The bond in Carbon Monoxide (CO) is a highly polar covalent bond, meaning the sharing of electrons between the atoms is unequal. This molecule is formed from two nonmetal atoms. While the bond does not involve the complete transfer of electrons characteristic of an ionic bond, the sharing is uneven enough that it exhibits some ionic-like characteristics.
The Fundamental Difference Between Covalent and Ionic Bonds
Chemical bonds are generally categorized into two types: ionic and covalent. Ionic bonds form when one or more valence electrons are completely transferred from one atom to another. This usually happens between a metal and a nonmetal, creating oppositely charged ions (cations and anions) held together by attraction.
Covalent bonds form when two atoms share electrons between them, typically occurring between two nonmetal atoms. In a pure covalent bond, sharing is equal. However, in most cases, sharing is unequal, resulting in a polar covalent bond where the electron cloud shifts toward one atom.
Determining Bond Type Using Electronegativity
Chemists use electronegativity to classify where a bond falls on the spectrum between ionic and covalent. Electronegativity is defined as an atom’s ability to attract a shared pair of electrons toward itself in a chemical bond. The Pauling scale assigns a numerical value to this attraction, ranging from about 0.7 to 3.98 for the most attractive element, Fluorine.
The difference between the electronegativity values (\(\Delta\text{EN}\)) of two bonded atoms is used to classify the bond type. A difference of zero indicates a nonpolar covalent bond. If the difference is between approximately 0.4 and 1.7, the bond is polar covalent, meaning electrons are shared unequally. A difference greater than about 1.7 is classified as an ionic bond, indicating complete electron transfer.
Analyzing the Carbon Monoxide Bond Character
To classify the bond in carbon monoxide, we compare the electronegativity values for Carbon (C) and Oxygen (O). Carbon’s value is approximately 2.55, and Oxygen’s value is 3.44. This difference reflects Oxygen’s stronger pull on the shared electrons.
Calculating the electronegativity difference (\(\Delta\text{EN}\)) yields \(3.44 – 2.55 = 0.89\). This value falls squarely within the range for a polar covalent bond, which is typically defined as \(0.4 < \Delta\text{EN} < 1.7[/latex]. The calculation confirms that the bond in CO is covalent because electrons are shared rather than fully transferred. However, the difference is substantial, meaning the bond possesses a high degree of ionic character, with electron density skewed toward the oxygen atom. This unequal sharing creates a partial negative charge ([latex]\delta^-[/latex]) on the more electronegative Oxygen atom and a partial positive charge ([latex]\delta^+[/latex]) on the Carbon atom. The classification as a polar covalent bond, rather than ionic, is based on the fact that the atoms are both nonmetals and the calculated [latex]\Delta\text{EN}[/latex] does not cross the threshold for complete electron transfer.
The Triple Bond Structure and High Polarity of CO
The specific Lewis structure of the carbon monoxide molecule features a triple bond between the carbon and oxygen atoms. This triple bond consists of two normal covalent bonds and one coordinate covalent bond, also known as a dative bond.
In a coordinate covalent bond, one atom contributes both of the shared electrons. For CO, oxygen donates a lone pair to carbon to satisfy the octet rule for both atoms. This unique electron arrangement leads to an unusual distribution of formal charges: the more electronegative oxygen atom carries a formal positive charge of [latex]+1\), while the less electronegative carbon atom carries a formal negative charge of \(-1\).
Despite the formal charges suggesting a positive charge on oxygen, the actual electron density is still pulled toward the highly electronegative oxygen atom, as predicted by the \(\Delta\text{EN}\) calculation. This conflict between the formal charge and the electronegativity difference results in a relatively small but measurable dipole moment for the molecule. The dipole moment arrow points toward the oxygen atom, confirming that the high electronegativity of oxygen ultimately dominates the electron distribution, making carbon monoxide a highly polar molecule.