Molecular Geometry CH7 by Owen Borville 10.15.2025
Molecular shape and geometry can be predicted by using the valence shell electron-pair repulsion (VSEPR) model.
ABx => where A is the central atom surrounded by x B atoms. x can have integer values of 2 to 6. Examples include:
AB2 = CO2, BeH2, BeCl2, SO2, H2O, NO2-
AB3 = BF3, NH3, ClF3, (SO3)2-
AB4 = CH4, SF4, CCl4, NH4+, SF4, XeF4, ClO4-
AB5 = PCl5, AsF5, SbCl5, IF5, SbF5, BrF5
AB6 = SF6, UF6, (TiCl6)3-
The basis of the VSEPR model is that electrons repel each other. Electrons are found in various domains, including lone pairs, single bonds, double bonds, and triple bonds.
A molecule with two double bonds has two electron domains on the central atom. A molecule with one single bond, one double bond, and one lone pair has three electron domains on the central atom. A molecule with three single bonds and one lone pair has four electron domains on the central atom.
Electrons will arrange themselves to be as far apart as possible in the molecular geometry. Arrangements minimize repulsive interactions:
Two electron domains create a linear molecular geometry.
Three electron domains create a trigonal planar molecular geometry.
Four electron domains will create a tetrahedral molecular geometry.
Five electron domains will create a trigonal bipyramidal molecular geometry.
Six electron domains will create octahedral molecular geometry.
The electron domain geometry is the arrangement of electron domains around the central atom. The molecular geometry is the arrangement of bonded atoms. In an ABx molecule, a bond angle is the angle between two adjacent A-B bonds.
AB5 molecules contain two different bond angles between adjacent bonds. Axial positions are perpendicular to the trigonal plane. Equatorial positions are three bonds arranged in a trigonal plane.
When the central atom in an ABx molecule bears one or more lone pairs, the electron-domain geometry and the molecular geometry are no longer the same. The electron-domain geometry is trigonal planar while the molecular geometry is bent.
Steps to determine the electron-domain and molecular geometries: (1) Draw the Lewis Structure of the molecule or polyatomic ion. (2) Count the number of electron domains on the central atom. (3) Determine the electron-domain geometry by applying the VSEPR model. (4) Determine the molecular geometry by considering the positions of the atoms only.
Some electron domains are better than others at repelling neighboring domains. Lone pairs take up more space than bonded pairs of electrons. Multiple bonds repel more strongly than single bonds.
The geometry of more complex molecules can be determined by treating them as though they have multiple central atoms.
Molecular polarity is one of the most important consequences of molecular geometry. A diatomic molecule is polar when the electronegativities of the two atoms are different, such as with the HF molecule, which is polar toward the fluorine atom.
The polarity of a molecule made up of three or more atoms depends on the polarity of the individual bonds and the molecular geometry. The bonds in CO2 are polar, but the molecule is non-polar and linear.
The bonds in H2O are polar AND the molecule is polar. The molecule geometry is also bent.
The bonds in BF3 are polar but the molecule is nonpolar.
The bonds in CCl4 are polar, but the molecule is nonpolar.
The bonds in CHCl3 are polar and the molecule is polar.
Dipole moments can be used to distinguish between structural isomers. Trans-dichloroethylene is nonpolar while cis-dichloroethylene is polar.
Intermolecular forces are attractive forces between neighboring molecules. These forces are known as van der Walls forces. The magnitude (and type) of intermolecular forces can be sufficient to hold the molecules of a substance together in a solid or liquid. Gases have no apparent intermolecular forces.
Dipole-dipole interactions are attractive forces that act between polar molecules. The magnitude of the attractive forces depends on the magnitude of the dipole. Red is high electron density (𝛿-) and blue is low density (𝛿+).
Hydrogen bonding is a special type of dipole-dipole interaction. Hydrogen bonding only occurs in molecules that contain H bonded to a small, highly electronegative atom such as N, O, or F.
