CH₂O Lewis Structure
Formaldehyde (Methanal)
The CH₂O Lewis structure has carbon as the central atom, two C–H single bonds, and one C=O double bond. Oxygen has two lone pairs. CH₂O has 12 valence electrons, a trigonal planar molecular geometry, and is polar.
Lewis Structure
Two C–H single bonds + One C=O double bond. Lone pairs: 2 on oxygen.
Molecular Geometry
Trigonal planar shape · ≈120° · AX₃
| Total valence electrons | 12 |
|---|---|
| Central atom | Carbon (C) |
| C–H | 2 × single bond |
| C=O | Double bond |
| Lone pairs on C | 0 |
| Lone pairs on O | 2 |
| Electron geometry | Trigonal planar |
| Molecular geometry | Trigonal planar |
| VSEPR | AX₃ |
| Bond angle | ≈120° |
| Polarity | Polar |
| C hybridization | sp² |
| O hybridization | sp² |
| Sigma bonds | 3 |
| Pi bonds | 1 |
| Resonance | Major H₂C=O; minor charge-separated form |
| Formal charges | 0 on every atom |
How to Draw the CH₂O Lewis Structure
Draw the Lewis structure of formaldehyde (CH₂O) step by step. Count 12 valence electrons, put carbon in the center, connect two hydrogens and one oxygen with single bonds, finish oxygen's octet, then form the C=O double bond so carbon reaches eight electrons too.
Step 1Count the valence electrons
Carbon brings 4 valence electrons, each hydrogen brings 1, and oxygen brings 6. CH₂O is neutral, so nothing is added or removed for charge: 4 + 2(1) + 6 = 12 valence electrons. Every valid Lewis drawing of formaldehyde must place all 12.
Atom Count Valence e⁻ Total C 1 4 4 H 2 1 2 O 1 6 6 Total 12 Step 2Choose the central atom
Carbon is less electronegative than oxygen and can form more than one bond, so carbon sits in the middle. Connect the two hydrogens and the oxygen to carbon first as three single bonds — a simple H–C(H)–O skeleton.
Step 3Complete the oxygen octet
Three single bonds use 6 of the 12 valence electrons, leaving 6 electrons. Place those six on oxygen as three lone pairs. Oxygen then has an octet (one bond + three lone pairs), but carbon still has only six electrons from the three single bonds.
Step 4Form the C=O double bond
Move one lone pair from oxygen into the C–O link to make a double bond. Carbon then reaches eight electrons (two C–H single bonds + one C=O double bond). Oxygen keeps two lone pairs. The double bond is required so carbon completes its octet.
Step 5Verify the structure
Carbon and oxygen each have 8 electrons; each hydrogen has 2. Bonding electrons total 8 and nonbonding electrons total 4, for 12 overall. Formal charge is 0 on every atom. The finished structure is H₂C=O — two C–H single bonds, one C=O double bond, and two lone pairs on oxygen.
- All 12 valence electrons placed (8 bonding + 4 nonbonding)
- Carbon octet satisfied
- Oxygen octet satisfied
- Each hydrogen duet satisfied
- Formal charges 0 on C, O, and both H
CH₂O Valence Electrons
Carbon contributes 4 valence electrons. The 2 hydrogen atoms contribute 2 × 1 = 2 valence electrons. Oxygen contributes 6 valence electrons. Added together, CH₂O starts with 12 valence electrons to place, and as a neutral molecule there is no charge to add or remove. Of those, 8 fill the bonds and 4 sit as lone pairs (2 on oxygen).
| carbon: 4 valence electrons |
| 2 × hydrogen: 2 × 1 = 2 |
| oxygen: 6 valence electrons |
| Total: 4 + 2 + 6 = 12 |
CH₂O Resonance Structures
The main Lewis structure of formaldehyde is H₂C=O, with formal charge 0 on every atom.
A secondary, charge-separated contributor can be written as H₂C⁺–O⁻, where the C=O π pair moves onto oxygen. That form puts +1 on carbon and −1 on oxygen. The two drawings are not equal partners the way SO₂ or O₃ resonance forms are. The neutral double-bond structure is the major contributor; the charge-separated form is minor. The double bond stays between carbon and oxygen — not on a hydrogen — and the two forms are not equal-weight resonance structures.
