10.2: VSEPR Theory - The Five Basic Shapes (2024)

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    Learning Objectives
    • To use the VSEPR model to predict molecular geometries.
    • To predict whether a molecule has a dipole moment.

    The Lewis electron-pair approach can be used to predict the number and types of bonds between the atoms in a substance, and it indicates which atoms have lone pairs of electrons. This approach gives no information about the actual arrangement of atoms in space, however. We continue our discussion of structure and bonding by introducing the valence-shell electron-pair repulsion (VSEPR) model (pronounced “vesper”), which can be used to predict the shapes of many molecules and polyatomic ions. Keep in mind, however, that the VSEPR model, like any model, is a limited representation of reality; the model provides no information about bond lengths or the presence of multiple bonds.

    The VSEPR Model

    The VSEPR model can predict the structure of nearly any molecule or polyatomic ion in which the central atom is a nonmetal, as well as the structures of many molecules and polyatomic ions with a central metal atom. The premise of the VSEPR theory is that electron pairs located in bonds and lone pairs repel each other and will therefore adopt the geometry that places electron pairs as far apart from each other as possible. This theory is very simplistic and does not account for the subtleties of orbital interactions that influence molecular shapes; however, the simple VSEPR counting procedure accurately predicts the three-dimensional structures of a large number of compounds, which cannot be predicted using the Lewis electron-pair approach.

    10.2: VSEPR Theory - The Five Basic Shapes (1)

    We can use the VSEPR model to predict the geometry of most polyatomic molecules and ions by focusing only on the number of electron pairs around the central atom, ignoring all other valence electrons present. According to this model, valence electrons in the Lewis structure form groups, which may consist of a single bond, a double bond, a triple bond, a lone pair of electrons, or even a single unpaired electron, which in the VSEPR model is counted as a lone pair. Because electrons repel each other electrostatically, the most stable arrangement of electron groups (i.e., the one with the lowest energy) is the one that minimizes repulsions. Groups are positioned around the central atom in a way that produces the molecular structure with the lowest energy, as illustrated in Figures \(\PageIndex{1}\) and \(\PageIndex{2}\).

    10.2: VSEPR Theory - The Five Basic Shapes (2)

    In the VSEPR model, the molecule or polyatomic ion is given an AXmEn designation, where A is the central atom, X is a bonded atom, E is a nonbonding valence electron group (usually a lone pair of electrons), and m and n are integers. Each group around the central atom is designated as a bonding pair (BP) or lone (nonbonding) pair (LP). From the BP and LP interactions we can predict both the relative positions of the atoms and the angles between the bonds, called the bond angles. Using this information, we can describe the molecular geometry, the arrangement of the bonded atoms in a molecule or polyatomic ion.

    VESPR Produce to predict Molecular geometry

    This VESPR procedure is summarized as follows:

    1. Draw the Lewis electron structure of the molecule or polyatomic ion.
    2. Determine the electron group arrangement around the central atom that minimizes repulsions.
    3. Assign an AXmEn designation; then identify the LP–LP, LP–BP, or BP–BP interactions and predict deviations from ideal bond angles.
    4. Describe the molecular geometry.

      We will illustrate the use of this procedure with several examples, beginning with atoms with two electron groups. In our discussion we will refer to Figure \(\PageIndex{2}\) and Figure \(\PageIndex{3}\), which summarize the common molecular geometries and idealized bond angles of molecules and ions with two to six electron groups.

      10.2: VSEPR Theory - The Five Basic Shapes (3)

      Two Electron Groups

      Our first example is a molecule with two bonded atoms and no lone pairs of electrons, \(BeH_2\).

      AX2 Molecules: BeH2

      1. The central atom, beryllium, contributes two valence electrons, and each hydrogen atom contributes one. The Lewis electron structure is

      10.2: VSEPR Theory - The Five Basic Shapes (4)

      3. Both groups around the central atom are bonding pairs (BP). Thus BeH2 is designated as AX2.

      4. From Figure \(\PageIndex{3}\) we see that with two bonding pairs, the molecular geometry that minimizes repulsions in BeH2 is linear.

      AX2 Molecules: CO2

      1. The central atom, carbon, contributes four valence electrons, and each oxygen atom contributes six. The Lewis electron structure is

      10.2: VSEPR Theory - The Five Basic Shapes (5)

      2. The carbon atom forms two double bonds. Each double bond is a group, so there are two electron groups around the central atom. Like BeH2, the arrangement that minimizes repulsions places the groups 180° apart.