Dispersion forces or London dispersion forces result from the Coulombic attractions between instantaneous dipoles of nonpolar molecules.
Ion-dipole interactions are Coulombic attractions between ions (either positive or negative) and polar molecules.
According to valence bond theory, atoms share electrons when atomic orbitals overlap. (1) A bond forms when single occupied atomic orbitals on two atoms overlap. (2) The two electrons shared in the region of orbital overlap must be of opposite spin. (3) Formation of a bond results in a lower potential energy for the system.
The H-H bond in H2 forms when the singly occupied 1s orbitals of the two H atoms overlap. The F-F bond in F2 forms when the singly occupied 2p orbitals of the two F atoms overlap.
The H-F bond in HF forms when the singly occupied 1s orbital on the H atom overlaps with the single occupied 2p orbital of the F atom.
A covalent bond will form if the potential energy of the molecule is lower than the combined potential energies of the isolated atoms.
Hybridization or mixing of different atomic orbitals to create new hybrid orbitals that are correctly oriented for covalent bonding according to VSEPR theory and can account for observed bond angles in molecules that could not be described by the direct overlap of atomic orbitals. Hybrid orbitals are created by combining an atom's atomic orbitals (such as s and p) to form new shapes and orientations suitable for bonding in a molecule.
BeCl2 => Lewis theory and VSEPR theory predict Cl-Be-Cl bond angle of 180 degrees with two electron domains and linear molecular geometry. A ground state beryllium atom cannot form two bonds because there are no unpaired electrons. An excited state configuration for Be has two unpaired electrons and can form two bonds. The two bonds formed would not be equivalent. Experimentally the bond in BeCl2 bonds is identical in length and strength. Mixing of one s orbital and one p orbital to yield two sp orbitals. The 2s orbital and one of the 2p orbitals on Be combine to form two sp hybrid orbitals. Like any two electron domains, the hybrid orbitals on Be are 180 degrees apart. The hybrid orbitals on Be each overlap with a singly occupied 3p orbital on a Cl atom. The energy required to form an excited state Be atom is more than compensated for by the energy given off when a bond forms.
BF3 => (1) Draw the Lewis structure. (2) Determine the number and type of hybrid orbitals necessary to rationalize the bonding in BF3. Count the number of electron domains on the central atom. Three hybrid orbitals are required. Determine the number and type of hybrid orbitals necessary to rationalize the bonding in BF3. (3) Draw the ground-state orbital diagram for the central atom. (4) Maximize the number of unpaired valence electrons by promotion. Determine the number and type of hybrid orbitals necessary to rationalize the bonding in BF3. (5) Combine the necessary number of atomic orbitals to generate the required number of hybrid orbitals. (6) Place electrons in the hybrid orbitals, putting one electron in each orbital before pairing any electrons. The mixing of one s orbital and two p orbitals yields three sp2 orbitals. Hybrid orbitals on boron overlap with 2p orbitals on fluorine.
CH4 => (1) Draw the Lewis structure. Determine the number and type of hybrid orbitals necessary to rationalize the bonding in CH4. (2) the number of electron domains on the central atom is the number of hybrid orbitals necessary to account for the molecule's geometry => Four electron domains, and four hybrid orbitals required. Determine the number and type of hybrid orbitals necessary to rationalize the bonding in CH4. (3) Draw the ground state orbital diagram for the central atom. (4) Maximize the number of unpaired electrons by promotion. Determine the number and type of hybrid orbitals necessary to rationalize the bonding in CH4. (5) Combine the necessary number of atomic orbitals to generate the required number of hybrid orbitals. (6) Place electrons in the hybrid orbitals, putting one electron in each orbital before pairing any electrons. The mixing of one s orbital and three p orbitals to yield four sp3 orbitals. Hybrid orbitals on carbon overlap with 1s orbitals on hydrogen.
