Transcript Chapter 9
Chapter
9 Molecular Geometry: Shape Determines Function
Chapter Outline
• • • • • • •
9.1 Molecular Shape 9.2 Valence-Shell Electron-Pair Repulsion Theory (VSEPR) 9.3 Polar Bonds and Polar Molecules 9.4 Valence Bond Theory 9.5 Shape and Interactions with Large Molecules 9.6 Chirality and Molecular Recognition 9.7 Molecular Orbital Theory © 2014 W. W. Norton Co., Inc.
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Molecular Shape
•
Chemical/physical properties are related to molecular shape.
•
Lewis structures
•
Show atoms and bonds, but not spatial orientations (3D).
•
Molecular models
•
Show orientations and bond angles; help us understand physicochemical properties.
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Lewis Structures vs. Models © 2014 W. W. Norton Co., Inc.
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Molecular Shape (cont.)
•
Bond angle:
•
Angle (in degrees) defined by lines joining the centers of two atoms to the center of a third atom to which they are covalently bonded
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Not always predictable from Lewis structures © 2014 W. W. Norton Co., Inc.
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Chapter Outline
• • • • • • •
9.1 Molecular Shape 9.2 Valence-Shell Electron-Pair Repulsion
• •
Theory (VSEPR) Central Atoms with No Lone Pairs Central Atoms with Lone Pairs 9.3 Polar Bonds and Polar Molecules 9.4 Valence Bond Theory 9.5 Shape and Interactions with Large Molecules 9.6 Chirality and Molecular Recognition 9.7 Molecular Orbital Theory © 2014 W. W. Norton Co., Inc.
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Valence-Shell Electron-Pair Repulsion Theory (VSEPR)
• • •
VSEPR Theory:
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A model predicting that the arrangement of valence electron pairs around a central atom minimizes repulsion to produce the lowest-energy orientation.
Electron-pair geometry:
•
Three-dimensional arrangement of bonding e – pairs and lone pairs electrons about a central atom Molecular geometry:
•
3-dimensional arrangement of atoms in a molecule.
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VSEPR: Electron-Pair Geometry
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To determine electron-pair geometry:
• •
Draw Lewis structure (see Chapter 8).
From Lewis structure, determine steric number (SN) :
• •
Determine optimal spatial arrangement of electron pairs (bonding + nonbonding) to minimize repulsion.
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Molecular Geometry: Central Atom with No Lone Pairs
• •
Molecular geometry = electron-pair geometry
• • • •
Determine steric number (SN):
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SN = 2 (two atoms bonded to central atom)
•
geometry
linear SN = 3 (three atoms bonded to central atom)
•
geometry
trigonal planar SN = 4
tetrahedral SN = 5
trigonal bipyramidal SN = 6
octahedral © 2014 W. W. Norton Co., Inc.
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Central Atoms with No Lone Pairs © 2014 W. W. Norton Co., Inc.
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Central Atoms with No Lone Pairs (cont.) © 2014 W. W. Norton Co., Inc.
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Geometric Forms
• Examples: • CO 2 BF 3 CCl 4 PF 5 SF 6
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Practice: Molecular Geometry (No Lone Pairs)
• •
Determine the molecular geometry of: a) H 2 CO (C is central atom) b) CH 4
- Collect and Organize: We are given molecular formulas and asked to predict their molecular geometry.
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Practice: Molecular Geometry (No Lone Pairs)
•
Determine the molecular geometry of: a) H 2 CO (C is central atom)
•
b) CH 4
- Analyze: We can use the periodic table to determine the number of valence electrons for each atom. From the molecular formula and valence electrons, we can draw Lewis structures. From the Lewis structures we can determine the SN. From the SN we can predict the electron pair geometry. Since there are no lone pairs, the electron pair geometry is the same as the molecular geometry.
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Practice: Molecular Geometry (No Lone Pairs)
• •
Determine the molecular geometry of: a) H 2 CO (C is central atom) b) CH 4 -Solve:
H 2 CO C is the central atom, with single bonds to each H atom and a double bond to O, SN = 3. Molecular geometry = trigonal planar .
CH 4 C is central atom, with single bonds to each H atom, SN = 4. Molecular geometry = tetrahedral.