      3. Once again, both groups around the central atom are bonding pairs (BP), so CO2 is designated as AX2.

      4. VSEPR only recognizes groups around the central atom. Thus the lone pairs on the oxygen atoms do not influence the molecular geometry. With two bonding pairs on the central atom and no lone pairs, the molecular geometry of CO2 is linear (Figure \(\PageIndex{3}\)). The structure of \(\ce{CO2}\) is shown in Figure \(\PageIndex{1}\).

      Three Electron Groups

      AX3 Molecules: BCl3

      1. The central atom, boron, contributes three valence electrons, and each chlorine atom contributes seven valence electrons. The Lewis electron structure is

      10.2: VSEPR Theory - The Five Basic Shapes (6)

      3. All electron groups are bonding pairs (BP), so the structure is designated as AX3.

      4. From Figure \(\PageIndex{3}\) we see that with three bonding pairs around the central atom, the molecular geometry of BCl3 is trigonal planar, as shown in Figure \(\PageIndex{2}\).

      AX3 Molecules: CO32

      1. The central atom, carbon, has four valence electrons, and each oxygen atom has six valence electrons. As you learned previously, the Lewis electron structure of one of three resonance forms is represented as

      10.2: VSEPR Theory - The Five Basic Shapes (7)

      3. All electron groups are bonding pairs (BP). With three bonding groups around the central atom, the structure is designated as AX3.

      4. We see from Figure \(\PageIndex{3}\) that the molecular geometry of CO32 is trigonal planar with bond angles of 120°.

      10.2: VSEPR Theory - The Five Basic Shapes (8)

      In our next example we encounter the effects of lone pairs and multiple bonds on molecular geometry for the first time.

      AX2E Molecules: SO2

      1. The central atom, sulfur, has 6 valence electrons, as does each oxygen atom. With 18 valence electrons, the Lewis electron structure is shown below.

      10.2: VSEPR Theory - The Five Basic Shapes (9)

      3. There are two bonding pairs and one lone pair, so the structure is designated as AX2E. This designation has a total of three electron pairs, two X and one E. Because a lone pair is not shared by two nuclei, it occupies more space near the central atom than a bonding pair (Figure \(\PageIndex{4}\)). Thus bonding pairs and lone pairs repel each other electrostatically in the order BP–BP < LP–BP < LP–LP. In SO2, we have one BP–BP interaction and two LP–BP interactions.

      4. The molecular geometry is described only by the positions of the nuclei, not by the positions of the lone pairs. Thus with two nuclei and one lone pair the shape is bent, or V shaped, which can be viewed as a trigonal planar arrangement with a missing vertex (Figures \(\PageIndex{2}\) and \(\PageIndex{3}\)). The O-S-O bond angle is expected to be less than 120° because of the extra space taken up by the lone pair.

      10.2: VSEPR Theory - The Five Basic Shapes (10)

      As with SO2, this composite model of electron distribution and negative electrostatic potential in ammonia shows that a lone pair of electrons occupies a larger region of space around the nitrogen atom than does a bonding pair of electrons that is shared with a hydrogen atom.

      Like lone pairs of electrons, multiple bonds occupy more space around the central atom than a single bond, which can cause other bond angles to be somewhat smaller than expected. This is because a multiple bond has a higher electron density than a single bond, so its electrons occupy more space than those of a single bond. For example, in a molecule such as CH2O (AX3), whose structure is shown below, the double bond repels the single bonds more strongly than the single bonds repel each other. This causes a deviation from ideal geometry (an H–C–H bond angle of 116.5° rather than 120°).

      10.2: VSEPR Theory - The Five Basic Shapes (11)

      Four Electron Groups

      One of the limitations of Lewis structures is that they depict molecules and ions in only two dimensions. With four electron groups, we must learn to show molecules and ions in three dimensions.

      AX4 Molecules: CH4

      1. The central atom, carbon, contributes four valence electrons, and each hydrogen atom has one valence electron, so the full Lewis electron structure is

      10.2: VSEPR Theory - The Five Basic Shapes (12)

      2. There are four electron groups around the central atom. As shown in Figure \(\PageIndex{2}\), repulsions are minimized by placing the groups in the corners of a tetrahedron with bond angles of 109.5°.

      3. All electron groups are bonding pairs, so the structure is designated as AX4.

      4. With four bonding pairs, the molecular geometry of methane is tetrahedral (Figure \(\PageIndex{3}\)).

      10.2: VSEPR Theory - The Five Basic Shapes (13)

      AX3E Molecules: NH3

      1. In ammonia, the central atom, nitrogen, has five valence electrons and each hydrogen donates one valence electron, producing the Lewis electron structure

      10.2: VSEPR Theory - The Five Basic Shapes (14)

      2. There are four electron groups around nitrogen, three bonding pairs and one lone pair. Repulsions are minimized by directing each hydrogen atom and the lone pair to the corners of a tetrahedron.