PCl5 => (1) To find the number of hybrid orbitals bonded, draw the Lewis Structure. (2) The number of electron domains on the central atom is the number of hybrid orbitals necessary to account for the molecule's geometry. Five electron domains identify five hybrid orbitals required. (3) Determine the number and type of hybrid orbitals necessary to rationalize the bonding in PCl5. Draw the ground state orbital diagram for the central atom. (4) Maximize the number of unpaired electrons by promotion. (5) Combine the necessary number of atomic orbitals to generate the required number of hybrid orbitals. (6) Place electrons in the hybrid orbitals, putting one electron in each orbital before pairing any electrons. Hybrid orbitals on phosphorus overlap with 3p orbitals on chlorine.
NH3 is a trigonal pyramidal molecule with three equivalent N-H bonds.
Valence bond theory and hybridization can be used to describe the bonding in molecules containing double and triple bonds.
For example, in ethylene (C2H4), each carbon has three electron domains: 2 single bonds and 1 double bond. Expect sp2 hybridization. Maximize unpaired electrons on carbon by promotion. Hybridize the required number of atomic orbitals (one for each electron domain on carbon). Hybridization scheme for the carbon atoms in ethylene: One unhybridized atomic 2p orbital gives rise to multiple bonds. Three equivalent sp2 hybrid orbitals explain three bonds around carbon.
Sigma (σ) bonds form when sp2 hybrid orbitals on the C atoms overlap. In a sigma bond, the shared electron density lies directly along the internuclear axis. The ethylene molecule contains five sigma bonds: 1 between the two carbon atoms (sp2 and sp2 overlap); 4 between the C and H atoms (sp2 and 1s overlap).
The remaining unhybridized p orbital is perpendicular to the plane in which the atoms of the molecule lie. The unhybridized p orbitals overlap in a sideways fashion to form a pi (π) bond.
Sigma (σ) bonds exhibit free rotation around the bond axis. Pi (π) bonds restrict free rotation around the bond axis.
The acetylene molecule (C2H2) is linear with sp hybridized carbons. Promotion of an electron maximizes the number of unpaired electrons. In acetylene, there are 3 sigma bonds, 1 between the two carbon atoms (sp and sp) and 2 sigma bonds between the C and H atoms (sp and 1s). There are 2 pi bonds: 2 in between the two carbon atoms (2p and 2p). The 2s orbital and one of the 2p orbitals then mix to form two sp hybrid orbitals. The acetylene molecule is linear with sp hybridized carbons. Two unhybridized atomic 2p orbitals gives rise to 2 pi bonds. Two equivalent sp hybrid orbitals give rise to 2 sigma bonds. Formation of the C-C sigma bond in acetylene: 3 sigma bonds, 1 between the two carbon atoms (sp and sp), 2 between the C and H atoms (sp and 1s). Formation of the C-C pi bond in acetylene: Two pi bonds, two between the two carbon atoms (2p and 2p).
Lewis structures and valence bond theory fail to predict some important properties of molecules. Paramagnetism is a result of a molecule's electron configuration. Species that contain one or more unpaired electrons are paramagnetic. Paramagnetic species are attracted to magnetic fields. Ex. The Lewis structure for O2 shows no unpaired electrons. The correct structure, based on experimental evidence, shows two unpaired electrons. O2 exhibits paramagnetism.
Lewis structures and valence bond theory fail to predict some important properties of molecules, such as diamagnetism. Species that contain paired electrons are diamagnetic. Diamagnetic species are weakly repelled by magnetic fields. The Lewis Structure for N2 shows no unpaired electrons. N2 exhibits diamagnetism.
Molecular orbital theory is another bonding theory is needed to describe the paramagnetism of O2. In molecular orbital theory, the atomic orbitals combine to form new orbitals that are the property of the entire molecule. The new orbitals are called molecular orbitals. Molecular orbitals have characteristics similar to atomic orbitals, including: specific shapes, specific energies, accommodation of a maximum of 2 electrons each, electron filling follows the Pauli exclusion principle, and the number of molecular orbitals obtained equals the number of orbitals combined.
