Carbon monosulfide (CS) is a polar molecule, with a measured dipole moment of 1.958 Debye. This surprises many students because the electronegativity difference between carbon and sulfur is nearly zero, which would normally suggest a nonpolar bond. The reality is more nuanced, and understanding why CS is polar teaches an important lesson about the limits of using electronegativity alone to predict polarity.
Why Electronegativity Alone Is Misleading Here
On the Pauling scale, carbon has an electronegativity of 2.55 and sulfur has an electronegativity of 2.58. That gives a difference of just 0.03, which is small enough that most textbooks would classify the bond as nonpolar covalent. For simple single bonds, this rule of thumb works well. A difference below about 0.4 generally means the two atoms share electrons fairly evenly.
But CS is not a simple single-bonded molecule. Carbon and sulfur form a triple bond (similar to carbon monoxide, CO), and in triple bonds, the distribution of electrons becomes more complex. Lone pairs, bonding pairs, and the way orbitals overlap all contribute to how charge is distributed across the molecule. In CS, this results in a significant shift of electron density that electronegativity values alone cannot predict.
The Measured Dipole Moment
The dipole moment of CS has been measured experimentally at 1.958 Debye (±0.005), according to NIST data. For context, water has a dipole moment of about 1.85 Debye, so CS is actually slightly more polar than water in terms of its individual bond. That is a remarkably large value for two atoms with nearly identical electronegativities.
This large dipole arises from the triple bond structure. When carbon and sulfur form a triple bond, the lone pair arrangements on each atom are asymmetric. Carbon ends up with a concentration of electron density that creates a meaningful charge separation across the molecule. The situation closely mirrors carbon monoxide (CO), where oxygen is more electronegative than carbon yet the dipole moment points in a direction that surprises students who rely only on electronegativity charts.
Molecular Shape and Symmetry
CS is a diatomic molecule, meaning it consists of just two atoms connected by a single bond (in this case, a triple bond). Its point group is C∞v, which is the symmetry classification for all heteronuclear diatomic molecules, meaning it looks the same when rotated around the bond axis but has no mirror symmetry that would cancel out a dipole. Because there is only one bond and no opposing bonds to cancel the charge separation, any polarity in the bond translates directly into a molecular dipole.
Compare this to a molecule like carbon dioxide (CO₂), where two polar bonds point in opposite directions and cancel each other out, making the overall molecule nonpolar. In CS, there is nothing to cancel. Whatever charge imbalance exists in the bond becomes the dipole of the whole molecule.
CS vs. CO: A Helpful Comparison
Carbon monoxide is the closest analog to carbon monosulfide. Both are diatomic molecules where carbon forms a triple bond with a group 16 element. CO has a dipole moment of about 0.11 Debye, which is surprisingly small given the electronegativity difference between carbon (2.55) and oxygen (3.44). In CO, the lone pair on carbon partially offsets the electronegativity-driven charge pull toward oxygen, resulting in a tiny net dipole.
CS flips this situation. The electronegativity difference is negligible, yet the lone pair and bonding electron distribution create a dipole moment of nearly 2 Debye. The takeaway is that for molecules with multiple bonds, lone pair effects and orbital interactions can dominate over simple electronegativity differences.
Where CS Exists
Carbon monosulfide is not a molecule you will encounter in everyday life. It is highly reactive under normal conditions and tends to polymerize (link together into chains) rather than remain as individual CS molecules. In the lab, it has been studied as a solid polymer that forms an orange-brown crystalline material.
CS is best known as an interstellar molecule. It was first detected in space in 1971 through radio astronomy, and it has since been observed in interstellar clouds and around carbon-rich stars. Astronomers use its microwave spectral lines to study the composition and conditions of molecular clouds throughout the galaxy. Its strong dipole moment is actually what makes it detectable: molecules with larger dipoles produce stronger rotational spectral signals, making CS a useful tracer molecule in astrophysics.
The Bottom Line on Polarity
CS is polar. If you are answering a homework question, the key facts are: the electronegativity difference is only 0.03, which would suggest a nonpolar bond, but the experimentally measured dipole moment is 1.958 Debye, confirming that the molecule is polar. The discrepancy comes from the triple bond structure, where lone pair and orbital effects create a charge separation that electronegativity values do not capture. This makes CS a textbook example of why electronegativity is a useful starting point for predicting polarity but not always the final word.