      3. With three bonding pairs and one lone pair, the structure is designated as AX3E. This designation has a total of four electron pairs, three X and one E. We expect the LP–BP interactions to cause the bonding pair angles to deviate significantly from the angles of a perfect tetrahedron.

      4. There are three nuclei and one lone pair, so the molecular geometry is trigonal pyramidal. In essence, this is a tetrahedron with a vertex missing (Figure \(\PageIndex{3}\)). However, the H–N–H bond angles are less than the ideal angle of 109.5° because of LP–BP repulsions (Figure \(\PageIndex{3}\) and Figure \(\PageIndex{4}\)).

      AX2E2 Molecules: H2O

      1. Oxygen has six valence electrons and each hydrogen has one valence electron, producing the Lewis electron structure

      10.2: VSEPR Theory - The Five Basic Shapes (15)

      3. With two bonding pairs and two lone pairs, the structure is designated as AX2E2 with a total of four electron pairs. Due to LP–LP, LP–BP, and BP–BP interactions, we expect a significant deviation from idealized tetrahedral angles.

      4. With two hydrogen atoms and two lone pairs of electrons, the structure has significant lone pair interactions. There are two nuclei about the central atom, so the molecular shape is bent, or V shaped, with an H–O–H angle that is even less than the H–N–H angles in NH3, as we would expect because of the presence of two lone pairs of electrons on the central atom rather than one. This molecular shape is essentially a tetrahedron with two missing vertices.

      10.2: VSEPR Theory - The Five Basic Shapes (16)

      Five Electron Groups

      In previous examples it did not matter where we placed the electron groups because all positions were equivalent. In some cases, however, the positions are not equivalent. We encounter this situation for the first time with five electron groups.

      AX5 Molecules: PCl5

      1. Phosphorus has five valence electrons and each chlorine has seven valence electrons, so the Lewis electron structure of PCl5 is

      10.2: VSEPR Theory - The Five Basic Shapes (17)

      3. All electron groups are bonding pairs, so the structure is designated as AX5. There are no lone pair interactions.

      4. The molecular geometry of PCl5 is trigonal bipyramidal, as shown in Figure \(\PageIndex{3}\). The molecule has three atoms in a plane in equatorial positions and two atoms above and below the plane in axial positions. The three equatorial positions are separated by 120° from one another, and the two axial positions are at 90° to the equatorial plane. The axial and equatorial positions are not chemically equivalent, as we will see in our next example.

      10.2: VSEPR Theory - The Five Basic Shapes (18)

      AX4E Molecules: SF4

      1. The sulfur atom has six valence electrons and each fluorine has seven valence electrons, so the Lewis electron structure is

      10.2: VSEPR Theory - The Five Basic Shapes (19)

      With an expanded valence, this species is an exception to the octet rule.

      2. There are five groups around sulfur, four bonding pairs and one lone pair. With five electron groups, the lowest energy arrangement is a trigonal bipyramid, as shown in Figure \(\PageIndex{2}\).

      3. We designate SF4 as AX4E; it has a total of five electron pairs. However, because the axial and equatorial positions are not chemically equivalent, where do we place the lone pair? If we place the lone pair in the axial position, we have three LP–BP repulsions at 90°. If we place it in the equatorial position, we have two 90° LP–BP repulsions at 90°. With fewer 90° LP–BP repulsions, we can predict that the structure with the lone pair of electrons in the equatorial position is more stable than the one with the lone pair in the axial position. We also expect a deviation from ideal geometry because a lone pair of electrons occupies more space than a bonding pair.

      10.2: VSEPR Theory - The Five Basic Shapes (20)

      At 90°, the two electron pairs share a relatively large region of space, which leads to strong repulsive electron–electron interactions.

      4. With four nuclei and one lone pair of electrons, the molecular structure is based on a trigonal bipyramid with a missing equatorial vertex; it is described as a seesaw. The Faxial–S–Faxial angle is 173° rather than 180° because of the lone pair of electrons in the equatorial plane.

      10.2: VSEPR Theory - The Five Basic Shapes (21)

      AX3E2 Molecules: BrF3

      1. The bromine atom has seven valence electrons, and each fluorine has seven valence electrons, so the Lewis electron structure is

      10.2: VSEPR Theory - The Five Basic Shapes (22)

      Once again, we have a compound that is an exception to the octet rule.

      2. There are five groups around the central atom, three bonding pairs and two lone pairs. We again direct the groups toward the vertices of a trigonal bipyramid.