H2 is the simplest homonuclear diatomic molecule. According to valence bond theory, H2 forms from the overlap of the 1s orbitals. According to molecular orbital theory, H2 forms when the 1s orbitals combine to give molecular orbitals. Molecular orbitals result from the constructive and destructive combination of atomic orbitals.
Constructive combination of the two 1s orbitals gives rise to a molecular orbital that lies along the internuclear axis. Constructive combination increases the electron density between the two nuclei. This molecular orbital is referred to as a bonding molecular orbital.
Destructive combination of the two 1s orbitals gives rise to a molecular orbital that lies along the internuclear axis, but does not lie between the two nuclei. Electron density in this molecular orbital pulls the two nuclei in opposite directions. This molecular orbital is referred to as an antibonding molecular orbital.
Molecular orbitals (σ) are orbitals that lie along the internuclear axis. Examples are σ1s bonding molecular orbital from the combination of two 1s orbitals. σ*1s antibonding molecular orbital from the combination of two 1s orbitals. The * distinguishes an antibonding molecular orbital from a bonding orbital.
Molecular orbitals have specific energies. Electrons in bonding molecular orbitals stabilize the molecule and are lower in energy than the isolated atomic orbitals. Electrons in antibonding molecular orbitals destabilize the molecule and are higher in energy than the isolated atomic orbitals.
Bond Order: The bond order indicates how stable the molecule is. The higher the bond order, the more stable a molecule is.
Bond Order = (number of electrons (e-) in bonding orbitals - number of electrons (e-) in antibonding orbitals)/2
According to molecular orbital theory, H2 is a stable molecule because bond order = (2-0)/2 = 1
However, He2 is not a stable molecule and does not exist, because bond order = (2-2)/2 = 0.
P atomic orbitals also form molecular orbitals by both constructive and destructive combination. The orientations of px, py, and pz give rise to two different types of molecular orbitals:
σ molecular orbitals - electron density along the internuclear axis. Px orbitals point towards each other. Bonding and antibonding σ molecular orbitals.
π molecular orbitals - electron density above and below the internuclear axis.
Molecular orbitals resulting from the combination of p atomic orbitals are higher than those resulting from the combination of s atomic orbitals. Some orders of orbital energies assume no mixing of s and p orbitals. S orbitals only interact with s orbitals. P orbitals only interact with p orbitals. This is found in O2, F2, and Ne2.
Some mixing of s and p orbitals occurs, such as in orbitals of Li2, B2, C2, and N2.
Filling molecular orbital diagrams follows the same rules as the filling of atomic orbitals. (1) lower energy orbitals fill first (2) each orbital can accommodate a maximum of two electrons with opposite spin. (3) Hund's rule is obeyed.
Molecular orbital diagrams for second-period homonuclear diatomic molecules.
Bonding Theories and Descriptions of Molecules with Delocalized Bonding:
Lewis Theory: Strength: Qualitative prediction of bond strength and bond length. Weakness: Two-dimensional model, real molecules are three-dimensional. Fails to explain why bonds form.
Valence-Shell Electron-Pair Repulsion Model (VSEPR): Strength: Predicts the shape of many molecules and polyatomic ions. Weakness: fails to explain why bonds form (based on Lewis theory).
Valence Bond Theory: Strength: Covalent bonds form when atomic orbitals overlap. Weakness: Fails to explain the bonding in many molecules.
Hybridization of Atomic Orbitals: Strength: An extension of valence bond theory. Using hybrid orbitals, it is possible to explain the bonding and geometry of more molecules. Weakness: Fails to predict some important properties, such as magnetism.
Molecular Orbital Theory: Strength: Accurately predicts the magnetic properties and other properties of molecules. Weakness: complex.
Some molecules are best described by using a combination of models. For example, benzene, C6H6, is represented with two resonance structures. The π bonds in benzene are delocalized or spread out over the entire molecule. Another example is the carbonate ion, (CO3)2-, which needs three resonance structures to represent the ion.