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Practice: Molecular Geometry (No Lone Pairs)
• •
Determine the molecular geometry of: a) H 2 CO (C is central atom) b) CH 4
- Think About It: The molecular geometries are consistent with VSEPR theory for 3 and 4 electron clouds. It is worth noting that the C atom always has 4 bonds, but a double bond counts as only one electron cloud, resulting in a trigonal planar geometry.
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Central Atoms with Lone Pairs
• •
Molecular geometry
electron-pair geometry
•
Replace bonding pair(s) with lone pair(s).
Example: SO 2
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(SN = 3) Three electron pairs (2 bonding + 1 lone pair) © 2014 W. W. Norton Co., Inc.
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Central Atoms w/ Lone Pairs (cont.)
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Bond angles less than predicted
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Electron pair repulsion!
• • •
Lone pair –lone pair = greatest repulsion.
Lone pair –bonding pair Bonding pair –bonding pair = least repulsion.
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Multiple bonds > single bonds 9 - 18 © 2014 W. W. Norton Co., Inc.
Molecular Geometry: SN = 4
Note: bond angles decrease as # of lone pairs increases.
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Molecular Geometry: SN = 4 (cont.)
Two lone pairs = greater repulsion, decreased bond angle.
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Molecular Geometry: SN = 5
Note: lone pairs occupy equatorial positions.
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Molecular Geometry: SN = 6
Note: bond angles = 90 (geometries w/ more than 2 lone pairs are possible.)
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Practice: Molecular Geometry
•
What are the molecular geometries of the ions: SCN – and NO 2 – ?
- Collect and Organize: geometries.
We are given the molecular formulas for two polyatomic ions and asked to predict the molecular
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Practice: Molecular Geometry
•
What are the molecular geometries of the ions: SCN – and NO 2 – ?
-Analyze: We can use the periodic table to determine the number of valence electrons for each atom. From the molecular formula and valence electrons, we can draw Lewis structures. From the Lewis structures, we can determine the SN. From the SN, we can predict the electron-pair geometry. Making note of the number of bonding pairs and lone pairs, we can identify the molecular geometry.
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Practice: Molecular Geometry
•
What are the molecular geometries of the ions: SCN – Solve:
SCN –
and NO 2 – ?
As the least electronegative element and the one with the greatest bonding capacity, C is the central atom. Although there are several possible resonance structures, they all have C with SN = 2 and no lone pairs. Molecular geometry = linear .
NO 2 – With N as the central atom, the Lewis structure has N with SN = 3 (two bonding pairs and one lone pair). Again, although there are two possible resonance structures, they both have the same SN value. Molecular geometry = bent, with bond angle <120 due to the extra repulsive energy of the lone pair on N .
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Practice: Molecular Geometry
•
What are the molecular geometries of the ions: SCN – and NO 2 – ?
-Think About It: number.
It is worth noting that both molecular structures have two bonding electron clouds, but different molecular geometries due to the differences in steric
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Chapter Outline
• • • • • • •
9.1 Molecular Shape 9.2 Valence-Shell Electron-Pair Repulsion Theory (VSEPR) 9.3 Polar Bonds and Polar Molecules
• •
What Makes a Molecule Polar?
Dipole Moments 9.4 Valence Bond Theory 9.5 Shape and Interactions with Large Molecules 9.6 Chirality and Molecular Recognition 9.7 Molecular Orbital Theory © 2014 W. W. Norton Co., Inc.
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Polar Bonds and Polar Molecules
• •
Requirements for polar molecule: 1. Polar bonds (i.e. covalent bond between atoms with ΔEN).
2. Nonuniform distribution of polar bonds.
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Polar Molecules (cont.)
• •
Bond dipole:
•
Separation of electrical charge created when atoms with different EN form a covalent bond.
Polar molecule:
•
Vectors of bond dipoles whose sum > zero.
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Measuring Polarity
•
Dipole moment (
) – a quantitative expression of the polarity of a molecule.
•
Units = debyes (D); 1 D = 3.34
×
10 –30 C·m ) © 2014 W. W. Norton Co., Inc.