      3. With three bonding pairs and two lone pairs, the structural designation is AX3E2 with a total of five electron pairs. Because the axial and equatorial positions are not equivalent, we must decide how to arrange the groups to minimize repulsions. If we place both lone pairs in the axial positions, we have six LP–BP repulsions at 90°. If both are in the equatorial positions, we have four LP–BP repulsions at 90°. If one lone pair is axial and the other equatorial, we have one LP–LP repulsion at 90° and three LP–BP repulsions at 90°:

      10.2: VSEPR Theory - The Five Basic Shapes (23)

      Structure (c) can be eliminated because it has a LP–LP interaction at 90°. Structure (b), with fewer LP–BP repulsions at 90° than (a), is lower in energy. However, we predict a deviation in bond angles because of the presence of the two lone pairs of electrons.

      4. The three nuclei in BrF3 determine its molecular structure, which is described as T shaped. This is essentially a trigonal bipyramid that is missing two equatorial vertices. The Faxial–Br–Faxial angle is 172°, less than 180° because of LP–BP repulsions (Figure \(\PageIndex{2}\).1).

      Because lone pairs occupy more space around the central atom than bonding pairs, electrostatic repulsions are more important for lone pairs than for bonding pairs.

      AX2E3 Molecules: I3

      1. Each iodine atom contributes seven electrons and the negative charge one, so the Lewis electron structure is

      10.2: VSEPR Theory - The Five Basic Shapes (24)

      2. There are five electron groups about the central atom in I3, two bonding pairs and three lone pairs. To minimize repulsions, the groups are directed to the corners of a trigonal bipyramid.

      3. With two bonding pairs and three lone pairs, I3 has a total of five electron pairs and is designated as AX2E3. We must now decide how to arrange the lone pairs of electrons in a trigonal bipyramid in a way that minimizes repulsions. Placing them in the axial positions eliminates 90° LP–LP repulsions and minimizes the number of 90° LP–BP repulsions.

      10.2: VSEPR Theory - The Five Basic Shapes (25)

      The three lone pairs of electrons have equivalent interactions with the three iodine atoms, so we do not expect any deviations in bonding angles.

      4. With three nuclei and three lone pairs of electrons, the molecular geometry of I3 is linear. This can be described as a trigonal bipyramid with three equatorial vertices missing. The ion has an I–I–I angle of 180°, as expected.

      10.2: VSEPR Theory - The Five Basic Shapes (26)

      Six Electron Groups

      Six electron groups form an octahedron, a polyhedron made of identical equilateral triangles and six identical vertices (Figure \(\PageIndex{2}\).)

      AX6 Molecules: SF6

      1. The central atom, sulfur, contributes six valence electrons, and each fluorine atom has seven valence electrons, so the Lewis electron structure is

      10.2: VSEPR Theory - The Five Basic Shapes (27)

      With an expanded valence, this species is an exception to the octet rule.

      2. There are six electron groups around the central atom, each a bonding pair. We see from Figure \(\PageIndex{2}\) that the geometry that minimizes repulsions is octahedral.

      3. With only bonding pairs, SF6 is designated as AX6. All positions are chemically equivalent, so all electronic interactions are equivalent.

      4. There are six nuclei, so the molecular geometry of SF6 is octahedral.

      10.2: VSEPR Theory - The Five Basic Shapes (28)

      AX5E Molecules: BrF5

      1. The central atom, bromine, has seven valence electrons, as does each fluorine, so the Lewis electron structure is

      10.2: VSEPR Theory - The Five Basic Shapes (29)

      With its expanded valence, this species is an exception to the octet rule.

      2. There are six electron groups around the Br, five bonding pairs and one lone pair. Placing five F atoms around Br while minimizing BP–BP and LP–BP repulsions gives the following structure:

      10.2: VSEPR Theory - The Five Basic Shapes (30)

      3. With five bonding pairs and one lone pair, BrF5 is designated as AX5E; it has a total of six electron pairs. The BrF5 structure has four fluorine atoms in a plane in an equatorial position and one fluorine atom and the lone pair of electrons in the axial positions. We expect all Faxial–Br–Fequatorial angles to be less than 90° because of the lone pair of electrons, which occupies more space than the bonding electron pairs.

      4. With five nuclei surrounding the central atom, the molecular structure is based on an octahedron with a vertex missing. This molecular structure is square pyramidal. The Faxial–B–Fequatorial angles are 85.1°, less than 90° because of LP–BP repulsions.

      10.2: VSEPR Theory - The Five Basic Shapes (31)

      AX4E2 Molecules: ICl4

      1. The central atom, iodine, contributes seven electrons. Each chlorine contributes seven, and there is a single negative charge. The Lewis electron structure is

      10.2: VSEPR Theory - The Five Basic Shapes (32)

      2. There are six electron groups around the central atom, four bonding pairs and two lone pairs. The structure that minimizes LP–LP, LP–BP, and BP–BP repulsions is

      10.2: VSEPR Theory - The Five Basic Shapes (33)

      3. ICl4 is designated as AX4E2 and has a total of six electron pairs. Although there are lone pairs of electrons, with four bonding electron pairs in the equatorial plane and the lone pairs of electrons in the axial positions, all LP–BP repulsions are the same. Therefore, we do not expect any deviation in the Cl–I–Cl bond angles.