Molecular shape and geometry can be predicted by using the valence shell electron-pair repulsion (VSEPR) model.
ABx => where A is the central atom surrounded by x B atoms. x can have integer values of 2 to 6. Examples include:
AB2 = CO2, BeH2, BeCl2, SO2, H2O, NO2-
AB3 = BF3, NH3, ClF3, (SO3)2-
AB4 = CH4, SF4, CCl4, NH4+, SF4, XeF4, ClO4-
AB5 = PCl5, AsF5, SbCl5, IF5, SbF5, BrF5
AB6 = SF6, UF6, (TiCl6)3-
The basis of the VSEPR model is that electrons repel each other. Electrons are found in various domains, including lone pairs, single bonds, double bonds, and triple bonds.
A molecule with two double bonds has two electron domains on the central atom. A molecule with one single bond, one double bond, and one lone pair has three electron domains on the central atom. A molecule with three single bonds and one lone pair has four electron domains on the central atom.
Electrons will arrange themselves to be as far apart as possible in the molecular geometry. Arrangements minimize repulsive interactions:
Two electron domains create a linear molecular geometry.
Three electron domains create a trigonal planar molecular geometry.
Four electron domains will create a tetrahedral molecular geometry.
Five electron domains will create a trigonal bipyramidal molecular geometry.
Six electron domains will create octahedral molecular geometry.
The electron domain geometry is the arrangement of electron domains around the central atom. The molecular geometry is the arrangement of bonded atoms. In an ABx molecule, a bond angle is the angle between two adjacent A-B bonds.
AB5 molecules contain two different bond angles between adjacent bonds. Axial positions are perpendicular to the trigonal plane. Equatorial positions are three bonds arranged in a trigonal plane.
When the central atom in an ABx molecule bears one or more lone pairs, the electron-domain geometry and the molecular geometry are no longer the same. The electron-domain geometry is trigonal planar while the molecular geometry is bent.
Steps to determine the electron-domain and molecular geometries: (1) Draw the Lewis Structure of the molecule or polyatomic ion. (2) Count the number of electron domains on the central atom. (3) Determine the electron-domain geometry by applying the VSEPR model. (4) Determine the molecular geometry by considering the positions of the atoms only.
Some electron domains are better than others at repelling neighboring domains. Lone pairs take up more space than bonded pairs of electrons. Multiple bonds repel more strongly than single bonds.
The geometry of more complex molecules can be determined by treating them as though they have multiple central atoms.
Molecular polarity is one of the most important consequences of molecular geometry. A diatomic molecule is polar when the electronegativities of the two atoms are different, such as with the HF molecule, which is polar toward the fluorine atom.
The polarity of a molecule made up of three or more atoms depends on the polarity of the individual bonds and the molecular geometry. The bonds in CO2 are polar, but the molecule is non-polar and linear.
The bonds in H2O are polar AND the molecule is polar. The molecule geometry is also bent.
The bonds in BF3 are polar but the molecule is nonpolar.
The bonds in CCl4 are polar, but the molecule is nonpolar.
The bonds in CHCl3 are polar and the molecule is polar.
Dipole moments can be used to distinguish between structural isomers. Trans-dichloroethylene is nonpolar while cis-dichloroethylene is polar.
Intermolecular forces are attractive forces between neighboring molecules. These forces are known as van der Walls forces. The magnitude (and type) of intermolecular forces can be sufficient to hold the molecules of a substance together in a solid or liquid. Gases have no apparent intermolecular forces.
Dipole-dipole interactions are attractive forces that act between polar molecules. The magnitude of the attractive forces depends on the magnitude of the dipole. Red is high electron density (𝛿-) and blue is low density (𝛿+).
Hydrogen bonding is a special type of dipole-dipole interaction. Hydrogen bonding only occurs in molecules that contain H bonded to a small, highly electronegative atom such as N, O, or F.
