Chemistry Second Edition Julia Burdge Lecture Power Points
Chemistry Second Edition Julia Burdge Lecture Power. Points Jason A. Kautz University of Nebraska-Lincoln 9 Chemical Bonding II: Molecular Geometry and Bonding Theories Copyright (c) The Mc. Graw-Hill Companies, Inc. Permission required for reproduction or display. 1
9 Chemical Bonding II: Molecular Geometry and Bonding Theories 9. 1 Molecular Geometry The VSEPR Model Electron-Domain Geometry and Molecular Geometry Deviation from Ideal Bond Angles Geometry of Molecules with More Than One Central Atom 9. 2 Molecular Geometry and Polarity 9. 3 Valence Bond Theory 9. 4 Hybridization of Atomic Orbitals Hybridization of s and p Orbitals Hybridization of s, p, and d Orbitals 9. 5 Hybridization of Molecules Containing Multiple Bonds 9. 6 Molecular Orbital Theory Bonding and Antibonding Molecular Orbitals σ Molecular Orbitals Bond Order π Molecular Orbitals Molecular Orbital Diagrams 9. 7 Bonding Theories and Descriptions of Molecules with Delocalized Bonding
9. 1 Molecular Geometry Molecular shape can be predicted by using the valence-shell electronpair repulsion (VSEPR) model. ABx A is the central atom surrounded by x B atoms x can have integer values of 2 to 6.
Molecular Geometry The basis of the VSEPR model is that electrons repel each other. Electrons are found in various domains. Lone pairs Single bonds 2 double bonds Double bonds Triple bonds 1 single bond 3 single bonds 1 double bond 1 lone pair 2 electron domains (central atom) 3 electron domains (central atom) 4 electron domains (central atom)
Molecular Geometry The basis of the VSEPR model is that electrons repel each other. Electrons will arrange themselves to be as far apart as possible. Arrangements minimize repulsive interactions. 2 electron domains Linear 3 electron domains Trigonal planar
Molecular Geometry 4 electron domains Tetrahedral 5 electron domains Trigonal bipyramidal 6 electron domains 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. 180° 120° Linear 109. 5° Trigonal planar 90° 120° Tetrahedral Trigonal bipyramidal 90° Octahedral
Molecular Geometry AB 5 molecules contain two different bond angles between adjacent bonds. Axial positions; perpendicular to the trigonal plane 90° Equatorial positions; three bonds arranged in a trigonal plane. 120° Trigonal bipyramidal
Molecular Geometry 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. • • • • O =O− O
Molecular Geometry
Molecular Geometry
Molecular Geometry The steps to determine the electron-domain and molecular geometries are as follows: Step 1: Draw the Lewis structure of the molecule or polyatomic ion. Step 2: Count the number of electron domains on the central atom. Step 3: Determine the electron-domain geometry by applying the VSEPR model. Step 4: Determine the molecular geometry by considering the positions of the atoms only.
Molecular Geometry Determine the shape of SF 4. Solution: Step 1: Draw the Lewis structure of the molecule or polyatomic ion. S F 6 x 1=6 7 x 4 = 28 34 total valence electrons 4 bonds x 2 e– = 8 e– 34 – 8 = 26 e– left over
Molecular Geometry Determine the shape of SF 4. Solution: Step 2: Count the number of electron domains on the central atom. 4 single bonds 1 lone pair 5 electron domains
Molecular Geometry Determine the shape of SF 4. Solution: Step 3: Determine the electron-domain geometry by applying the VSEPR model. 5 electron domains Trigonal bipyrimidal electron domain geometry
Molecular Geometry Determine the shape of SF 4. Solution: Step 4: Determine the molecular geometry by considering the positions of the atoms only. 5 electron domains Trigonal bipyrimidal electron domain geometry See-saw shape
Molecular Geometry 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.
Molecular Geometry The geometry of more complex molecules can be determined by treating them as though they have multiple central atoms. Central O atom No. of electron domains: 4 Electron-domain geometry: tetrahedral Molecular geometry: Central C atom No. of electron domains: 4 Electron-domain geometry: tetrahedral Molecular geometry: tetrahedral bent
9. 2 Molecular Geometry and Polarity Molecular polarity is one of the most important consequences of molecular geometry. A diatomic molecule is polar when the electronegativites of the two atoms are different. • • H−F • • δ− • • H−F • • δ+
Molecular Geometry and Polarity The polarity of a molecule made up of three or more atoms depends on: (1) the polarity of the individual bonds (2) the molecular geometry Carbon dioxide, CO 2 The bonds in CO 2 are polar, the molecule is nonpolar
Molecular Geometry and Polarity The polarity of a molecule made up of three or more atoms depends on: (1) the polarity of the individual bonds (2) the molecular geometry Water, H 2 O The bonds in H 2 O are polar, the molecule is polar
Molecular Geometry and Polarity Dipole moments can be used to distinguish between structural isomers. trans-dichloroethylene nonpolar cis-dichloroethylene polar
9. 3 Valence Bond Theory 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.
