BF₃ Lewis Structure (Boron Trifluoride)
BF₃ has 24 valence electrons, three B–F single bonds, and three lone pairs on each fluorine. Boron has only six electrons around it—an accepted octet-rule exception. BF₃ is trigonal planar, with 120° bond angles, and is nonpolar.
Lewis Structure
Three B–F single bonds. Lone pairs: 3 on each fluorine.
Molecular Geometry
Trigonal planar shape · 120° · AX₃
| Total valence electrons | 24 |
|---|---|
| Central atom | Boron (B) |
| Bonds | Three B–F single bonds |
| Lone pairs on B | 0 |
| Lone pairs on F | 3 |
| Electron geometry | Trigonal planar |
| Molecular geometry | Trigonal planar |
| VSEPR | AX₃ |
| Bond angle | 120° |
| Polarity | Nonpolar |
| B hybridization | sp² |
| Resonance | Single preferred form; minor charge-separated alternatives |
| Formal charges | 0 on every atom |
| Octet rule | Exception — B has 6 electrons |
How to Draw the BF₃ Lewis Structure
Draw the Lewis structure of boron trifluoride step by step. Count the valence electrons, put boron in the center, form three B–F single bonds, complete each fluorine octet, then accept the incomplete octet on boron.
Step 1Count the valence electrons
Boron brings 3 valence electrons and each fluorine brings 7. With three fluorines that is 3 + 3(7) = 24 electrons to place — and BF₃ is neutral, so there is no charge to add or remove. Every valid Lewis drawing of boron trifluoride must use all 24.
Atom Count Valence e⁻ Total B 1 3 3 F 3 7 21 Total 24 Step 2Choose the central atom
Boron is less electronegative than fluorine and can form more than one bond, so boron sits in the middle. The three fluorines attach around it.
Step 3Draw three B–F single bonds
Connect each fluorine to boron with a single bond. Those three B–F bonds use 6 of the 24 valence electrons, leaving 18 electrons still to place.
Step 4Complete the fluorine octets
Give the remaining 18 electrons to the three fluorines as lone pairs — three lone pairs on each F. Every fluorine then has an octet (one bond + three lone pairs). Nothing is left over for a lone pair on boron.
Step 5Check boron's incomplete octet
Boron is surrounded by only six bonding electrons. That is the accepted Lewis structure for BF₃: three single bonds, formal charge 0 on every atom, and an incomplete octet on boron. Do not force three B=F double bonds just to reach eight on B — that would create large formal charges and is not the preferred structure.
- Lewis structure complete — all 24 valence electrons placed
- Each fluorine octet satisfied (1 bond + 3 lone pairs)
- Boron has 6 electrons (incomplete octet — accepted exception)
- Formal charges 0 on B and every F
BF₃ Valence Electrons
Boron contributes 3 valence electrons. The 3 fluorine atoms contribute 3 × 7 = 21 valence electrons. Added together, BF₃ starts with 24 valence electrons to place, and as a neutral molecule there is no charge to add or remove. Of those, 6 fill the bonds and 18 sit as lone pairs (3 on each fluorine).
| boron: 3 valence electrons |
| 3 × fluorine: 3 × 7 = 21 |
| Total: 3 + 21 = 24 |
BF₃ Resonance Structures
BF₃ is usually written as one Lewis structure with three B–F single bonds, not a set of equivalent resonance forms like O₃ or NO₃⁻.
Some texts also show minor charge-separated forms in which one fluorine double-bonds to boron (B −1, that F +1) so boron reaches an octet. Those forms are secondary. Boron is a period-2 element and does not expand its octet the way sulfur can in SO₂. For a single answer, use three single bonds and formal charge 0 on every atom.
Which BF₃ Lewis Structure Is Correct?
The Lewis structure to learn first for BF₃ has three B–F single bonds and formal charge 0 on every atom. Boron keeps only six electrons — an accepted incomplete-octet exception, not an error. Drawings that add a B=F double bond so boron reaches eight electrons introduce charge separation; they are minor alternatives, not the main structure.
| Representation | B octet | Formal charges | Use |
|---|---|---|---|
| Three B–F single bonds | 6 e⁻ on B | All 0 | Preferred Lewis structure |
| One B=F + two B–F | 8 e⁻ on B | B −1, double-bond F +1 | Minor charge-separated alternative |
BF₃ Lewis Structure with Formal Charges
Formula
Formal charge = valence e⁻ − nonbonding e⁻ − ½ × bonding e⁻
| Atom | Calculation | Formal charge |
|---|---|---|
| B | 3 − 0 − ½(6) = 0 | 0 |
| F | 7 − 6 − ½(2) = 0 | 0 |
| F | 7 − 6 − ½(2) = 0 | 0 |
| F | 7 − 6 − ½(2) = 0 | 0 |
The formal charges sum to 0, matching the neutral molecule. That makes this the preferred BF₃ Lewis structure: zero formal charge on every atom.