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Chapter Outline
• • • • • • •
9.1 Molecular Shape 9.2 Valence-Shell Electron-Pair Repulsion Theory (VSEPR) 9.3 Polar Bonds and Polar Molecules 9.4 Valence Bond Theory
• •
Bonds from Orbital Overlap Hybridization (sp 3 , sp 2 , sp, sp 3 d, sp 3
d
2 ) 9.5 Shape and Interactions with Large Molecules 9.6 Chirality and Molecular Recognition 9.7 Molecular Orbital Theory © 2014 W. W. Norton Co., Inc.
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Atomic Orbitals and Bonds
• •
Valence Bond Theory (Linus Pauling)
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Quantum mechanics-based model
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Covalent bond = overlap of half-filled orbitals Sigma (
) bond:
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Covalent bond in which the highest electron density lies between the two atoms along the bond axis.
Overlap of 1s orbitals
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Hybridization:
sp 3
Orbitals
•
Hybridization – the mixing of atomic orbitals to generate new sets of orbitals that form covalent bonds with other atoms.
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Tetrahedral Sigma Bonds
•
Tetrahedral orientation of sp 3 hybridized orbitals = tetrahedral molecular geometry
Overlap of 1s with
sp
3 orbitals
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Other
sp 3
Hybrid Examples
Note: lone pairs (non bonding)
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Trigonal Planar: Hybridization
sp
2
Unhybridized p orbitals form double bonds.
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sp2 Hybridization (cont.)
•
pi (
) bond : a covalent bond in which electron density is greatest around —not along —the bonding axis. © 2014 W. W. Norton Co., Inc.
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Linear:
sp
Hybridization
Form triple bond (one and two π bonds).
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Trigonal Bipyramidal: Hybridization
sp
3
d
Formed by mixing one s, one d, and three p orbitals. Example: PF 5 – five sigma bonds
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Octahedral:
sp
3
d
2 Hybridization
Formed by mixing one s, two d, and three p orbitals.
Example: SF 6 – six sigma bonds
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Practice: Hybrid Orbitals
What are the hybridizations of the central atoms of the ions: SCN – and NO 2 – ?
-
Collect and Organize:
Note that these are the same molecules for which we determined molecular geometry back in section 9.2. Using Lewis structures and VSEPR, we determined that the electronic geometry around the central carbon in SCN electronic geometry around the central N atom in NO 2 – – was linear (SN = 2), and the was trigonal planar (SN = 3).
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Practice: Hybrid Orbitals
What are the hybridizations of the central atoms of the ions: SCN – and NO 2 – ?
- Analyze: geometry.
From the steric number of the central atoms and valence bond theory, we can determine the hybridization around the central atom based on electron-pair
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Practice: Hybrid Orbitals
What are the hybridizations of the central atoms of the ions: SCN – and NO 2 – ?
Solve:
SCN – is linear (SN = 2), so the hybridization must be
sp
.
NO 2 – is trigonal planar (SN=3), so the hybridization must be
sp 2 .
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Chapter Outline
• • • • • • •
9.1 Molecular Shape 9.2 Valence-Shell Electron-Pair Repulsion Theory (VSEPR) 9.3 Polar Bonds and Polar Molecules 9.4 Valence Bond Theory 9.5 Shape and Interactions with Large Molecules
• •
Shape and Molecular Recognition Delocalized Electrons 9.6 Chirality and Molecular Recognition 9.7 Molecular Orbital Theory © 2014 W. W. Norton Co., Inc.
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Molecular Recognition
•
Molecular recognition:
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The process by which molecules interact with other molecules in living tissues to produce a biological effect.
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Example: ethylene (ripening agent)
All atoms in same plane.
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Delocalization of Electrons
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Delocalization: spreading of electrons in alternating single and double bonds over three or more atoms in a molecule
a) b)
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Aromatic Compounds
•
Aromatic compound:
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A cyclic, planar compound with delocalized
electrons above and below the plane of the molecule.
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Example: polycyclic aromatic hydrocarbons (PAH)
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Planar shape may allow intercalation in DNA © 2014 W. W. Norton Co., Inc.
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Chapter Outline
• • • • • • •
9.1 Molecular Shape 9.2 Valence-Shell Electron-Pair Repulsion Theory (VSEPR) 9.3 Polar Bonds and Polar Molecules 9.4 Valence Bond Theory 9.5 Shape and Interactions with Large Molecules 9.6 Chirality and Molecular Recognition
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Optical Isomerism 9.7 Molecular Orbital Theory © 2014 W. W. Norton Co., Inc.