Valence Bond Theory The H−H bond in H 2 forms when the singly occupied 1 s orbitals of the two H atoms overlap:
Valence Bond Theory The F−F bond in F 2 forms when the singly occupied 2 p orbitals of the two F atoms overlap:
Valence Bond Theory The H−F bond in HF forms when the singly occupied 1 s orbital on the H atom overlaps with the single occupied 2 p orbital of the F atom:
9. 4 Hybridization of Atomic Orbitals Hybridization or mixing of atomic orbitals can account for observed bond angles in molecules that could not be described by the direct overlap of atomic orbitals.
Hybridization of Atomic Orbitals Be. Cl 2 Lewis theory and VSEPR theory predict Cl−Be−Cl bond angle of 180° 2 electron domains Linear molecular geometry A ground state beryllium atom can not form two bonds; there are no unpaired electrons.
Hybridization of Atomic Orbitals Be. Cl 2 An excited state configuration for Be has two unpaired electrons and can form to bonds. The two bonds formed would not be equivalent.
Hybridization of Atomic Orbitals Be. Cl 2 Experimentally the bond in Be. Cl 2 bonds are identical in length and strength. Mixing of one s orbital and one p orbital to yield two sp orbitals.
Hybridization of Atomic Orbitals Be. Cl 2 The 2 s orbital and one of the 2 p orbitals on Be combine to form two sp hybrid orbitals. Like any two electron domains, the hybrid orbitals on Be are 180° apart.
Hybridization of Atomic Orbitals Be. Cl 2 The hybrid orbitals on Be each overlap with a singly occupied 3 p 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.
Hybridization of Atomic Orbitals Determine the number and type of hybrid orbitals necessary to rationalize the bonding in BF 3 Step 1: Draw the Lewis structure:
Hybridization of Atomic Orbitals Determine the number and type of hybrid orbitals necessary to rationalize the bonding in BF 3 Step 2: The number of electron domains on the central atom is the number of hybrid orbitals necessary to account for the molecule’s geometry. Three electron domains Three hybrid orbitals required
Hybridization of Atomic Orbitals Determine the number and type of hybrid orbitals necessary to rationalize the bonding in BF 3 Step 3: Draw the ground-state orbital diagram for the central atom. Step 4: Maximize the number of unpaired electrons by promotion.
Hybridization of Atomic Orbitals Determine the number and type of hybrid orbitals necessary to rationalize the bonding in BF 3 Step 5: Combine the necessary number of atomic orbitals to generate the required number of hybrid orbitals. Step 6: Place electrons in the hybrid orbitals, putting one electron in each orbital before pairing any electrons.
Hybridization of Atomic Orbitals Mixing of one s orbital and two p orbitals to yield three sp 2 orbitals.
Hybridization of Atomic Orbitals Hybrid orbitals on boron overlap with 2 p orbitals on fluorine.
Hybridization of Atomic Orbitals Determine the number and type of hybrid orbitals necessary to rationalize the bonding in CH 4. Step 1: Draw the Lewis structure:
Hybridization of Atomic Orbitals Determine the number and type of hybrid orbitals necessary to rationalize the bonding in CH 4. Step 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 Four hybrid orbitals required
Hybridization of Atomic Orbitals Determine the number and type of hybrid orbitals necessary to rationalize the bonding in CH 4. Step 3: Draw the ground-state orbital diagram for the central atom. Step 4: Maximize the number of unpaired electrons by promotion.
Hybridization of Atomic Orbitals Determine the number and type of hybrid orbitals necessary to rationalize the bonding in CH 4. Step 5: Combine the necessary number of atomic orbitals to generate the required number of hybrid orbitals. Step 6: Place electrons in the hybrid orbitals, putting one electron in each orbital before pairing any electrons.
Hybridization of Atomic Orbitals Mixing of one s orbital and three p orbitals to yield four sp 3 orbitals.
