Molecular Geometry & Bonding Theories

Molecular Geometry & Bonding Theories

Molecular ShapesVSEPR Model

Molecular PolarityCovalent BondingHybrid OrbitalsMultiple Bonds

Molecular ShapesVSEPR Model

Molecular PolarityCovalent BondingHybrid OrbitalsMultiple Bonds

IntroductionIntroduction• Molecules have shapes and sizes that are defined by the angles and distances between nuclei of atoms.

• Shape, size, and strength and polarity of bonds determine properties of a substance.

• We start with Lewis structures to determine the number and types of bonds between atoms.

• Molecules have shapes and sizes that are defined by the angles and distances between nuclei of atoms.

• Shape, size, and strength and polarity of bonds determine properties of a substance.

• We start with Lewis structures to determine the number and types of bonds between atoms.

Electron DomainElectron Domain

• Electron domain: a region in which e-’s will most likely be found.

• Bonding pair: e- domain between two atoms.

• Non-bonding pair (lone pair): e-

domain located mainly on one atom.• An e- domain consists of a nonbonding pair, a single bond, or a multiple bond.

• Electron domain: a region in which e-’s will most likely be found.

• Bonding pair: e- domain between two atoms.

• Non-bonding pair (lone pair): e-

domain located mainly on one atom.• An e- domain consists of a nonbonding pair, a single bond, or a multiple bond.

Determining Number of Electron DomainsDetermining Number of Electron Domains

• VSEPR Model = “Valence Shell Electron Pair Repulsion” Model

• e-’s are negatively charged, so they repel each other.

• VSEPR Model says that the best arrangement of a given number of e- domains is the one that minimizes the repulsions among them.

• VSEPR Model = “Valence Shell Electron Pair Repulsion” Model

• e-’s are negatively charged, so they repel each other.

• VSEPR Model says that the best arrangement of a given number of e- domains is the one that minimizes the repulsions among them.

VSEPR Model Using BalloonsVSEPR Model Using Balloons•Electron domain geometry: arrangement of e- domains around a central atom.

•Electron domain geometry: arrangement of e- domains around a central atom.

Electron-Domain Geometries

Electron-Domain Geometries

Molecular GeometryMolecular Geometry

• The actual spacial arrangement of atoms.

• Molecular geometry is determined from e--domain geometry.

• To predict molecular shapes with VSEPR model, draw Lewis structure, count e--domains, then use the arrangement to determine molecular geometry.

• The actual spacial arrangement of atoms.

• Molecular geometry is determined from e--domain geometry.

• To predict molecular shapes with VSEPR model, draw Lewis structure, count e--domains, then use the arrangement to determine molecular geometry.

Molecular Shapes from Electron-Domain Geometries

Molecular Shapes from Electron-Domain Geometries

More Molecular Shapes from Electron-Domain Geometries

More Molecular Shapes from Electron-Domain Geometries

Molecular Geometry PracticeMolecular Geometry Practice• Predict the molecular geometries:

• Predict the molecular geometries:

•SnCl3-

•SeCl2

•CO32-

•O3

Lone Pair Electrons on Trigonal Bipyramidal Structures

Lone Pair Electrons on Trigonal Bipyramidal Structures

Nonbonding pairs always occupy equatorial positions on a trigonal bipyramidal structure.

Can you explain why?

Nonbonding pairs always occupy equatorial positions on a trigonal bipyramidal structure.

Can you explain why?

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Lone Pair Electrons on Octahedral StructuresLone Pair Electrons on Octahedral Structures

• Nonbonding pairs always occupy the axial positions first.

• Why?

• Nonbonding pairs always occupy the axial positions first.

• Why?

Effect of Non-bonding Electrons & Multiple Bonds

Effect of Non-bonding Electrons & Multiple Bonds•Non-bonding pairs and multiple bonds exert greater repulsive forces on adjacent domains and tend to compress bond angles.

•Lone pair > triple bond > double bond > single bond

•Non-bonding pairs and multiple bonds exert greater repulsive forces on adjacent domains and tend to compress bond angles.

•Lone pair > triple bond > double bond > single bond

Bond Angle of Water Bond Angle of Water• Water’s H-O-H bond angle is always 104.5º due to two lone electron pairs.

• Water’s H-O-H bond angle is always 104.5º due to two lone electron pairs.

Bond Angle of Methane and AmmoniaBond Angle of Methane and Ammonia

• NH3 has a smaller bond angle than methane due to the lone pair of electrons.

• NH3 has a smaller bond angle than methane due to the lone pair of electrons.

Bond DipolesBond Dipoles

• Bond dipoles and dipole moments are vectors (magnitude + direction).

• Overall dipole moment of a molecule is the sum of its bond dipoles.

• Bond dipoles and dipole moments are vectors (magnitude + direction).

• Overall dipole moment of a molecule is the sum of its bond dipoles.

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Polarity of WaterPolarity of Water• If the water molecule were linear, water would NOT be polar.

• If the water molecule were linear, water would NOT be polar.

Polar or Nonpolar?Polar or Nonpolar?• BrCl• Yes. All diatomics with polar bonds are polar molecules.