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Isomerism
•
Isomers – compounds with same molecular formula but with atoms arranged differently in three-dimensional space.
Optical isomers
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Optical Isomers = Chirality
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Chirality – property of a molecule that is not superimposable on its mirror image.
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Chapter Outline
• • • • • • •
9.1 Molecular Shape 9.2 Valence-Shell Electron-Pair Repulsion Theory (VSEPR) 9.3 Polar Bonds and Polar Molecules 9.4 Valence Bond Theory 9.5 Shape and Interactions with Large Molecules 9.6 Chirality and Molecular Recognition 9.7 Molecular Orbital Theory
• • • •
Molecular Orbitals of Hydrogen and Helium MOs of Homonuclear Diatomic Molecules MOs of Heteronuclear Diatomic Molecules MOs of N2+ and Spectra of Auroras © 2014 W. W. Norton Co., Inc.
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Molecular Orbital (MO) Theory
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Based on mixing of atomic orbitals of similar shapes and energies to form molecular orbitals (MOs) that belong to the molecule as a whole
•
The number of MOs formed is equal to the number of atomic orbitals combined.
•
MOs represent discrete energy states; orbitals spread out over entire molecule.
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Types of Molecular Orbitals
• •
Bonding orbitals:
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Region of increased electron density between nuclear centers that hold atoms together
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Are lower in energy (more stable) than atomic orbitals from which they are formed Antibonding orbitals:
•
Region of electron density that destabilize the molecule because they do not increase electron density between nuclear centers
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Less stable than atomic orbitals from which they are formed 9 - 55 © 2014 W. W. Norton Co., Inc.
Molecular Orbital Diagrams
• • •
MO diagram:
•
Energy-level diagram for molecular showing the relative energies and electron occupancy of the MOs for a molecule.
Sigma ( σ) bond:
•
Covalent bond with the highest electron density along the bond axis.
Pi ( π) bond:
•
Bond formed by mixing of atomic orbitals not oriented along the bonding axis in a molecule.
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Molecular Orbital Diagram: H
2 The two 1s orbitals mix to yield two sigma MOs (1 bonding/1 antibonding). 9 -57 © 2014 W. W. Norton Co., Inc.
Bond Order and Stability
Bond order = 1/2 (# bonding e – – # antibonding e – )
Bond order in H 2 (Stable) – = ½
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Bond order in He 2 (Not stable) = 0
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MO Guidelines
1.
2.
3.
4.
5.
The total # of MOs = the # of AOs orbitals mixed.
Orbitals with similar energy/shape mix more effectively than do those of different energy/shape.
Orbitals of different n (different sizes/energies) result in less effective mixing.
A MO can accommodate two electrons.
Electrons fill MO diagrams according to Hund’s rule.
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MO Scheme for O
2
• •
Electron configuration for O 2 : ( σ 2s ) 2 ( σ 2s * ) 2 ( σ 2p ) 2 ( π 2p ) 4 ( π 2p * ) 2 Bond order = ½ (8 – 4) = 2
•
O 2 has two bonds
•
O 2 has two unpaired electrons in π 2p * © 2014 W. W. Norton Co., Inc.
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MO Scheme for N
2
• • Electron configuration for N 2 : ( σ 2
s
) 2 ( σ 2
s
* ) 2 ( σ 2p ) 2 ( π 2p ) 4 Bond order = ½ (8 – 2) = 3 N 2 has three bonds.
N 2 has no unpaired electrons.
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Paramagnetic
vs.
Diamagnetic
• •
Paramagnetism:
•
Atoms or molecules having unpaired electrons are attracted to magnetic fields.
•
Example: O 2 Diamagnetism:
•
Atoms or molecules having all paired electrons are repelled by magnetic fields.
•
Example: N 2 © 2014 W. W. Norton Co., Inc.
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MO for NO
• • •
Z
eff alters MO diagram; atomic orbitals for O are lower in energy.
Odd electron in atom.
π* 2p , closer in energy to the 2p atomic orbitals of N Bond order = ½ (8 – 3) = 2.5
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MO for N 2 , N 2 + : Emission Spectra
Crimson red
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Blue-violet
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