Linear, bent, and trigonal planar are three common molecular shapes that describe how atoms are arranged around a central atom in a molecule. The main difference is the angle between the bonds and whether the atoms lie along a straight line, form an angle, or spread out in a flat, triangular arrangement.
A linear molecule has a bond angle of 180°, a bent molecule has an angle smaller than 180°, and a trigonal planar molecule has three bonds arranged around a central atom at approximately 120° angles. These shapes influence how molecules interact with one another, whether they have an overall electrical polarity, and how they behave in chemical reactions.
Understanding these differences begins with the forces that determine molecular geometry.
What determines a molecule’s shape?
Molecules do not form their shapes randomly. Their atoms arrange themselves according to the distribution of electrons, particularly the electron pairs around a central atom.
A useful model for predicting molecular geometry is the valence shell electron pair repulsion (VSEPR) theory. It states that regions of electron density around a central atom arrange themselves to minimize repulsion. These regions include bonding pairs, which help form chemical bonds, and lone pairs, which are valence electrons not shared in a bond.
A single bond, double bond, or triple bond generally counts as one region of electron density when predicting molecular shape. Although multiple bonds contain more than one shared electron pair, they occupy one general direction around the central atom.
Lone pairs also occupy space and repel other electron regions. Because lone pairs are associated with only one atom, their electron density can exert a stronger influence on nearby bonding regions than a bonding pair does. This helps explain why some molecules bend even when their electron regions would otherwise favor a more symmetrical arrangement.
It is important to distinguish electron geometry from molecular geometry. Electron geometry describes the arrangement of all electron regions around a central atom, including lone pairs. Molecular geometry describes the positions of the atoms alone. The two can differ when the central atom has lone pairs.
Linear molecular shape: Atoms in a straight line
A linear molecule has atoms arranged along a single straight line. When a central atom is bonded to two surrounding atoms, the ideal bond angle is 180°.
The simplest example is carbon dioxide (CO₂). Its carbon atom sits between two oxygen atoms, forming the arrangement O=C=O. The two regions of electron density associated with the carbon–oxygen double bonds point in opposite directions, producing a linear structure.
Another example is hydrogen cyanide (HCN), which has the arrangement H–C≡N. Carbon is bonded to hydrogen and nitrogen, and the angle between those bonds is approximately 180°.
In both examples, the central atom has two bonding regions and no lone pairs that alter the arrangement. The regions spread as far apart as possible, creating a straight-line geometry.
Linear geometry can also occur when the central atom has more than two electron regions, provided some of those regions are lone pairs. For example, a molecule such as xenon difluoride (XeF₂) is linear because its three lone pairs and two bonding regions produce a symmetrical arrangement in which the two fluorine atoms lie opposite one another.
The linear shape has important consequences for molecular polarity. In carbon dioxide, each carbon–oxygen bond is polar because oxygen attracts electrons more strongly than carbon does. However, the two bond dipoles point in opposite directions and have equal magnitudes, so they cancel. The molecule as a whole is nonpolar.
This cancellation is not guaranteed for every linear molecule. If the two ends contain different atoms, or if the bond dipoles differ in magnitude, the molecule may have a net dipole and therefore be polar.
Bent molecular shape: Atoms forming an angle
A bent molecule has two atoms bonded to a central atom, but the three atoms do not lie in a straight line. Instead, the bonds form an angle that is less than 180°.
Water (H₂O) is a familiar example. Its oxygen atom is bonded to two hydrogen atoms and also has two lone pairs. The four electron regions around oxygen have a roughly tetrahedral electron geometry, but the positions of the atoms form a bent molecular shape. The H–O–H bond angle is approximately 104.5°.
The two lone pairs help explain why water is bent rather than linear. Their electron density repels the bonding regions, changing the arrangement of the hydrogen atoms. The resulting angle is smaller than the ideal tetrahedral angle of about 109.5°.
Sulfur dioxide (SO₂) is another example of a bent molecule. Its central sulfur atom is surrounded by two bonding regions and a lone pair in the basic VSEPR description. The electron regions have a trigonal planar arrangement, while the positions of the atoms form a bent shape. Its bond angle is close to 120°, though the precise geometry depends on the molecule’s electronic structure.
These examples illustrate why bent molecules do not all have the same bond angle. The angle depends on the number of electron regions, the presence and arrangement of lone pairs, and the nature of the bonds. Water’s angle is substantially smaller than that of sulfur dioxide because their central atoms have different electron-region arrangements.
Bent geometry also affects polarity. In water, each oxygen–hydrogen bond is polar, and the bond dipoles point in directions that do not cancel. Water therefore has a net dipole and is a polar molecule. This polarity contributes to water’s ability to interact strongly with other polar molecules and form hydrogen bonds.
A bent molecule is not necessarily polar in every conceivable case. Polarity depends on the identities of the atoms and the vector sum of the bond dipoles. However, bent molecules with polar bonds commonly have a net dipole because their bond directions are not opposite one another.
Trigonal planar molecular shape: Three atoms in a flat arrangement
A trigonal planar molecule has a central atom bonded to three surrounding atoms. All four atoms lie in the same plane, with the three bonds spreading around the central atom at approximately 120° intervals.
Boron trifluoride (BF₃) is a classic example. Boron occupies the center, and three fluorine atoms surround it in a flat triangular arrangement. The three B–F bonds point in different directions, separated by bond angles of approximately 120°.
In the usual VSEPR description, boron has three bonding regions and no lone pairs. These regions spread evenly around the central atom in a plane, minimizing their mutual repulsion.
