Which Of The Following Bonds Is A Nonpolar Covalent Bond

You're staring at a multiple-choice question on a chemistry exam. Four bond options. One answer. And your brain freezes because electronegativity difference* sounded clear in lecture but feels slippery now.

Been there. Most of us have.

The question "which of the following bonds is a nonpolar covalent bond" shows up constantly — high school quizzes, AP Chemistry, college gen chem, even the MCAT. It's not a trick question. But it is the kind where a shaky grasp of the fundamentals makes every option look plausible. Small thing, real impact.

Let's fix that. Not with a chart you'll memorize and forget, but with the actual logic so you can walk into any version of this question and know exactly what to look for.

What Is a Nonpolar Covalent Bond

At its simplest, a nonpolar covalent bond is a partnership of equals. The electron density sits right in the middle. Here's the thing — no partial charges. Day to day, no dipole moment. On the flip side, two atoms share a pair of electrons, and neither atom pulls that shared pair closer to itself. Just a clean, symmetric share.

That's the textbook definition. In practice, it happens when the two atoms have identical* or nearly identical* electronegativity values.

Electronegativity, if you need the refresher, is an atom's pull on shared electrons. Practically speaking, 98 (sometimes rounded to 4. The bigger the difference between two atoms, the more the bond tilts toward ionic. 0), cesium and francium down near 0.But linus Pauling gave us the scale we still use — fluorine at 3. The smaller the difference, the more covalent. And when the difference is essentially zero? 7. Nonpolar covalent.

The Cutoff That Matters

Here's the number you'll see in almost every textbook: 0.4.

An electronegativity difference (ΔEN) of 0.4 or less → nonpolar covalent.
Between 0.4 and ~1.Consider this: 7–1. 9 → polar covalent.
Above that → ionic character dominates.

Some sources use 0.5 as the line. On top of that, others say "less than 0. 5." The exact boundary isn't a law of physics — it's a convention. What matters is that you know your* professor's or exam's cutoff. Think about it: if you're studying for the AP exam, 0. 4 is the number. If your textbook says 0.5, use 0.Think about it: 5. Consistency beats precision here.

Why It Matters / Why People Care

You might wonder: does a 0.Practically speaking, 3 vs. 0.5 difference actually change anything in the real world?

Yes. And not just for test points.

Molecular polarity drives solubility, boiling points, membrane permeability, protein folding, drug design — the list goes on. Day to day, a molecule full of nonpolar covalent bonds (like O₂ or CH₄) behaves totally differently from one with polar bonds (like H₂O or NH₃). Oil and water don't mix because oil is nonpolar and water is polar. That's not a fun fact — it's the basis of cell membranes, detergent action, and why your salad dressing separates.

On exams, this concept is a gateway. So naturally, misidentify the bond type, and you'll mispredict molecular geometry, polarity, intermolecular forces, and reactivity. One wrong ΔEN calculation cascades into three wrong answers downstream.

So it's worth nailing.

How to Tell If a Bond Is Nonpolar Covalent

The process is mechanical. Do it the same way every time and you'll stop second-guessing.

Step 1: Identify the Two Atoms

Look at the bond in question. Not the whole molecule — just the two atoms directly bonded. Now, if the question gives you a molecule like CH₃Cl, you're evaluating C–H and C–Cl separately. Each bond gets its own classification.

Step 2: Look Up Electronegativity Values

Use the Pauling scale. You don't need to memorize the whole table, but you should* know the big ones cold:

If you're taking a standardized test, they'll either provide a table or expect you to know the trend: EN increases across a period, decreases down a group.

Continue exploring with our guides on what are the monomers of proteins called and what is a natural exponential function.

Step 3: Calculate the Difference

Subtract the smaller from the larger. Absolute value — direction doesn't matter for magnitude.

ΔEN = |EN₁ – EN₂|

Step 4: Compare to the Cutoff

≤ 0.4 (or 0.5) → nonpolar covalent.

cutoff → polar covalent (or ionic if way above).

That's it. Even so, four steps. The trap isn't the math — it's rushing or misidentifying the atoms.

Worked Examples

H–H
EN(H) = 2.20, EN(H) = 2.20
ΔEN = 0.00 → nonpolar covalent. Always. Diatomic elements (H₂, N₂, O₂, F₂, Cl₂, Br₂, I₂) are the easiest points on the test.

C–H
EN(C) = 2.55, EN(H) = 2.20
ΔEN = 0.35 → nonpolar covalent (by the 0.4 rule).
This one trips people up. Carbon and hydrogen are close enough that C–H bonds are treated as nonpolar in most introductory contexts. But — in more advanced organic chemistry, the slight polarity of C–H matters for acidity and reactivity. For gen chem? Nonpolar.

C–Cl
EN(C) = 2.55, EN(Cl) = 3.16
ΔEN = 0.61 → polar covalent. Not nonpolar.

O–H
EN(O) = 3.44, EN(H) = 2.20
ΔEN = 1.24 → polar covalent. Very polar. This is why water is weird.

Na–Cl
EN(Na) ≈ 0.93, EN(Cl) = 3.16
ΔEN = 2.23 → ionic. Not covalent at all.

Common Examples You'll See on Tests

Exam writers love certain bonds. Recognize them on sight and you save seconds per question.

The

The C–H Bond
As noted earlier, C–H bonds are often mistakenly labeled polar. While carbon (2.55) and hydrogen (2.20) have a ΔEN of 0.35—just below the 0.4 threshold—the bond’s slight polarity becomes significant in organic chemistry. As an example, in alkanes like methane (CH₄), the symmetry of the molecule cancels out any minor dipole moments, making the molecule nonpolar overall. That said, in molecules like chloroform (CHCl₃), the C–H bond’s polarity contributes to the molecule’s reactivity. In gen chem, though, C–H is typically classified as nonpolar to simplify discussions of molecular geometry and intermolecular forces.

The O–H Bond
With a ΔEN of 1.24, O–H is a classic polar covalent bond. This high polarity explains water’s unique properties, such as hydrogen bonding and high surface tension. The O–H bond’s polarity also drives acidity in compounds like alcohols and carboxylic acids, as the hydrogen can be donated in reactions.

The C–Cl Bond
At ΔEN = 0.61, C–Cl bonds are polar covalent. Chlorine’s electronegativity pulls electron density away from carbon, creating a dipole moment. This polarity affects molecular geometry and reactivity—for example, in chloroalkanes, the C–Cl bond’s dipole influences solvation and reaction pathways.

The Na–Cl Bond
With a ΔEN of ~2.23, Na–Cl is ionic, not covalent. Sodium’s low electronegativity (≈0.93) and chlorine’s high value create a full transfer of electrons, forming Na⁺ and Cl⁻ ions. This ionic character underpins the compound’s high melting point and conductivity in aqueous solution.

Conclusion
Mastering bond polarity is foundational to understanding molecular behavior. By systematically applying the four-step process—identifying atoms, referencing electronegativity values, calculating ΔEN, and comparing to the cutoff—you can avoid cascading errors in predicting molecular geometry, polarity, intermolecular forces, and reactivity. While exceptions and nuances exist (e.g., C–H in advanced contexts), sticking to the mechanical process ensures accuracy in introductory and standardized settings. Remember: precision in the basics empowers confidence in tackling complex chemical systems.