Formal charge = valence e⁻ − nonbonding e⁻ − ½(bonding e⁻).
Boron: 3 − 0 − ½(6) = 0. Boron has no lone pairs and three single bonds (6 bonding electrons).
Each fluorine: 7 − 6 − ½(2) = 0. Fluorine has 3 lone pairs (6 nonbonding electrons) and one single bond.
Every atom is zero and the charges sum to 0, so the three-single-bond drawing is the preferred Lewis structure for BF₃. Adding a B=F double bond so boron reaches eight electrons creates charge separation (B −1, double-bonded F +1). That form is a minor alternative, not the main structure.
Does BF₃ Follow the Octet Rule?
BF₃ does not give boron a full octet, and that is not a drawing mistake.
Boron forms three B–F single bonds, so only six electrons surround the center. Period-2 boron commonly keeps this incomplete octet rather than forcing extra bonds. You can write a charge-separated form with one B=F double bond so boron reaches eight electrons, but that puts formal charge −1 on boron and +1 on the double-bonded fluorine. The preferred structure keeps every formal charge at zero — three single bonds and six electrons on B.
That does not mean the real molecule has no π character at all. Experimental B–F bonds are shorter than a pure single bond would suggest, which is often explained by fluorine lone-pair donation into boron's empty p orbital. When you draw the Lewis structure, still start from three single bonds and formal charge 0 on every atom.
Three B–F single bonds
Boron has 6 electrons; every formal charge is 0. This is the preferred Lewis structure.
One B=F double bond + two single bonds
Boron reaches 8 electrons, but B is −1 and the double-bonded F is +1. A minor charge-separated alternative.
BF₃ Molecular Geometry
| Electron geometry | Trigonal planar |
|---|---|
| Molecular geometry / shape | Trigonal planar |
| VSEPR notation | AX₃ |
| Bond angle | 120° |
On the central boron, 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. BF₃ has 3 bonding domains and no lone pairs (AX₃), for 3 electron domains in total.
Those domains arrange to minimize repulsion around the boron, 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 BF₃ is AX₃: boron is the central atom, and the three B–F bonds are three bonding domains, with no lone pairs on boron. Three domains need about 120° of space, so both the electron geometry and the molecular geometry are trigonal planar. The F–B–F angle is 120°. An incomplete octet on boron does not change the domain count — VSEPR still reads three bonding regions and predicts the flat triangle. Ammonia makes a useful contrast: NH₃ also has three bonding domains, but nitrogen keeps one lone pair (AX₃E), so the molecule is trigonal pyramidal with an angle near 107° instead of a planar 120°.
Why Is BF₃ Trigonal planar?
Why is BF₃ trigonal planar rather than pyramidal or tetrahedral? Three bonding domains surround boron and none are lone pairs. Those three domains repel equally and settle in one plane, 120° apart. A tetrahedral layout needs four domains; a trigonal pyramidal layout needs a lone pair on the center. Because every domain is a bond to fluorine, the electron geometry and the molecular geometry are the same: both trigonal planar.
BF₃ Bond Angle
The F–B–F bond angle is 120°. That is the ideal trigonal planar angle: three equal bonding pairs and no lone pairs on boron, so nothing compresses the angle the way a lone pair does in NH₃ (~107°). For three electron domains with no lone pairs, 120° is the standard angle for boron trifluoride.
| BF₃ | 120° |
|---|---|
| NH₃ | ≈107° |
| CH₄ | 109.5° |
Is BF₃ Polar or Nonpolar?
BF₃ is nonpolar.
BF₃ is nonpolar. Its three B–F bonds are polar, but their bond dipoles cancel because the molecule has a symmetrical trigonal planar shape.
Fluorine (electronegativity 3.98) pulls shared electrons harder than boron (2.04), so every B–F bond dipole points toward fluorine (δ+ on B, δ− on each F). Bond polarity alone does not decide the molecule. Because BF₃ is trigonal planar (AX₃) with three identical B–F bonds and no lone pairs on boron, those three equal dipoles sit 120° apart in one plane and cancel. The net molecular dipole is zero.
CO₂ is also nonpolar, but for a linear reason: two equal C=O dipoles point opposite ways along one axis. NH₃ has polar bonds too, yet one lone pair on nitrogen (AX₃E) makes ammonia trigonal pyramidal and polar. Same three bonding domains as BF₃, different shape and polarity — BF₃ has no lone pair on the center.