Dispersion forces or London dispersion forces result from the Coulombic attractions between instantaneous dipoles of nonpolar molecules.
Ion-dipole interactions are Coulombic attractions between ions (either positive or negative) and polar molecules.
According to valence bond theory, atoms share electrons when atomic orbitals overlap. (1) A bond forms when single occupied atomic orbitals on two atoms overlap. (2) The two electrons shared in the region of orbital overlap must be of opposite spin. (3) Formation of a bond results in a lower potential energy for the system.
The H-H bond in H2 forms when the singly occupied 1s orbitals of the two H atoms overlap. The F-F bond in F2 forms when the singly occupied 2p orbitals of the two F atoms overlap.
The H-F bond in HF forms when the singly occupied 1s orbital on the H atom overlaps with the single occupied 2p orbital of the F atom.
A covalent bond will form if the potential energy of the molecule is lower than the combined potential energies of the isolated atoms.
Hybridization or mixing of different atomic orbitals to create new hybrid orbitals that are correctly oriented for covalent bonding according to VSEPR theory and can account for observed bond angles in molecules that could not be described by the direct overlap of atomic orbitals. Hybrid orbitals are created by combining an atom's atomic orbitals (such as s and p) to form new shapes and orientations suitable for bonding in a molecule.
BeCl2 => Lewis theory and VSEPR theory predict Cl-Be-Cl bond angle of 180 degrees with two electron domains and linear molecular geometry. A ground state beryllium atom cannot form two bonds because there are no unpaired electrons. An excited state configuration for Be has two unpaired electrons and can form two bonds. The two bonds formed would not be equivalent. Experimentally the bond in BeCl2 bonds is identical in length and strength. Mixing of one s orbital and one p orbital to yield two sp orbitals. The 2s orbital and one of the 2p orbitals on Be combine to form two sp hybrid orbitals. Like any two electron domains, the hybrid orbitals on Be are 180 degrees apart. The hybrid orbitals on Be each overlap with a singly occupied 3p orbital on a Cl atom. The energy required to form an excited state Be atom is more than compensated for by the energy given off when a bond forms.
BF3 => (1) Draw the Lewis structure. (2) Determine the number and type of hybrid orbitals necessary to rationalize the bonding in BF3. Count the number of electron domains on the central atom. Three hybrid orbitals are required. Determine the number and type of hybrid orbitals necessary to rationalize the bonding in BF3. (3) Draw the ground-state orbital diagram for the central atom. (4) Maximize the number of unpaired valence electrons by promotion. Determine the number and type of hybrid orbitals necessary to rationalize the bonding in BF3. (5) Combine the necessary number of atomic orbitals to generate the required number of hybrid orbitals. (6) Place electrons in the hybrid orbitals, putting one electron in each orbital before pairing any electrons. The mixing of one s orbital and two p orbitals yields three sp2 orbitals. Hybrid orbitals on boron overlap with 2p orbitals on fluorine.
CH4 => (1) Draw the Lewis structure. Determine the number and type of hybrid orbitals necessary to rationalize the bonding in CH4. (2) the number of electron domains on the central atom is the number of hybrid orbitals necessary to account for the molecule's geometry => Four electron domains, and four hybrid orbitals required. Determine the number and type of hybrid orbitals necessary to rationalize the bonding in CH4. (3) Draw the ground state orbital diagram for the central atom. (4) Maximize the number of unpaired electrons by promotion. Determine the number and type of hybrid orbitals necessary to rationalize the bonding in CH4. (5) Combine the necessary number of atomic orbitals to generate the required number of hybrid orbitals. (6) Place electrons in the hybrid orbitals, putting one electron in each orbital before pairing any electrons. The mixing of one s orbital and three p orbitals to yield four sp3 orbitals. Hybrid orbitals on carbon overlap with 1s orbitals on hydrogen.
