Shapes of Molecules

Water and carbon dioxide are both triatomic molecules. The similarities go even further - they are both partially formed from oxygen and they both contain covalent bonds. However, their molecules are very different in shape. Whilst the atoms in carbon dioxide are held in a straight line, water is a bent molecule. To understand why this is the case, we need to take a look at electron pair repulsion, VSEPR.

Get started Sign up for free
Shapes of Molecules Shapes of Molecules

Create learning materials about Shapes of Molecules with our free learning app!

  • Instand access to millions of learning materials
  • Flashcards, notes, mock-exams and more
  • Everything you need to ace your exams
Create a free account

Millions of flashcards designed to help you ace your studies

Sign up for free

Convert documents into flashcards for free with AI!

Table of contents

    Valence shell electron pair repulsion theory, or VSEPR, is a model used in chemistry to predict the shape of molecules.

    If we break that term down a little, we can understand what it means.

    You should know that electrons tend to go around in pairs. This is because orbitals, which are regions of space where electrons can be found 95 percent of the time, can contain at most two electrons (check out Electron Shells, Subshells, and Orbitals for a refresher). Because electrons are charged particles, electron pairs will repel each other and try to be as far away from each other as possible. An atom’s outer shell of electrons is known as its valence shell. Because the valence electrons in a simple covalent molecule are the bonded electrons, electron pair repulsion determines the way in which the bonds are positioned. This dictates the shape of the molecule.

    VSEPR states that electron pairs all repel each other and will try to take up positions as far away from each other as possible, in order to minimise repulsion. It simply uses our knowledge of the behaviour of electrons to predict the shape of simple covalent compounds. Check out Covalent and Dative Bonding to remind yourself of how atoms share electrons in order to achieve stable electron configurations.

    How do you draw the shapes of molecules in 3D?

    Before we look at any examples of covalent structures, we need to learn how you can represent them. You might remember that we can draw covalent bonds as a line between two atoms. This gives a simple picture of molecules. However, if we want to better show a molecule’s 3D shape, we can use wedged and dotted lines.

    • Wedged lines show a bond coming out of the screen or page towards you.
    • Dotted or dashed lines show a bond going into the screen or page away from you.
    • Lone pairs of electrons are shown as dots.
    • Any standard straight lines simply show a planar bond.

    The methane molecule is a good example of this:

    Shapes of Molecules, methane, StudySmarterA methane molecule, CH4. The central wedged bond protrudes out of the screen whilst the right hand dashed bond extends backwards.

    The various shapes of molecules

    If all the pairs of valence electrons in an atom are bonded, they will all repel each other mutually. This results in bonds spaced equally far apart. The number of bonded electron pairs affects the shape of the molecule and the angle between the bonding pairs.

    Let’s take a look at some of the most common shapes. However, you should bear in mind that these rules only apply to molecules with no lone pairs of electrons. Lone pairs of electrons are unshared pairs that aren’t covalently bonded. We’ll explore their effect further later.


    If a molecule only has two bonded electron pairs (and no lone pairs), it forms a linear molecule. The simplest example is beryllium chloride, \(BeCl_2\) . Although beryllium is a metal, it can bond covalently to chlorine. Beryllium only has two electrons in its valence shell and so forms two bonds. The electron pairs will repel each other equally, resulting in an angle between the two bonds of 180°.

    Shapes of molecules, beryllium chloride linear, StudySmarterBeryllium chloride. Each of beryllium's valence electrons forms a covalent bond with a chlorine atom. The angle between the bonds around the central beryllium atom is 180°.StudySmarter Originals

    Trigonal planar

    Molecules with three bonded electron pairs are known as trigonal planar. This is because the bond angle between each bond is 120°, so the bonds lie flat on a plane. You could stack the molecules up one on top of the other like sheets of paper. Boron trifluoride is an example.