Ethene (C₂H₄), also called ethylene, provides another example of trigonal planar geometry. Each carbon atom is bonded to two hydrogen atoms and the other carbon atom. Around each carbon, there are three regions of electron density, so the local arrangement is approximately trigonal planar. The entire molecule is planar because the atoms around both carbon atoms lie in the same plane.
Trigonal planar geometry often appears in molecules containing atoms that form three bonds with no lone pairs in the relevant arrangement. It is also common around carbon atoms involved in double bonds, where the bonding structure favors a flat arrangement.
The shape has consequences for polarity. Boron trifluoride has three polar B–F bonds, but its symmetrical trigonal planar arrangement causes the bond dipoles to cancel. The molecule is therefore nonpolar overall, even though its individual bonds are polar.
A trigonal planar molecule with three different surrounding atoms may not have the same symmetry, so its bond dipoles may not cancel. As with linear and bent molecules, molecular polarity depends on both bond polarity and three-dimensional arrangement.
How the three shapes compare
The easiest way to distinguish linear, bent, and trigonal planar geometries is to count the atoms attached to the central atom, consider the lone pairs, and examine the bond angles.
| Feature | Linear | Bent | Trigonal planar |
|---|---|---|---|
| Atoms bonded to the central atom | Usually 2 | Usually 2 | 3 |
| Basic molecular arrangement | Straight line | V-shaped or angular | Flat triangular arrangement |
| Typical bond angle | 180° | Less than 180° | Approximately 120° |
| Example | CO₂ | H₂O | BF₃ |
| Typical electron geometry in the basic examples | Linear | Tetrahedral for H₂O; trigonal planar for SO₂ | Trigonal planar |
| Effect of lone pairs | May be absent or help produce a linear shape | Often responsible for the bend | Absent in the basic example |
The table describes common examples rather than every possible molecular arrangement. In particular, the bond angle of a bent molecule varies, and linear molecules can have lone pairs on the central atom. The key distinction is the arrangement of the atoms, not simply the total number of electrons or bonds.
Why bond angles differ
Bond angles arise from the way electron regions distribute themselves around a central atom. In the simplest cases, two regions point in opposite directions at 180°, while three equivalent regions spread around a plane at 120°. Four electron regions generally adopt a tetrahedral arrangement, with ideal angles of about 109.5°.
Real molecules may depart from these ideal angles because electron regions do not always repel one another equally. Lone pairs tend to exert greater repulsive influence on nearby bonding regions than bonding pairs exert on one another. Multiple bonds can also influence the surrounding geometry because their electron density is distributed differently from that of single bonds.
Water demonstrates the effect clearly. Its four electron regions have a tetrahedral arrangement, but two are lone pairs. The two O–H bonds are pushed closer together, reducing the bond angle to approximately 104.5°. By contrast, sulfur dioxide has a bent shape derived from a trigonal planar electron arrangement, so its bond angle is much closer to 120°.
These differences show why molecular geometry cannot always be predicted by counting bonded atoms alone. A complete prediction considers both bonding and nonbonding electron regions.
How molecular shape affects polarity and chemical behavior
Molecular shape influences more than appearance. It determines the directions in which bond dipoles point, affecting whether a molecule has an overall dipole moment.
A bond dipole results when electrons are shared unequally between two atoms. The more electronegative atom attracts the shared electrons more strongly, creating a separation of electrical charge. A molecule’s overall dipole depends on the combined effect of all its bond dipoles.
In a symmetrical linear molecule such as carbon dioxide, equal bond dipoles point in opposite directions and cancel. In symmetrical trigonal planar boron trifluoride, the three equal bond dipoles also cancel because they are evenly distributed around the central atom.
Water behaves differently. Its bent shape prevents its two O–H bond dipoles from canceling, leaving a net dipole. This contributes to water’s strong intermolecular attractions, including hydrogen bonding, which helps explain its relatively high boiling point compared with many similarly sized molecules.
Molecular shape can also influence how substances react. The positions of atoms and the distribution of electron density affect which parts of a molecule can approach a reaction partner and how chemical bonds can form or break. In larger molecules, geometry can help determine the fit between interacting molecules and the orientation needed for a reaction.
Shape alone does not determine every chemical property. Bond strength, electron distribution, molecular size, and surrounding conditions also matter. Nevertheless, geometry provides an essential starting point for understanding why molecules with different structures can behave differently.
How to identify a molecular shape
To identify whether a molecule is linear, bent, or trigonal planar, begin by drawing its Lewis structure. This representation shows which atoms are connected, the bonds between them, and the lone pairs of valence electrons.
Next, identify the central atom and count the regions of electron density around it. Each bond counts as one region, even if it is a double or triple bond, and each lone pair counts as one region.
Then use the arrangement of these regions to predict the electron geometry. Two regions generally favor a linear arrangement, three favor a trigonal planar arrangement, and four favor a tetrahedral arrangement.
Finally, determine the molecular geometry by considering the positions of the atoms while excluding lone pairs from the shape’s name. Two bonds around a central atom produce a linear shape if their directions are opposite, but a bent shape if they form an angle. Three bonds around a central atom produce a trigonal planar shape when the three bonded atoms occupy a flat triangular arrangement.
For example, carbon dioxide has two bonding regions around carbon and no lone pairs, so it is linear. Water has two bonding regions and two lone pairs around oxygen, so its electron geometry is tetrahedral but its molecular shape is bent. Boron trifluoride has three bonding regions and no lone pairs around boron, giving it a trigonal planar shape.
The central lesson is that molecular geometry describes the arrangement of atoms, while electron geometry describes the arrangement of all electron regions. Once that distinction is clear, linear, bent, and trigonal planar shapes become easier to recognize and understand.