      4. With five nuclei, the ICl4− ion forms a molecular structure that is square planar, an octahedron with two opposite vertices missing.

      10.2: VSEPR Theory - The Five Basic Shapes (34)

      The relationship between the number of electron groups around a central atom, the number of lone pairs of electrons, and the molecular geometry is summarized in Figure \(\PageIndex{6}\).

      10.2: VSEPR Theory - The Five Basic Shapes (35)

      Figure \(\PageIndex{6}\): Overview of Molecular Geometries

      Example \(\PageIndex{1}\)

      Using the VSEPR model, predict the molecular geometry of each molecule or ion.

      1. PF5 (phosphorus pentafluoride, a catalyst used in certain organic reactions)
      2. H3O+ (hydronium ion)

      Given: two chemical species

      Asked for: molecular geometry

      Strategy:
      1. Draw the Lewis electron structure of the molecule or polyatomic ion.
      2. Determine the electron group arrangement around the central atom that minimizes repulsions.
      3. Assign an AXmEn designation; then identify the LP–LP, LP–BP, or BP–BP interactions and predict deviations in bond angles.
      4. Describe the molecular geometry.
      Solution:
      1. A The central atom, P, has five valence electrons and each fluorine has seven valence electrons, so the Lewis structure of PF5 is
        10.2: VSEPR Theory - The Five Basic Shapes (36)

        C All electron groups are bonding pairs, so PF5 is designated as AX5. Notice that this gives a total of five electron pairs. With no lone pair repulsions, we do not expect any bond angles to deviate from the ideal.

        D The PF5 molecule has five nuclei and no lone pairs of electrons, so its molecular geometry is trigonal bipyramidal.

        10.2: VSEPR Theory - The Five Basic Shapes (37)

      2. A The central atom, O, has six valence electrons, and each H atom contributes one valence electron. Subtracting one electron for the positive charge gives a total of eight valence electrons, so the Lewis electron structure is
        10.2: VSEPR Theory - The Five Basic Shapes (38)

        B There are four electron groups around oxygen, three bonding pairs and one lone pair. Like NH3, repulsions are minimized by directing each hydrogen atom and the lone pair to the corners of a tetrahedron.

        C With three bonding pairs and one lone pair, the structure is designated as AX3E and has a total of four electron pairs (three X and one E). We expect the LP–BP interactions to cause the bonding pair angles to deviate significantly from the angles of a perfect tetrahedron.

        D There are three nuclei and one lone pair, so the molecular geometry is trigonal pyramidal, in essence a tetrahedron missing a vertex. However, the H–O–H bond angles are less than the ideal angle of 109.5° because of LP–BP repulsions:

        10.2: VSEPR Theory - The Five Basic Shapes (39)

      Exercise \(\PageIndex{1}\)

      Using the VSEPR model, predict the molecular geometry of each molecule or ion.

      1. XeO3
      2. PF6
      3. NO2+
      Answer a

      trigonal pyramidal

      Answer b

      octahedral

      Answer c

      linear

      Example \(\PageIndex{2}\)

      Predict the molecular geometry of each molecule.

      1. XeF2
      2. SnCl2

      Given: two chemical compounds

      Asked for: molecular geometry

      Strategy:

      Use the strategy given in Example\(\PageIndex{1}\).

      Solution:
      1. A Xenon contributes eight electrons and each fluorine seven valence electrons, so the Lewis electron structure is

        10.2: VSEPR Theory - The Five Basic Shapes (40)

        B There are five electron groups around the central atom, two bonding pairs and three lone pairs. Repulsions are minimized by placing the groups in the corners of a trigonal bipyramid.

        C From B, XeF2 is designated as AX2E3 and has a total of five electron pairs (two X and three E). With three lone pairs about the central atom, we can arrange the two F atoms in three possible ways: both F atoms can be axial, one can be axial and one equatorial, or both can be equatorial:

        10.2: VSEPR Theory - The Five Basic Shapes (41)

        The structure with the lowest energy is the one that minimizes LP–LP repulsions. Both (b) and (c) have two 90° LP–LP interactions, whereas structure (a) has none. Thus both F atoms are in the axial positions, like the two iodine atoms around the central iodine in I3. All LP–BP interactions are equivalent, so we do not expect a deviation from an ideal 180° in the F–Xe–F bond angle.