CH₂O Lewis Structure with Formal Charges
Formula
Formal charge = valence e⁻ − nonbonding e⁻ − ½ × bonding e⁻
| Atom | Calculation | Formal charge |
|---|---|---|
| C | 4 − 0 − ½(8) = 0 | 0 |
| H | 1 − 0 − ½(2) = 0 | 0 |
| H | 1 − 0 − ½(2) = 0 | 0 |
| O | 6 − 4 − ½(4) = 0 | 0 |
The formal charges sum to 0, matching the neutral molecule. That makes this the preferred CH₂O Lewis structure: zero formal charge on every atom.
Formal charge = valence e⁻ − nonbonding e⁻ − ½(bonding e⁻).
Carbon: 4 − 0 − ½(8) = 0. Carbon has no lone pairs and four bonding pairs around it (two single bonds + one double bond).
Each hydrogen: 1 − 0 − ½(2) = 0.
Oxygen: 6 − 4 − ½(4) = 0. Oxygen has two lone pairs (4 nonbonding electrons) and one double bond (4 bonding electrons).
Every atom is zero and the charges sum to 0, so H₂C=O is the preferred Lewis structure. A charge-separated H₂C⁺–O⁻ drawing is only a minor contributor beside that form.
CH₂O Molecular Geometry
| Electron geometry | Trigonal planar |
|---|---|
| Molecular geometry / shape | Trigonal planar |
| VSEPR notation | AX₃ |
| Bond angle | ≈120° |
On the central carbon, count electron domains the VSEPR way: each bonded neighbor is one domain (a double or triple bond still counts as one), and each lone pair is one more. CH₂O has 3 bonding domains and no lone pairs (AX₃), for 3 electron domains in total.
Those domains arrange to minimize repulsion around the carbon, so the molecule is trigonal planar with a bond angle of ≈120°. With no lone pairs on the central atom, the electron geometry and the molecular geometry are the same: both trigonal planar.
In VSEPR terms CH₂O is AX₃: carbon is the central atom, and the three bonding domains are the two C–H bonds plus the C=O bond, with no lone pairs on carbon. Three domains need about 120° of space, so both the electron geometry and the molecular geometry are trigonal planar. The ideal bond angle is ≈120°.
A double bond still counts as one electron domain, not two. Oxygen's lone pairs sit on oxygen, not on carbon, so they are not part of carbon's domain count. That is why CH₂O is trigonal planar, not bent and not tetrahedral.
Why Is CH₂O Trigonal planar?
Why is CH₂O trigonal planar rather than bent or tetrahedral? Three bonding domains surround carbon and none are lone pairs. Those three domains repel and settle in one plane, about 120° apart. A tetrahedral layout needs four domains on carbon; a bent layout needs lone pairs on the center (as in H₂O). Oxygen's two lone pairs affect only the oxygen end — they are not counted when naming the shape around carbon. Because every domain on carbon is a bond, electron geometry and molecular geometry match: both trigonal planar.
CH₂O Bond Angle
The ideal H–C–H and H–C–O angles are about 120° — the standard trigonal planar angle for three bonding domains with no lone pairs on carbon. Real gas-phase formaldehyde angles are close but not identical: the C=O double bond takes slightly more space than a C–H single bond, so the measured angles deviate a little from the ideal 120°. The usual VSEPR value remains ≈120°.
| CH₂O | ≈120° |
|---|---|
| BF₃ | 120° |
| NH₃ | ≈107° |
| CH₄ | 109.5° |
Is CH₂O Polar or Nonpolar?
CH₂O is polar.
CH₂O is polar. The C=O bond produces a strong bond dipole toward oxygen, and the molecule's trigonal planar arrangement does not cancel the overall dipole moment.
Oxygen (electronegativity 3.44) pulls shared electrons harder than carbon (2.55), so the carbonyl bond has δ− on O and δ+ on the carbonyl carbon. The two C–H bonds are only weakly polar and do not sit opposite the C=O vector the way the two C=O bonds cancel in linear CO₂. Because CH₂O is trigonal planar (AX₃) with unequal outer atoms — two hydrogens and one oxygen — the bond dipoles leave a net molecular dipole pointing toward oxygen.
CO₂ is also built from polar C=O bonds, but its linear, symmetric shape cancels those dipoles and CO₂ is nonpolar. H₂O is polar for a bent reason: two lone pairs on oxygen leave the O–H dipoles unbalanced. Same idea, different shape — CH₂O is planar and still polar because the ends are not equal.
