The C–H bond is technically slightly polar, but it’s so close to nonpolar that chemists treat it as nonpolar in nearly every practical context. The electronegativity difference between carbon (2.55) and hydrogen (2.20) is only 0.35 on the Pauling scale, which falls well below the threshold where polarity starts to matter in chemical behavior.
Why the C–H Bond Is Nearly Nonpolar
Polarity in a bond comes down to how unevenly two atoms share their electrons. When one atom pulls electrons more strongly than the other, the bond develops a slight charge imbalance: one end becomes slightly negative, the other slightly positive. The bigger the difference in pull (measured as electronegativity), the more polar the bond.
Carbon has an electronegativity of 2.55 and hydrogen sits at 2.20, giving a difference of just 0.35. For comparison, an O–H bond has a difference of about 1.4, and an Na–Cl bond comes in around 2.1. The C–H bond’s 0.35 gap is small enough that most chemistry textbooks and courses classify it as nonpolar covalent. A single C–H bond does have a tiny dipole moment of about 1.57 debye in isolation, but this small charge separation rarely influences how a molecule behaves.
How Symmetry Cancels Out What Little Polarity Exists
Even if individual C–H bonds carry a whisper of polarity, molecules built from them often have no net polarity at all. Methane (CH₄) is the classic example. Its four C–H bonds point outward in a perfectly symmetrical tetrahedral shape, meaning the tiny pull of each bond is exactly counterbalanced by the others. The measured dipole moment of methane is zero. If the molecule had any other shape, those small individual bond dipoles wouldn’t cancel perfectly, and you’d detect some overall polarity. The fact that methane registers at exactly zero confirms its perfect symmetry.
This pattern holds for other simple hydrocarbons too. Ethane, propane, and other molecules built mostly from C–H and C–C bonds (where the electronegativity difference is 0.0) behave as nonpolar compounds.
How C–H Compares to Truly Polar Bonds
Putting C–H next to other common bonds makes its near-nonpolar nature obvious:
- C–C: 0.0 difference, completely nonpolar
- C–H: 0.4 difference, essentially nonpolar
- C–N: 0.5 difference, slightly polar
- N–H: 0.9 difference, polar
- C–O: 1.0 difference, polar
- O–H: 1.4 difference, strongly polar
The jump from C–H at 0.4 to N–H at 0.9 is where polarity starts to meaningfully affect a molecule’s properties. Bonds with differences above roughly 0.5 begin participating in stronger intermolecular attractions and influencing solubility, boiling points, and reactivity in noticeable ways.
Why It Matters: Solubility and Molecular Behavior
The practical consequence of C–H being nonpolar is that molecules dominated by C–H bonds don’t mix with water. Water is highly polar, and the rule “like dissolves like” means polar substances dissolve in polar solvents while nonpolar substances dissolve in nonpolar solvents. This is exactly why oil (long chains of C–H bonds) and water don’t mix. Hydrocarbons like hexane, benzene, and toluene are classic nonpolar solvents precisely because their C–H frameworks carry no significant charge imbalance.
The hydrocarbon portion of any organic molecule is described as hydrophobic, meaning it resists dissolving in water. When you add a polar group like O–H to a hydrocarbon chain (creating an alcohol, for instance), the molecule gains some water solubility. But the C–H sections still resist it. Short alcohols like methanol and ethanol dissolve freely in water because their polar O–H group dominates the small hydrocarbon portion. Longer alcohols become increasingly insoluble as the nonpolar C–H chain grows.
The “Weakly Polar” Nuance
Some chemistry sources describe C–H as “weakly polar” rather than flatly nonpolar, and that’s technically more precise. Carbon is slightly more electronegative than hydrogen, so it does pull the shared electrons a tiny bit closer to itself. This gives the carbon end a faint negative charge and the hydrogen end a faint positive charge. In certain specialized contexts, this matters. For example, C–H bonds next to strongly electron-withdrawing groups (like fluorine or oxygen) can become polar enough to participate in weak hydrogen bonding. Chemists studying crystal structures and protein folding sometimes account for these subtle C–H interactions.
For general chemistry purposes, though, the C–H bond is nonpolar. The electronegativity difference is too small to drive meaningful polarity in the molecules where these bonds appear, and symmetric arrangements cancel out even that small effect. If you’re answering an exam question or trying to predict whether a molecule dissolves in water, treating C–H as nonpolar will give you the right answer.