        D With two nuclei about the central atom, the molecular geometry of XeF2 is linear. It is a trigonal bipyramid with three missing equatorial vertices.

      2. A The tin atom donates 4 valence electrons and each chlorine atom donates 7 valence electrons. With 18 valence electrons, the Lewis electron structure is
        10.2: VSEPR Theory - The Five Basic Shapes (42)

        B There are three electron groups around the central atom, two bonding groups and one lone pair of electrons. To minimize repulsions the three groups are initially placed at 120° angles from each other.

        C From B we designate SnCl2 as AX2E. It has a total of three electron pairs, two X and one E. Because the lone pair of electrons occupies more space than the bonding pairs, we expect a decrease in the Cl–Sn–Cl bond angle due to increased LP–BP repulsions.

        D With two nuclei around the central atom and one lone pair of electrons, the molecular geometry of SnCl2 is bent, like SO2, but with a Cl–Sn–Cl bond angle of 95°. The molecular geometry can be described as a trigonal planar arrangement with one vertex missing.

      Exercise \(\PageIndex{2}\)

      Predict the molecular geometry of each molecule.

      1. SO3
      2. XeF4
      Answer a

      trigonal planar

      Answer b

      square planar

      Molecules with No Single Central Atom

      The VSEPR model can be used to predict the structure of somewhat more complex molecules with no single central atom by treating them as linked AXmEn fragments. We will demonstrate with methyl isocyanate (CH3–N=C=O), a volatile and highly toxic molecule that is used to produce the pesticide Sevin. In 1984, large quantities of Sevin were accidentally released in Bhopal, India, when water leaked into storage tanks. The resulting highly exothermic reaction caused a rapid increase in pressure that ruptured the tanks, releasing large amounts of methyl isocyanate that killed approximately 3800 people and wholly or partially disabled about 50,000 others. In addition, there was significant damage to livestock and crops.

      We can treat methyl isocyanate as linked AXmEn fragments beginning with the carbon atom at the left, which is connected to three H atoms and one N atom by single bonds. The four bonds around carbon mean that it must be surrounded by four bonding electron pairs in a configuration similar to AX4. We can therefore predict the CH3–N portion of the molecule to be roughly tetrahedral, similar to methane:

      10.2: VSEPR Theory - The Five Basic Shapes (43)

      The nitrogen atom is connected to one carbon by a single bond and to the other carbon by a double bond, producing a total of three bonds, C–N=C. For nitrogen to have an octet of electrons, it must also have a lone pair:

      10.2: VSEPR Theory - The Five Basic Shapes (44)

      Because multiple bonds are not shown in the VSEPR model, the nitrogen is effectively surrounded by three electron pairs. Thus according to the VSEPR model, the C–N=C fragment should be bent with an angle less than 120°.

      The carbon in the –N=C=O fragment is doubly bonded to both nitrogen and oxygen, which in the VSEPR model gives carbon a total of two electron pairs. The N=C=O angle should therefore be 180°, or linear. The three fragments combine to give the following structure:

      10.2: VSEPR Theory - The Five Basic Shapes (45)
      10.2: VSEPR Theory - The Five Basic Shapes (46)

      Certain patterns are seen in the structures of moderately complex molecules. For example, carbon atoms with four bonds (such as the carbon on the left in methyl isocyanate) are generally tetrahedral. Similarly, the carbon atom on the right has two double bonds that are similar to those in CO2, so its geometry, like that of CO2, is linear. Recognizing similarities to simpler molecules will help you predict the molecular geometries of more complex molecules.

      Example \(\PageIndex{3}\)

      Use the VSEPR model to predict the molecular geometry of propyne (H3C–C≡CH), a gas with some anesthetic properties.

      Given: chemical compound

      Asked for: molecular geometry

      Strategy:

      Count the number of electron groups around each carbon, recognizing that in the VSEPR model, a multiple bond counts as a single group. Use Figure \(\PageIndex{3}\) to determine the molecular geometry around each carbon atom and then deduce the structure of the molecule as a whole.

      Solution:

      Because the carbon atom on the left is bonded to four other atoms, we know that it is approximately tetrahedral. The next two carbon atoms share a triple bond, and each has an additional single bond. Because a multiple bond is counted as a single bond in the VSEPR model, each carbon atom behaves as if it had two electron groups. This means that both of these carbons are linear, with C–C≡C and C≡C–H angles of 180°.

      Exercise \(\PageIndex{3}\)

      Predict the geometry of allene (H2C=C=CH2), a compound with narcotic properties that is used to make more complex organic molecules.

      Answer

      The terminal carbon atoms are trigonal planar, the central carbon is linear, and the C–C–C angle is 180°.