The C=O bond dipole points toward O. The trigonal planar layout is not symmetric enough to cancel it — net molecular dipole remains.
| CH₂O | Trigonal planar · Polar |
|---|---|
| CO₂ | Linear · Nonpolar |
| H₂O | Bent · Polar |
CH₂O Hybridization
Carbon in CH₂O is sp²-hybridized; oxygen is sp²-hybridized.
| Carbon (C) | sp² |
|---|---|
| Oxygen (O) | sp² |
Carbon is sp² hybridized in CH₂O. The central carbon has three bonding electron domains — two C–H single bonds and one C=O double bond — so it mixes one 2s orbital with two 2p orbitals into three sp² hybrids that point about 120° apart in a plane. Those hybrids form the three σ bonds (two to H, one to O). The leftover p orbital on carbon overlaps sideways with a p orbital on oxygen to make the π bond of C=O.
Oxygen is often described as sp² hybridized as well: one hybrid forms the C–O σ bond, two hybrids hold the lone pairs, and a p orbital forms the π bond. Count for the whole molecule: 3 σ bonds and 1 π bond.
Sigma and Pi Bonds in CH₂O
Count σ and π from the finished Lewis structure.
Each C–H single bond is 1 σ (no π).
The C=O double bond is 1 σ + 1 π.
Whole molecule: 3 σ and 1 π. That matches the sp² picture on carbon: three hybrids build the σ framework, and one leftover p orbital on carbon forms the π bond with oxygen.
| Sigma bonds (σ) | 3 |
|---|---|
| Pi bonds (π) | 1 |
The sp² hybrid orbitals on carbon make the σ framework; any leftover p orbitals on the multiply bonded atoms form the π bonds.
CH₂O Lewis Structure Summary
| Chemical name | Formaldehyde |
|---|---|
| Total valence electrons | 12 |
| Central atom | Carbon |
| C–H | 2 × single bond |
| C=O | Double bond |
| Lone pairs on carbon | 0 |
| Lone pairs on outer atoms | 2 on oxygen |
| Electron geometry | Trigonal planar |
| Molecular geometry | Trigonal planar |
| VSEPR | AX₃ |
| Bond angle | ≈120° |
| Hybridization | Carbon sp²; Oxygen sp² |
| Polarity | Polar |
| Sigma bonds | 3 |
| Pi bonds | 1 |
| Formal charge | 0 on every atom |
| Resonance | Major H₂C=O; minor charge-separated form |
| Overall charge | 0 |
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Related Lewis Structures
Common next structures from general chemistry — each link opens that molecule so you can compare geometry, polarity, or the drawing steps.
Frequently Asked Questions
What is the Lewis structure of CH2O?
The Lewis structure of CH₂O (formaldehyde, also called methanal or H₂CO) places carbon in the center with two C–H single bonds and one C=O double bond. Oxygen has two lone pairs. The drawing uses all 12 valence electrons, and every formal charge is 0.
How many valence electrons does CH2O have?
CH₂O has 12 valence electrons: 4 from carbon, 1 from each hydrogen (2 total), and 6 from oxygen. Eight of them are bonding (two C–H single bonds and one C=O double bond); four sit as two lone pairs on oxygen.
How many lone pairs does CH2O have?
CH₂O has two lone pairs, both on oxygen. Carbon and the hydrogens have no lone pairs in the preferred Lewis structure.
What is the molecular geometry of CH2O?
CH₂O has a trigonal planar molecular geometry. Carbon has three bonding domains and no lone pairs (VSEPR AX₃), so the electron geometry and the molecular shape are both trigonal planar with ideal angles near 120°.
What is the bond angle of CH2O?
The ideal bond angles in CH₂O are about 120°, the standard trigonal planar VSEPR value. Real H–C–H and H–C–O angles are close to that ideal but not exactly equal, because the C=O double bond and the C–H bonds are not identical.
Is CH2O polar or nonpolar?
CH₂O is polar. The C=O bond has a strong dipole toward oxygen, and the trigonal planar layout with two hydrogens and one oxygen does not cancel that dipole, so a net molecular dipole remains.
What is the hybridization of CH2O?
The central carbon in CH₂O is sp²-hybridized. Three electron domains mix one s and two p orbitals into three sp² hybrids; the remaining p orbital forms the π bond of C=O. Oxygen is often described as sp² hybridized as well.
Does CH2O have resonance structures?
The main structure is H₂C=O with formal charges of zero. A minor charge-separated form is H₂C⁺–O⁻, but the two forms are not equal contributors the way SO₂ resonance forms are.