    Shapes of molecules, boron trifluoride trigonal planar, StudySmarterBoron trifluoride. The bonds are held apart at an angle of 120°.StudySmarter Originals


    Molecules with four bonded electron pairs and no lone pairs form a tetrahedral shape. This is a regular triangular-based pyramid. All the bond angles are 109.5°. For example, the carbon in methane \(CH_4\) has four valence electrons, and each electron is part of a pair bonded covalently to a hydrogen atom. It is a tetrahedral molecule.

    Shapes of molecules, methane tetrahedral, StudySmarterMethane. The angle between each bond around the central atom is 109.5°.StudySmarter Originals

    Trigonal bipyramidal

    Molecules with five bonded electron pairs form a trigonal bipyramid. This shape is similar to a trigonal planar molecule but with two further bonds held at 90° extending above and below the plane. Phosphorus(V) pentachloride is a good example.

    Shapes of molecules, phosphorus(V) pentachloride trigonal bipyramid, StudySmarterPhosphorus(V) pentachloride. Three planar bonds have angles of 120° between them, whilst two further bonds are held at right angles to the plane.StudySmarter Originals


    If a molecule has six bonding pairs around a central atom, it forms an octahedral structure. All of the bonds are at right angles to each other, as shown in sulfur hexafluoride.

    Shapes of molecules, sulfur hexafluoride octahedral, StudySmarterSulfur hexafluoride has six bonded electron pairs. All bond angles are 90°. StudySmarter Originals

    Lone pairs of electrons

    All of our above examples use molecules that don’t have any lone pairs of electrons. All their valence electrons are bonded. But what happens if a molecule does have a lone pair? Let’s take a molecule with four electron pairs as an example.

    We now know that if all of the electrons are part of bonding pairs, the molecule will be tetrahedral and have bond angles of 109.5°. However, if one of the electron pairs is in fact a lone pair, the bond angles are reduced to 107°. This is because lone pairs repel each other more strongly than shared pairs, squeezing the bonds together. Each lone electron pair in a molecule with eight valence electrons reduces the bond angle by 2.5°, so a molecule with two bonding pairs and two lone pairs will have a bond angle of 104.5°. The following table shows the relative strength of repulsion between combinations of bonded and lone pairs of electrons.

    Shapes of molecules, lone electron pair repulsion, StudySmarterA table comparing the strength of repulsion between bonded and lone pairs of electrons. StudySmarter Originals

    Let’s now look at the shapes formed by molecules with lone pairs.


    A molecule with three bonded electron pairs and one lone electron pair around a central atom has an angle of 107° between each bond. An example is ammonia, \(NH_3\) . The nitrogen atom contains five valence electrons. Three are covalently bonded to hydrogen atoms and the remaining two form a lone pair. This lone pair repels the bonding pairs more strongly than the bonding pairs repel each other, reducing the bond angle and forming a pyramidal molecule.

    Shapes of molecules, ammonia pyramidal, StudySmarterAn ammonia molecule. Compared to a tetrahedral molecule with no lone pairs, the bond angle is reduced by 2.5°. StudySmarter Originals


    A molecule with two lone pairs and two bonding pairs has its bond angle reduced even further to 104.5°. This forms a v-shaped molecule, such as water, \(H_2O\) .

    Shapes of molecules, water v shaped, StudySmarterA v-shaped water molecule. StudySmarter Originals

    The following diagram summarises the different shapes of molecules.

    Shapes of molecules, shapes of molecules table summary, StudySmarterA table summarising the shapes of different molecules. StudySmarter Originals

    Examples of the shapes of molecules

    Let’s go back to our original molecules, water and carbon dioxide. We’ve already discovered that water has a v-shaped structure due to the effect of its lone electron pairs on the bonding pairs. But what sort of shape does carbon dioxide have?

    By drawing a dot and cross diagram we can see that carbon dioxide, \(CO_2\) , has two double bonds. These double bonds can be thought of as single units when it comes to shape. Like single bond electron pairs, these groups of four electrons will want to be as far apart from each other as possible. This forms a linear molecule with a bond angle of 180°.

    structure of Carbon dioxide molecule, StudySmarterCarbon dioxide. Although it contains four bonding pairs of electrons, the pairs are arranged as two double bonds. Each double bond is considered as a single unit, so the molecule is linear.