      Molecular Dipole Moments

      You previously learned how to calculate the dipole moments of simple diatomic molecules. In more complex molecules with polar covalent bonds, the three-dimensional geometry and the compound’s symmetry determine whether there is a net dipole moment. Mathematically, dipole moments are vectors; they possess both a magnitude and a direction. The dipole moment of a molecule is therefore the vector sum of the dipole moments of the individual bonds in the molecule. If the individual bond dipole moments cancel one another, there is no net dipole moment. Such is the case for CO2, a linear molecule (Figure \(\PageIndex{8a}\)). Each C–O bond in CO2 is polar, yet experiments show that the CO2 molecule has no dipole moment. Because the two C–O bond dipoles in CO2 are equal in magnitude and oriented at 180° to each other, they cancel. As a result, the CO2 molecule has no net dipole moment even though it has a substantial separation of charge. In contrast, the H2O molecule is not linear (Figure \(\PageIndex{8b}\)); it is bent in three-dimensional space, so the dipole moments do not cancel each other. Thus a molecule such as H2O has a net dipole moment. We expect the concentration of negative charge to be on the oxygen, the more electronegative atom, and positive charge on the two hydrogens. This charge polarization allows H2O to hydrogen-bond to other polarized or charged species, including other water molecules.

      10.2: VSEPR Theory - The Five Basic Shapes (47)

      Other examples of molecules with polar bonds are shown in Figure \(\PageIndex{9}\). In molecular geometries that are highly symmetrical (most notably tetrahedral and square planar, trigonal bipyramidal, and octahedral), individual bond dipole moments completely cancel, and there is no net dipole moment. Although a molecule like CHCl3 is best described as tetrahedral, the atoms bonded to carbon are not identical. Consequently, the bond dipole moments cannot cancel one another, and the molecule has a dipole moment. Due to the arrangement of the bonds in molecules that have V-shaped, trigonal pyramidal, seesaw, T-shaped, and square pyramidal geometries, the bond dipole moments cannot cancel one another. Consequently, molecules with these geometries always have a nonzero dipole moment. Molecules with asymmetrical charge distributions have a net dipole moment.

      10.2: VSEPR Theory - The Five Basic Shapes (48)

      Example \(\PageIndex{4}\)

      Which molecule(s) has a net dipole moment?

      1. \(\ce{H2S}\)
      2. \(\ce{NHF2}\)
      3. \(\ce{BF3}\)

      Given: three chemical compounds

      Asked for: net dipole moment

      Strategy:

      For each three-dimensional molecular geometry, predict whether the bond dipoles cancel. If they do not, then the molecule has a net dipole moment.

      Solution:
      1. The total number of electrons around the central atom, S, is eight, which gives four electron pairs. Two of these electron pairs are bonding pairs and two are lone pairs, so the molecular geometry of \(\ce{H2S}\) is bent (Figure \(\PageIndex{6}\)). The bond dipoles cannot cancel one another, so the molecule has a net dipole moment.

        10.2: VSEPR Theory - The Five Basic Shapes (49)

      2. Difluoroamine has a trigonal pyramidal molecular geometry. Because there is one hydrogen and two fluorines, and because of the lone pair of electrons on nitrogen, the molecule is not symmetrical, and the bond dipoles of NHF2 cannot cancel one another. This means that NHF2 has a net dipole moment. We expect polarization from the two fluorine atoms, the most electronegative atoms in the periodic table, to have a greater affect on the net dipole moment than polarization from the lone pair of electrons on nitrogen.

        10.2: VSEPR Theory - The Five Basic Shapes (50)

      3. The molecular geometry of BF3 is trigonal planar. Because all the B–F bonds are equal and the molecule is highly symmetrical, the dipoles cancel one another in three-dimensional space. Thus BF3 has a net dipole moment of zero:

      10.2: VSEPR Theory - The Five Basic Shapes (51)

      Exercise \(\PageIndex{4}\)

      Which molecule(s) has a net dipole moment?

      • \(\ce{CH3Cl}\)
      • \(\ce{SO3}\)
      • \(\ce{XeO3}\)
      Answer

      \(\ce{CH3Cl}\) and \(\ce{XeO3}\)

      Summary

      Lewis electron structures give no information about molecular geometry, the arrangement of bonded atoms in a molecule or polyatomic ion, which is crucial to understanding the chemistry of a molecule. The valence-shell electron-pair repulsion (VSEPR) model allows us to predict which of the possible structures is actually observed in most cases. It is based on the assumption that pairs of electrons occupy space, and the lowest-energy structure is the one that minimizes electron pair–electron pair repulsions. In the VSEPR model, the molecule or polyatomic ion is given an AXmEn designation, where A is the central atom, X is a bonded atom, E is a nonbonding valence electron group (usually a lone pair of electrons), and m and n are integers. Each group around the central atom is designated as a bonding pair (BP) or lone (nonbonding) pair (LP). From the BP and LP interactions we can predict both the relative positions of the atoms and the angles between the bonds, called the bond angles. From this we can describe the molecular geometry. The VSEPR model can be used to predict the shapes of many molecules and polyatomic ions, but it gives no information about bond lengths and the presence of multiple bonds. A combination of VSEPR and a bonding model, such as Lewis electron structures, is necessary to understand the presence of multiple bonds.