PCl5 => (1) To find the number of hybrid orbitals bonded, draw the Lewis Structure. (2) The number of electron domains on the central atom is the number of hybrid orbitals necessary to account for the molecule's geometry. Five electron domains identify five hybrid orbitals required. (3) Determine the number and type of hybrid orbitals necessary to rationalize the bonding in PCl5. Draw the ground state orbital diagram for the central atom. (4) Maximize the number of unpaired electrons by promotion. (5) Combine the necessary number of atomic orbitals to generate the required number of hybrid orbitals. (6) Place electrons in the hybrid orbitals, putting one electron in each orbital before pairing any electrons. Hybrid orbitals on phosphorus overlap with 3p orbitals on chlorine.
NH3 is a trigonal pyramidal molecule with three equivalent N-H bonds.
Valence bond theory and hybridization can be used to describe the bonding in molecules containing double and triple bonds.
For example, in ethylene (C2H4), each carbon has three electron domains: 2 single bonds and 1 double bond. Expect sp2 hybridization. Maximize unpaired electrons on carbon by promotion. Hybridize the required number of atomic orbitals (one for each electron domain on carbon). Hybridization scheme for the carbon atoms in ethylene: One unhybridized atomic 2p orbital gives rise to multiple bonds. Three equivalent sp2 hybrid orbitals explain three bonds around carbon.
Sigma (σ) bonds form when sp2 hybrid orbitals on the C atoms overlap. In a sigma bond, the shared electron density lies directly along the internuclear axis. The ethylene molecule contains five sigma bonds: 1 between the two carbon atoms (sp2 and sp2 overlap); 4 between the C and H atoms (sp2 and 1s overlap).
The remaining unhybridized p orbital is perpendicular to the plane in which the atoms of the molecule lie. The unhybridized p orbitals overlap in a sideways fashion to form a pi (π) bond.
Sigma (σ) bonds exhibit free rotation around the bond axis. Pi (π) bonds restrict free rotation around the bond axis.
The acetylene molecule (C2H2) is linear with sp hybridized carbons. Promotion of an electron maximizes the number of unpaired electrons. In acetylene, there are 3 sigma bonds, 1 between the two carbon atoms (sp and sp) and 2 sigma bonds between the C and H atoms (sp and 1s). There are 2 pi bonds: 2 in between the two carbon atoms (2p and 2p). The 2s orbital and one of the 2p orbitals then mix to form two sp hybrid orbitals. The acetylene molecule is linear with sp hybridized carbons. Two unhybridized atomic 2p orbitals gives rise to 2 pi bonds. Two equivalent sp hybrid orbitals give rise to 2 sigma bonds. Formation of the C-C sigma bond in acetylene: 3 sigma bonds, 1 between the two carbon atoms (sp and sp), 2 between the C and H atoms (sp and 1s). Formation of the C-C pi bond in acetylene: Two pi bonds, two between the two carbon atoms (2p and 2p).
Lewis structures and valence bond theory fail to predict some important properties of molecules. Paramagnetism is a result of a molecule's electron configuration. Species that contain one or more unpaired electrons are paramagnetic. Paramagnetic species are attracted to magnetic fields. Ex. The Lewis structure for O2 shows no unpaired electrons. The correct structure, based on experimental evidence, shows two unpaired electrons. O2 exhibits paramagnetism.
Lewis structures and valence bond theory fail to predict some important properties of molecules, such as diamagnetism. Species that contain paired electrons are diamagnetic. Diamagnetic species are weakly repelled by magnetic fields. The Lewis Structure for N2 shows no unpaired electrons. N2 exhibits diamagnetism.
Molecular orbital theory is another bonding theory is needed to describe the paramagnetism of O2. In molecular orbital theory, the atomic orbitals combine to form new orbitals that are the property of the entire molecule. The new orbitals are called molecular orbitals. Molecular orbitals have characteristics similar to atomic orbitals, including: specific shapes, specific energies, accommodation of a maximum of 2 electrons each, electron filling follows the Pauli exclusion principle, and the number of molecular orbitals obtained equals the number of orbitals combined.
