantity that represents the overall magnitude and direction of the charge separation. In carbon monoxide, the two competing electronic effects—the electronegativity difference and the formal charge distribution—almost perfectly counteract each other. Oxygen’s higher electronegativity suggests a \(\text{C}^{(+)}\text{O}^{(-)}\) polarization, while the formal charges suggest a \(\text{C}^{(-)}\text{O}^{(+)}\) polarization.
Because carbon monoxide is a diatomic molecule, its molecular geometry is inherently linear, meaning the bond dipole is equivalent to the net molecular dipole. The slight imbalance between the two opposing electronic effects results in a very small, non-zero net dipole moment, measured to be approximately \(0.122\) Debye (D) in the gas phase. This value is significantly lower than what a simple polar bond might suggest, confirming that the formal charge contributions play a substantial role in reducing the overall polarity.
The direction of this small net dipole moment is also counterintuitive: it points away from the oxygen and toward the carbon, meaning the carbon end carries a small net negative charge (\(\delta^-\)), and the oxygen end carries a small net positive charge (\(\delta^+\)). This indicates that the effect of the formal charges slightly outweighs the effect of the electronegativity difference. Thus, carbon monoxide is classified as a weakly polar molecule because the electronic tug-of-war results in a residual, permanent charge separation.
Why Carbon Monoxide’s Polarity Matters
The slight, permanent polarity of carbon monoxide has consequences for its chemical reactivity and biological effects. This small net dipole moment allows the molecule to interact with external electric fields and other polar molecules, influencing properties such as its solubility in various solvents and its ability to act as a ligand in coordination chemistry.
Most notably, carbon monoxide’s structure and electronic distribution enable its toxicity to humans and animals. The \(\text{C}^{(-)}\text{O}^{(+)}\) polarization makes it an excellent ligand for binding to metal ions. In the body, carbon monoxide binds to the iron (\(\text{Fe}^{2+}\)) atom in the heme group of hemoglobin, which transports oxygen in the blood, forming a stable complex called carboxyhemoglobin. This strong, preferential binding prevents oxygen from being delivered to tissues, leading to cellular hypoxia and poisoning.