• SO2

• Yes. Bent molecule. O’s more neg.• NF3

• Yes. Trigonal bipyramidal geometry.• BCl3

• No. Symmetry in trigonal planar geometry• SF6

• No. Symmetry in octahedral arrangement.

• BrCl• Yes. All diatomics with polar bonds are polar molecules.

• SO2

• Yes. Bent molecule. O’s more neg.• NF3

• Yes. Trigonal bipyramidal geometry.• BCl3

• No. Symmetry in trigonal planar geometry• SF6

• No. Symmetry in octahedral arrangement.

Molecular Shape & Molecular PolarityMolecular Shape & Molecular Polarity• Molecular polarity has a significant effect on physical and chemical properties.

• For a molecule with more than two atoms, dipole moment depends on polarities of individual bonds and the molecular geometry.

• Molecular polarity has a significant effect on physical and chemical properties.

• For a molecule with more than two atoms, dipole moment depends on polarities of individual bonds and the molecular geometry.

Valence-Bond TheoryValence-Bond Theory

• Covalent bonds form when a valence atomic orbital of one atom merges with that of another atom.

• The orbitals overlap (share a region of space).

• The overlap allows two e-’s of opposite spin to share common space.

• Covalent bonds form when a valence atomic orbital of one atom merges with that of another atom.

• The orbitals overlap (share a region of space).

• The overlap allows two e-’s of opposite spin to share common space.

Formation of Bonds in H2

Formation of Bonds in H2

• The bond in H2 forms from the overlap of two 1s orbitals from two hydrogen atoms.

• The bond in H2 forms from the overlap of two 1s orbitals from two hydrogen atoms.

Bonds From Orbital OverlapBonds From Orbital Overlap

HybridizationHybridization

• To explain geometries, we assume atomic orbitals mix to form new hybrid orbitals.

• Hybridization: the process of mixing and changing atomic orbitals as atoms approach to form bonds.

• To explain geometries, we assume atomic orbitals mix to form new hybrid orbitals.

• Hybridization: the process of mixing and changing atomic orbitals as atoms approach to form bonds.

Hybridization: spHybridization: sp

• Linear arrangement of e- domains means sp hybridization.

• One s-orbital and one p-orbital hybridize to form two equivalent sp hybrid orbitals.

• Linear arrangement of e- domains means sp hybridization.

• One s-orbital and one p-orbital hybridize to form two equivalent sp hybrid orbitals.

Hybridization: sp2 Hybridization: sp2

• Trigonal planar arrangement means sp2 hybridization.

• One s-orbital and two p-orbitals hybridize to form three sp2 hybrids.

• Trigonal planar arrangement means sp2 hybridization.

• One s-orbital and two p-orbitals hybridize to form three sp2 hybrids.

Hybridization: sp3

Hybridization: sp3

• Tetrahedral arrangement means sp3 hybridization.

• Four sp3 hybrids form from one s-orbital and three p-orbitals.

• Tetrahedral arrangement means sp3 hybridization.

• Four sp3 hybrids form from one s-orbital and three p-orbitals.

Sigma () BondSigma () Bond

• Sigma () bond: e- density concentrated symmetricaly around the line connecting the nuclei (internuclear axis).

• Single bonds are bonds.• Can be made from s- or p-orbitals.

• Allows rotation at bond.

• Sigma () bond: e- density concentrated symmetricaly around the line connecting the nuclei (internuclear axis).

• Single bonds are bonds.• Can be made from s- or p-orbitals.

• Allows rotation at bond.

Pi (∏) BondPi (∏) Bond

• ∏ bond: covalent bond in which there is a side-to-side overlap of p-orbitals.

• p-orbitals are perpendicular to internuclear axis.

• Overlap regions lie above and below internuclear axis.

• Less orbital overlap than in bonds, so ∏ bonds are weaker.

• Does NOT allow rotation around bond.

• ∏ bond: covalent bond in which there is a side-to-side overlap of p-orbitals.

• p-orbitals are perpendicular to internuclear axis.

• Overlap regions lie above and below internuclear axis.

• Less orbital overlap than in bonds, so ∏ bonds are weaker.

• Does NOT allow rotation around bond.

Double & Triple Bonds

Double & Triple Bonds

• Double bonds consists of one bond and one ∏ bond.

• Triple bonds consist of one bond and two ∏ bonds.

• In a double bond, one set of p-orbitals overlap above and below the internuclear axis.

• In a triple bond, the second set of p-orbitals overlap in front of and behind the internuclear axis.

• Double bonds consists of one bond and one ∏ bond.

• Triple bonds consist of one bond and two ∏ bonds.

• In a double bond, one set of p-orbitals overlap above and below the internuclear axis.

• In a triple bond, the second set of p-orbitals overlap in front of and behind the internuclear axis.

Bonding in EthyleneBonding in Ethylene

Ethylene is planar.

Bonding in AcetyleneBonding in Acetylene

Delocalized ∏ BondingDelocalized ∏ Bonding

• Resonance structures with ∏ bonds have delocalization of electrons.

• The electrons in these bonds extend over more than two bonded atoms.

• Resonance structures with ∏ bonds have delocalization of electrons.

• The electrons in these bonds extend over more than two bonded atoms.