H2 is the simplest homonuclear diatomic molecule. According to valence bond theory, H2 forms from the overlap of the 1s orbitals. According to molecular orbital theory, H2 forms when the 1s orbitals combine to give molecular orbitals. Molecular orbitals result from the constructive and destructive combination of atomic orbitals.
Constructive combination of the two 1s orbitals gives rise to a molecular orbital that lies along the internuclear axis. Constructive combination increases the electron density between the two nuclei. This molecular orbital is referred to as a bonding molecular orbital.
Destructive combination of the two 1s orbitals gives rise to a molecular orbital that lies along the internuclear axis, but does not lie between the two nuclei. Electron density in this molecular orbital pulls the two nuclei in opposite directions. This molecular orbital is referred to as an antibonding molecular orbital.
Molecular orbitals (σ) are orbitals that lie along the internuclear axis. Examples are σ1s bonding molecular orbital from the combination of two 1s orbitals. σ*1s antibonding molecular orbital from the combination of two 1s orbitals. The * distinguishes an antibonding molecular orbital from a bonding orbital.
Molecular orbitals have specific energies. Electrons in bonding molecular orbitals stabilize the molecule and are lower in energy than the isolated atomic orbitals. Electrons in antibonding molecular orbitals destabilize the molecule and are higher in energy than the isolated atomic orbitals.
Bond Order: The bond order indicates how stable the molecule is. The higher the bond order, the more stable a molecule is.
Bond Order = (number of electrons (e-) in bonding orbitals - number of electrons (e-) in antibonding orbitals)/2
According to molecular orbital theory, H2 is a stable molecule because bond order = (2-0)/2 = 1
However, He2 is not a stable molecule and does not exist, because bond order = (2-2)/2 = 0.
P atomic orbitals also form molecular orbitals by both constructive and destructive combination. The orientations of px, py, and pz give rise to two different types of molecular orbitals:
σ molecular orbitals - electron density along the internuclear axis. Px orbitals point towards each other. Bonding and antibonding σ molecular orbitals.
π molecular orbitals - electron density above and below the internuclear axis.
Molecular orbitals resulting from the combination of p atomic orbitals are higher than those resulting from the combination of s atomic orbitals. Some orders of orbital energies assume no mixing of s and p orbitals. S orbitals only interact with s orbitals. P orbitals only interact with p orbitals. This is found in O2, F2, and Ne2.
Some mixing of s and p orbitals occurs, such as in orbitals of Li2, B2, C2, and N2.
Filling molecular orbital diagrams follows the same rules as the filling of atomic orbitals. (1) lower energy orbitals fill first (2) each orbital can accommodate a maximum of two electrons with opposite spin. (3) Hund's rule is obeyed.
Molecular orbital diagrams for second-period homonuclear diatomic molecules.
Bonding Theories and Descriptions of Molecules with Delocalized Bonding:
Lewis Theory: Strength: Qualitative prediction of bond strength and bond length. Weakness: Two-dimensional model, real molecules are three-dimensional. Fails to explain why bonds form.
Valence-Shell Electron-Pair Repulsion Model (VSEPR): Strength: Predicts the shape of many molecules and polyatomic ions. Weakness: fails to explain why bonds form (based on Lewis theory).
Valence Bond Theory: Strength: Covalent bonds form when atomic orbitals overlap. Weakness: Fails to explain the bonding in many molecules.
Hybridization of Atomic Orbitals: Strength: An extension of valence bond theory. Using hybrid orbitals, it is possible to explain the bonding and geometry of more molecules. Weakness: Fails to predict some important properties, such as magnetism.
Molecular Orbital Theory: Strength: Accurately predicts the magnetic properties and other properties of molecules. Weakness: complex.
Some molecules are best described by using a combination of models. For example, benzene, C6H6, is represented with two resonance structures. The π bonds in benzene are delocalized or spread out over the entire molecule. Another example is the carbonate ion, (CO3)2-, which needs three resonance structures to represent the ion.