      Molecules with polar covalent bonds can have a dipole moment, an asymmetrical distribution of charge that results in a tendency for molecules to align themselves in an applied electric field. Any diatomic molecule with a polar covalent bond has a dipole moment, but in polyatomic molecules, the presence or absence of a net dipole moment depends on the structure. For some highly symmetrical structures, the individual bond dipole moments cancel one another, giving a dipole moment of zero.

      10.2: VSEPR Theory - The Five Basic Shapes (2024)

      FAQs

      What are the 5 basic shapes of VSEPR theory? ›

      The VSEPR theory describes five main shapes of simple molecules: linear, trigonal planar, tetrahedral, trigonal bipyramidal, and octahedral.

      What is the 10.2 VSEPR theory? ›

      10.2: VSEPR Theory and the Effect of Lone Pairs

      Molecular structure describes the location of the atoms, not the electrons. The geometry that includes all electron pairs is the electron-pair geometry. The electron-pair geometries describe all regions where electrons are located, bonds as well as lone pairs.

      What is the VSEPR theory short answer? ›

      The valence shell electron pair repulsion (VSEPR) theory is a model used to predict 3-D molecular geometry based on the number of valence shell electron bond pairs among the atoms in a molecule or ion. This model assumes that electron pairs will arrange themselves to minimize repulsion effects from one another.

      What are the 5 postulates of VSEPR theory? ›

      The VSEPR theory's postulates :

      The total number of valence shell electron pairs determines the form of the molecule. electron pairs seek to position themselves in such a way that the electron-electron repulsion between them is minimised while the distance between them is maximised.

      What is the VSEPR formula? ›

      The Theory uses the letter A to represent central atoms and X to represent peripheral atoms and E to represent lone pairs to describe the structure of the molecule. There is one central atom (Cl) and 3 peripheral Oxygen atoms and one lone pair of electrons so the VSEPR code would be AX3E.

      What is the VSEPR theory quizlet? ›

      VSEPR Theory. states that the repulsion between electron pairs causes molecular shapes to adjust so that the valence-electron pairs stay as far away as possible.

      What are 6 basic molecular shapes? ›

      The 6 basic molecular shapes are linear, trigonal planar, angular (bent), tetrahedral, trigonal pyramidal, and trigonal bipyramidal.

      What is VSEPR in short notes? ›

      VSEPR theory is used to predict the arrangement of electron pairs around central atoms in molecules, especially simple and symmetric molecules. A central atom is defined in this theory as an atom which is bonded to two or more other atoms, while a terminal atom is bonded to only one other atom.

      What is the VSEPR theory for kids? ›

      Lesson Summary

      VSEPR stands for Valence Shell Electron Pair Repulsion. It states that electron pairs in the valence shell of an atom repel each other; their molecular geometry is the result of this repulsion.

      What are the 5 main VSEPR shapes? ›

      Five Shapes of VSEPR:
      • Linear.
      • Trigonal planar.
      • Tetrahedral.
      • Trigonal bipyramidal.
      • Octahedral.

      What is the postulate 5 in short answer? ›

      Postulate 5: If a straight line falling on two straight lines makes the interior angles on the same side of it taken together less than two right angles, then the two straight lines, if produced indefinitely, meet on that side on which the sum of angles is less than two right angles.

      What are the first 5 postulates? ›

      The five postulates on which Euclid based his geometry are:
      • To draw a straight line from any point to any point.
      • To produce a finite straight line continuously in a straight line.
      • To describe a circle with any center and distance.
      • That all right angles are equal to one another.

      What shapes of molecules have 5 electron pairs? ›

      Five electron pairs give a starting point that is a trigonal bipyramidal structure.

      What are the basic features of Vsepr model? ›

      The main postulates of VSEPR theory are as follows : i The shape of a molecule depends upon the number of valence shell electron pairs around the central atom. ii Pairs of electrons in the valence shell repel one another since their electron clouds are negatively charged.

      What is the steric number 5? ›

      A molecule with steric number 5 has a trigonal bipyramidal shape. A molecule with steric number 6 has an octahedral shape. A molecule with steric number 7 has a pentagonal bipyramidal shape.

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