Three equal B–F bond dipoles point toward F. The trigonal planar shape cancels them completely — net dipole 0.
BF₃ Hybridization
Boron in BF₃ is sp²-hybridized.
Boron in BF₃ is sp²-hybridized: three B–F bonds and no lone pairs give three electron domains. One 2s and two 2p orbitals mix into three sp² hybrids that point 120° apart in a plane and form the three B–F σ bonds. The remaining empty 2p orbital on boron sits perpendicular to that plane — that vacant orbital is why BF₃ can accept an electron pair and act as a Lewis acid. BF₃ has no π bonds in the usual single-bond Lewis drawing; every B–F link is a single σ bond.
Why Is BF₃ a Lewis Acid?
BF₃ is a Lewis acid, not a Lewis base. The incomplete octet leaves boron electron-deficient, so boron can accept a lone pair from a donor such as ammonia. In the adduct F₃B←NH₃, nitrogen shares its lone pair with boron; boron’s coordination number becomes 4 and the local geometry around boron shifts from trigonal planar toward tetrahedral.
NH₃ is the complementary Lewis base: it has a lone pair ready to donate. NH₃ and BF₃ are often drawn together because they form that adduct — not because free BF₃ needs a second Lewis structure.
| Property | BF₃ | NH₃ |
|---|---|---|
| Bonding domains | 3 | 3 |
| Lone pairs on center | 0 | 1 |
| VSEPR | AX₃ | AX₃E |
| Molecular geometry | Trigonal planar | Trigonal pyramidal |
| Bond angle | 120° | ≈107° |
| Polarity | Nonpolar | Polar |
| Lewis role | Lewis acid (electron-pair acceptor) | Lewis base (electron-pair donor) |
BF₃ Lewis Structure Summary
| Chemical name | Boron trifluoride |
|---|---|
| Total valence electrons | 24 |
| Central atom | Boron |
| Bond type | Three B–F single bonds |
| Lone pairs on boron | 0 |
| Lone pairs on outer atoms | 3 on each fluorine |
| Electron geometry | Trigonal planar |
| Molecular geometry | Trigonal planar |
| VSEPR | AX₃ |
| Bond angle | 120° |
| Hybridization | sp² |
| Polarity | Nonpolar |
| Formal charge | 0 on every atom |
| Resonance | Single preferred form; minor charge-separated alternatives |
| Overall charge | 0 |
| Octet exception | Incomplete octet on B (6 e⁻) — accepted |
Try Another Lewis Structure
Enter another formula to draw its Lewis structure — bonds, lone pairs, formal charges, and the steps behind them.
Need the mass instead? Open the molar mass calculator for BF₃ →
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 BF3?
The Lewis structure of BF₃ places boron in the center with three B–F single bonds and three lone pairs on each fluorine. It uses all 24 valence electrons. Boron has only six electrons around it (an incomplete octet), and every formal charge is 0.
How many valence electrons does BF3 have?
BF₃ has 24 valence electrons: 3 from boron and 7 from each fluorine (3 × 7 = 21). Six of them fill the three B–F single bonds; the other eighteen sit as three lone pairs on each fluorine.
Does BF3 follow the octet rule?
Not on boron. Each fluorine has an octet, but boron is surrounded by only six electrons from three single bonds. That incomplete octet is an accepted exception for boron, not a wrong Lewis structure.
What is the molecular geometry of BF3?
BF₃ has a trigonal planar molecular geometry. Boron has three bonding domains and no lone pairs (VSEPR AX₃), so the electron geometry and the molecular shape are both trigonal planar with an F–B–F angle of 120°.
What is the bond angle of BF3?
The F–B–F bond angle is 120°. That is the ideal angle for three equal bonding pairs with no lone pairs on the central atom.
Is BF3 polar or nonpolar?
BF₃ is nonpolar. Each B–F bond is polar, but the three bonds sit symmetrically in a trigonal planar AX₃ shape, so the dipoles cancel and the net molecular dipole is zero.
What is the hybridization of boron in BF3?
Boron in BF₃ is sp²-hybridized. Three electron domains (three B–F bonds) mix one s and two p orbitals into three sp² hybrids; the remaining empty p orbital can accept a lone pair when BF₃ acts as a Lewis acid.
Does BF3 have resonance structures?
The main Lewis structure is a single drawing with three B–F single bonds and formal charges of zero. Minor charge-separated forms with one B=F double bond appear in some texts, but they are not the preferred structure.
Why is BF3 a Lewis acid rather than a Lewis base?
BF₃ is a Lewis acid because boron is electron-deficient (only six electrons in the preferred structure) and has an empty p orbital that can accept a lone pair from a donor such as NH₃. A Lewis base donates a pair; BF₃ accepts one.