Hybridization of Atomic Orbitals Hybrid orbitals on carbon overlap with 1 s orbitals on hydrogen.
Hybridization of Atomic Orbitals Determine the number and type of hybrid orbitals necessary to rationalize the bonding in PCl 5. Step 1: Draw the Lewis structure:
Hybridization of Atomic Orbitals Determine the number and type of hybrid orbitals necessary to rationalize the bonding in PCl 5. Step 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 Five hybrid orbitals required
Hybridization of Atomic Orbitals Determine the number and type of hybrid orbitals necessary to rationalize the bonding in PCl 5. Step 3: Draw the ground-state orbital diagram for the central atom. Step 4: Maximize the number of unpaired electrons by promotion.
Hybridization of Atomic Orbitals Determine the number and type of hybrid orbitals necessary to rationalize the bonding in PCl 5. Step 5: Combine the necessary number of atomic orbitals to generate the required number of hybrid orbitals. Step 6: Place electrons in the hybrid orbitals, putting one electron in each orbital before pairing any electrons.
Hybridization of Atomic Orbitals Hybrid orbitals on phosphorus overlap with 3 p orbitals on chlorine.
Hybridization of Atomic Orbitals
9. 5 Hybridization in Molecules Containing Multiple Bonds Valence bond theory and hybridization can be used to describe the bonding in molecules containing double and triple bonds. Each carbon has three electron domains: 2 single bonds 1 double bond Expect sp 2 hybridization ethylene (C 2 H 4)
Hybridization in Molecules Containing Multiple Bonds Maximize unpaired electrons on carbon by promotion: Hybridize the required number of atomic orbitals (one for each electron domain on carbon)
Hybridization in Molecules Containing Multiple Bonds Hybridization scheme for a the carbon atoms in ethylene: One unhybridized atomic 2 p orbital gives rise to multiple bonds ethylene (C 2 H 4) Three equivalent sp 2 hybrid orbitals explain three bonds around carbon
Hybridization in Molecules Containing Multiple Bonds A sigma (σ) bond forms when sp 2 hybrid orbitals on the C atoms overlap.
Hybridization in Molecules Containing Multiple Bonds A sigma (σ) bond occurs when electron density is concentrated directly along the internuclear axis
Hybridization in Molecules Containing Multiple Bonds The ethylene molecule contains five sigma bonds: 1 between the two carbon atoms (sp 2 and sp 2 overlap) 4 between the C and H atoms (sp 2 and 1 s overlap) ethylene (C 2 H 4) Unhybridized 2 p orbitals, NOT involved in sigma bonding (no overlap)
Hybridization in Molecules Containing Multiple Bonds 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.
Hybridization in Molecules Containing Multiple Bonds A pi (π) bond occurs when the regions of electron density are concentrated above and below the plane of the molecule. internuclear axis
Hybridization in Molecules Containing Multiple Bonds A sigma and a pi bond together constitute a double bond. ethylene (C 2 H 4)
Hybridization in Molecules Containing Multiple Bonds Sigma bonds exhibit free rotation around the bond axis. All three Lewis structures represent the same molecule.
Hybridization in Molecules Containing Multiple Bonds Pi bonds restrict free rotation around the bond axis. cis-1, 2 -dichloroethylene trans-1, 2 -dichloroethylene There are two isomers of 1, 2 -dichloroethylene
Hybridization in Molecules Containing Multiple Bonds The acetylene molecule is linear with sp hybridized carbons. acetylene (C 2 H 2) Promotion of an electron maximizes the number of unpaired electrons:
Hybridization in Molecules Containing Multiple Bonds The acetylene molecule is linear with sp hybridized carbons. 3 sigma bonds: 1 between the two carbon atoms (sp and sp) 2 between the C and H atoms (sp and 1 s) acetylene (C 2 H 2) 2 pi bonds 2 between the two carbon atoms (2 p and 2 p) The 2 s orbital and one of the 2 p orbitals then mix to form two sp hybrid orbitals:
Hybridization in Molecules Containing Multiple Bonds The acetylene molecule is linear with sp hybridized carbons. acetylene (C 2 H 2) Two unhybridized atomic 2 p orbitals gives rise to 2 pi bonds Two equivalent sp hybrid orbitals give rise to 2 sigma bonds
Hybridization in Molecules Containing Multiple Bonds Formation of the C−C sigma bond in acetylene: acetylene (C 2 H 2) 3 sigma bonds: 1 between the two carbon atoms (sp and sp) 2 between the C and H atoms (sp and 1 s)
Hybridization in Molecules Containing Multiple Bonds Formation of the C−C pi bond in acetylene: acetylene (C 2 H 2) 2 pi bonds 2 between the two carbon atoms (2 p and 2 p)
9. 6 Molecular Orbital Theory Lewis structures and valence bond theory fail to predict some important properties of molecules. Paramagnetic species are attracted to magnet fields. Paramagnetism is a result of a molecules electron configuration. Species that contain one or more unpaired electrons are paramagnetic. The Lewis structure for O 2 shows no unpaired electrons. O 2 exhibits paramagnetism.
