How to Draw the Lewis Structure for the Polyatomic Formyl (HCO)

The formyl radical (HCO) has 13 total valence electrons — 1 from hydrogen, 4 from carbon, and 6 from oxygen — which means it is an odd-electron (radical) species that cannot satisfy the octet rule for every atom. To draw the Lewis structure for the polyatomic formyl, place carbon in the center bonded to hydrogen and oxygen, then distribute the remaining electrons so carbon and oxygen each hold a single unpaired electron to reach the most stable arrangement.

Count Valence Electrons

Start by tallying the electrons each atom brings. Hydrogen contributes 1 valence electron, carbon contributes 4, and oxygen contributes 6. That gives a total of 13 valence electrons for HCO.

That odd total is the single most important clue about this molecule. An even electron count usually lets every atom pair up its electrons into bonds and lone pairs. An odd count means at least one electron must remain unpaired — which is exactly what makes formyl a radical (a species with an unpaired electron). The International Union of Pure and Applied Chemistry classifies such odd-electron species as radicals precisely because of this unpaired electron.

Because 13 electrons can’t be split evenly into pairs, you already know the finished structure will contain a lone, unbonded electron. Keep that in mind as you build the skeleton.

Skeleton and Bonding

Carbon sits in the center because it forms the most bonds and has the lowest electronegativity of the two heavy atoms. Arrange the atoms as H–C–O.

Step 1: Single bonds

Connect H to C and C to O with single bonds. Each bond uses 2 electrons, so two bonds consume 4 electrons, leaving 9 electrons to place.

Step 2: Complete the outer atoms

Hydrogen can only hold 2 electrons, and its single bond already gives it that. Oxygen needs a full octet, so give oxygen three lone pairs (6 electrons). That uses 6 more electrons, leaving 3 electrons.

Those remaining 3 electrons go on carbon as one lone pair plus one single unpaired electron. At this stage carbon has only 6 electrons around it (4 from bonds, 2 from the lone pair), so it falls short of an octet.

Step 3: Form the double bond

To give carbon a better electron count, move one of oxygen’s lone pairs into a second C–O bond. Now carbon–oxygen is a double bond, and carbon’s count improves. Oxygen still keeps two lone pairs and a full octet. This rearrangement is the standard bonding pattern for HCO; the Royal Society of Chemistry maintains reference data confirming the H–C=O connectivity of formyl.

After this step, count again: 2 (H–C bond) + 4 (C=O double bond) + 4 (oxygen’s two lone pairs) + 3 (carbon’s lone pair plus unpaired electron) = 13 electrons. All electrons are accounted for.

Formal Charges and Best Structure

Formal charge tells you which arrangement is most stable. Use the formula: Formal charge = (valence electrons) − (lone pair electrons) − ½(bonding electrons).

Work through each atom in the double-bonded structure:

• Hydrogen: 1 − 0 − ½(2) = 0
• Carbon: 4 − 2 − ½(6) = −1
• Oxygen: 6 − 4 − ½(4) = 0

This gives carbon a formal charge of −1 and an unpaired electron. An alternative single-bonded structure (C–O with carbon holding a lone pair and three unpaired electrons) is higher in energy and less favorable, so the double-bonded version is preferred.

Resonance considerations

Formyl’s unpaired electron is not strictly fixed on carbon. A minor resonance form places it on oxygen, with the C–O bond becoming single and carbon taking a full lone pair. The major contributor, however, keeps the C=O double bond with the radical centered on carbon. This is consistent with how the American Chemical Society describes radical stabilization: delocalization spreads the unpaired electron, but one structure dominates.

The takeaway is that the best structure has a C=O double bond, a C–H single bond, an unpaired electron on carbon, and two lone pairs on oxygen.

Final Lewis Structure

Here is the completed HCO Lewis structure, described atom by atom:

• H–C bond: a single line (2 electrons) between hydrogen and carbon.
• C=O bond: a double line (4 electrons) between carbon and oxygen.
• Oxygen: two lone pairs (4 electrons), giving oxygen a full octet and a formal charge of 0.
• Carbon: one lone pair (2 electrons) and one unpaired dot (1 electron), giving carbon a formal charge of −1.
• Hydrogen: no lone pairs; satisfied with its 2 shared electrons.

Written in line notation, the structure looks like this:

H − C(=O)• with a lone pair on carbon

Where:

• The = marks the double bond to oxygen.
• The • marks carbon’s single unpaired (radical) electron.
• Carbon also carries one lone pair.

Total electrons: 2 (H–C) + 4 (C=O) + 4 (O lone pairs) + 2 (C lone pair) + 1 (C unpaired) = 13, matching the original count. Hydrogen is satisfied, oxygen has a complete octet, and carbon carries the radical character with a −1 formal charge — the defining features of the formyl radical.

Why Formyl Is a Radical, Not a Typical Molecule

Most Lewis structures you meet in an introductory chemistry course have an even number of valence electrons, so every electron finds a partner. Formyl breaks that pattern. With 13 electrons, one must remain alone, and that lone electron is what chemists call the radical center. Radicals are unusually reactive because the unpaired electron is hungry for a partner — it will readily combine with another radical, abstract an atom from a neighboring molecule, or add to a double bond.

This reactivity explains why formyl appears as a short-lived intermediate rather than a stable compound you can bottle. It shows up in combustion chemistry, in atmospheric reactions, and in the breakdown of formaldehyde (H₂C=O), where a hydrogen atom is stripped away to leave HCO.

Radical vs. ion: a quick contrast

It helps to compare formyl with species that look similar but behave very differently.

• Formyl radical (HCO): neutral overall, 13 electrons, unpaired electron on carbon, formal charge −1 on carbon balanced by the radical.
• Formyl cation (HCO⁺): 12 electrons, loses the radical electron, often written with a C≡O triple bond.
• Formyl anion (HCO⁻): 14 electrons, gains an electron, giving carbon a full lone pair and a complete octet.

These species differ by both charge and electron count, which is why the radical label matters — it is not the same as an ion.

Key Takeaways

• Formyl (HCO) has 13 valence electrons, making it an odd-electron radical with one unpaired electron.
• Carbon is the central atom, bonded to one hydrogen and one oxygen.
• The most stable structure uses a C=O double bond and a C–H single bond.
• Oxygen holds two lone pairs with a formal charge of 0; carbon carries one lone pair, one unpaired electron, and a −1 formal charge.
• A minor resonance form shifts the radical toward oxygen, but the carbon-centered form is the major contributor.
• Hydrogen follows the duet rule; oxygen satisfies the octet rule; carbon cannot fully complete its octet because of the odd electron count.