    Another example is xenon tetrafluoride, \(XeF_4\) . Xenon contains eight electrons in its valence shell. Four form bonds with fluorine atoms and four remain as two lone pairs. This forms what is known as a square planar arrangement, with the lone pairs at 180° to each other, and the angle between the bonding pairs at 90°. Note its similarity to an octahedral arrangement.

    Shapes of molecules, xenon tetrafluoride square planar, StudySmarterXenon tetrafluoride. The lone pairs of electrons are positioned above and below the plane.StudySmarter Originals

    Shapes of Molecules - Key takeaways

    • VSEPR, also known as valence shell electron pair repulsion theory, states that electron pairs repel each other and will try to take up positions as far away from each other as possible, in order to minimise repulsion. This influences the shapes of molecules.
    • You can use straight lines to represent covalent bonds. Wedged lines show a bond protruding out of the page and dashed or dotted lines show a bond extending backwards.
    • Lone pairs of electrons repel each other more strongly than bonding pairs. Each lone pair reduces the bond angle by 2.5° in molecules with four electron pairs.
    • Common molecule shapes with no lone pairs of electrons include linear, trigonal planar, tetrahedral, trigonal bipyramidal, and octahedral.
    • Common molecule shapes with lone pairs of electrons include pyramidal and v-shaped.
    Frequently Asked Questions about Shapes of Molecules

    What is the shape of a water molecule?

    Water molecules are v-shaped.

    What is the shape of the DNA molecule called?

    DNA forms a double helix shape.

    What is the shape of the methane molecule?

    Methane molecules are tetrahedral in shape.

    What is the shape of the xenon tetrafluoride molecule?

    Xenon tetrafluoride is square planar in shape.

    How do you work out the shape of a molecule?

    To work out the shape of a molecule, identify how many lone and bonding pairs of electrons it has. This dictates its shape. For example, the oxygen atom in a water molecule has two lone pairs and two bonding pairs. This gives it a v-shaped structure.

    How can the shape of a molecule affect its polarity?

    Molecules with polar bonds are often polar molecules. However, if the molecule is symmetrical, the charges of the polar bonds cancel out and the molecule is non-polar overall.

    Test your knowledge with multiple choice flashcards

    What is a valence electron?

    Electron domain geometry is the same as molecular shape geometry.

    What is the angle created by the bonds in H2O?


    Discover learning materials with the free StudySmarter app

    Sign up for free
    About StudySmarter

    StudySmarter is a globally recognized educational technology company, offering a holistic learning platform designed for students of all ages and educational levels. Our platform provides learning support for a wide range of subjects, including STEM, Social Sciences, and Languages and also helps students to successfully master various tests and exams worldwide, such as GCSE, A Level, SAT, ACT, Abitur, and more. We offer an extensive library of learning materials, including interactive flashcards, comprehensive textbook solutions, and detailed explanations. The cutting-edge technology and tools we provide help students create their own learning materials. StudySmarter’s content is not only expert-verified but also regularly updated to ensure accuracy and relevance.

    Learn more
    StudySmarter Editorial Team

    Team Chemistry Teachers

    • 9 minutes reading time
    • Checked by StudySmarter Editorial Team
    Save Explanation Save Explanation

    Study anywhere. Anytime.Across all devices.

    Sign-up for free

    Sign up to highlight and take notes. It’s 100% free.

    Join over 22 million students in learning with our StudySmarter App

    The first learning app that truly has everything you need to ace your exams in one place

    • Flashcards & Quizzes
    • AI Study Assistant
    • Study Planner
    • Mock-Exams
    • Smart Note-Taking
    Join over 22 million students in learning with our StudySmarter App
    Sign up with Email

    Get unlimited access with a free StudySmarter account.

    • Instant access to millions of learning materials.
    • Flashcards, notes, mock-exams, AI tools and more.
    • Everything you need to ace your exams.
    Second Popup Banner