Molecular Orbital Theory Lewis structures and valence bond theory fail to predict some important properties of molecules. Diamagnetic species are weakly repelled by magnetic fields. Species that contain paired electrons are diamagnetic. Nitrogen, N 2, is diamagnetic.
Molecular Orbital Theory Another bonding theory is needed to describe the paramagnetism of O 2. In molecular orbital theory, the atomic orbitals combine to form new orbitals that are the “property” of the entire molecule. The news orbitals are called molecular orbitals. Molecular orbitals have characteristics similar to atomic and hybrid orbitals: specific shapes specific energies accommodate a maximum of 2 electrons each electron filling follows the Pauli exclusion principle the number of molecular orbitals obtained equals the number of orbitals combined
Molecular Orbital Theory H 2 is the simplest homonuclear diatomic molecule. Valence bond theory: H 2 forms when from the overlap of the 1 s orbitals. Molecular obrital theory: H 2 forms when the 1 s orbitals combine to give molecular orbitals. Molecular orbitals result from the constructive and destructive combination of atomic orbitals.
Molecular Orbital Theory Constructive combination of the two 1 s 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.
Molecular Orbital Theory Destructive combination of the two 1 s 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 Orbital Theory Molecular orbitals that lie along the internuclear axis are referred to as σ molecular orbitals. Examples: σ1 s bonding molecular orbital from the combination of two 1 s orbitals σ*1 s antibonding molecular orbital from the combination of two 1 s orbitals The asterisk distinguishes an antibonding molecular orbital from a bonding orbital.
Molecular Orbital Theory 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 are higher in energy than the isolated atomic orbitals.
Molecular Orbital Theory The bond order indicates how stable a molecule is The higher the bond order, the more stable a molecule is.
Molecular Orbital Theory H 2 According to molecular orbital theory, H 2 is a stable molecule.
Molecular Orbital Theory He 2 According to molecular orbital theory, He 2 is NOT a stable molecule and does not exist.
Molecular Orbital Theory 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
Molecular Orbital Theory p atomic orbitals also form molecular orbitals by both constructive and destructive combination. π molecular orbitals – electron density above and below the internuclear axis.
Molecular Orbital Theory Molecular orbitals resulting from the combination of p atomic orbitals are higher than those resulting from the combination of s atomic orbitals. This order of orbital energies assumes 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 O 2, Fe 2, Ne 2.
Molecular Orbital Theory Molecular orbitals resulting from the combination of p atomic orbitals are higher than those resulting from the combination of s atomic orbitals. This order of orbital energies assumes some mixing of s and p orbitals. This arrangement of orbitals is found in Li 2, B 2, C 2 and N 2.
Molecular Orbital Theory 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 Theory Molecular orbital diagrams for second-period homonuclear diatomic molecules.
9. 7 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 Strength: predict the shape of many molecules and polyatomic ions. Weakness: fails to explain why bonds form (based on Lewis theory).
Bonding Theories and Descriptions of Molecules with Delocalized Bonding 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.
Bonding Theories and Descriptions of Molecules with Delocalized Bonding Molecular Orbital Theory Strength: accurately predict the magnetic and other properties of molecules Weakness: complex
Bonding Theories and Descriptions of Molecules with Delocalized Bonding Some molecules are best described using a combination of models. Benzene, C 6 H 6, is represented with two resonance structures: The π bonds in benzene are delocalized -- spread out over the entire molecule.
Bonding Theories and Descriptions of Molecules with Delocalized Bonding The π bonds in benzene are delocalized -- spread out over the entire molecule.
9 Chapter Summary: Key Points Molecular Geometry The VSEPR Model Bond Angles Molecular Polarity Valence Bond Theory Hybridization Multiple Bonding Molecular Orbital Theory Bonding Molecular Orbitals Antibonding Molecular Orbitals Bond Order Molecular Orbital Diagrams Delocalized Bonding
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