PRINCIPLES OF CHEMISTRY I CHEM 1211 CHAPTER 10
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PRINCIPLES OF CHEMISTRY I
CHEM 1211
CHAPTER 10
DR. AUGUSTINE OFORI AGYEMANAssistant professor of chemistryDepartment of natural sciences
Clayton state university
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CHAPTER 10
MOLECULAR STRUCTURE AND
BONDING THEORIES
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Valence Shell Electron Pair Repulsion (VSEPR) Theory
- Used to predict molecular structure (geometry)
- That is the three-dimensional arrangement of atoms within molecules
- The specific arrangements depend on the number of valence electron pairs present
Stearic Number= number of lone pairs on central atom
+ number of atoms bonded to central atom
ELECTRON PAIRS
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ELECTRON PAIRS
Two Electron Pairs (2 Electron Domains)
- Predicted to be as far apart as possible from one another
- Gives 180o angles to one another (opposite sides of the central atom)
- This electron pair arrangement is said to be linear
: :
180o
central atom
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Three Electron Pairs (3 Electron Domains)
- Predicted to be as far apart as possible
- Found at the corners of an equilateral triangle (separated by 120o angles)
- This electron pair arrangement is said to be trigonal planar
..
::
120o
ELECTRON PAIRS
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Four Electron Pairs (4 Electron Domains)
- Predicted to be as far apart as possible
- Found at the corners of a tetrahedron (separated by 109o angles)
- This electron pair arrangement is said to be tetrahedral
: :
:
:
109o
ELECTRON PAIRS
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Five Electron Pairs (5 Electron Domains)
- Separated by 90o and 120o
- This electron pair arrangement is said to be trigonal bipyramidal
ELECTRON PAIRS
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Six Electron Pairs (6 Electron Domains)
- Separated by 90o
- This electron pair arrangement is said to be octahedral
ELECTRON PAIRS
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VSEPR ELECTRON GROUPS
- Electrons present in a specific localized region about a central atom
Single bond - VSEPR electron group containing two electrons
- Represents one electron group
Double bond- VSEPR electron group containing four electrons
- Represents one electron group
VSEPR MODEL
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VSEPR ELECTRON GROUPS
Triple bond- VSEPR electron group containing six electrons
- Represents one electron group
Nonbonding Electron Pair Included when determining the number of electron groups
- Each pair represents one electron group
VSEPR MODEL
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Molecules with Two VSEPR Electron Groups
- These molecules are linear
ExamplesCO2 (carbon dioxide)
HCN (hydrogen cyanide)BeCl2 (beryllium chloride)
VSEPR MODEL
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Molecules with Three VSEPR Electron Groups
These molecules are
- trigonal planar (all electron groups are bonding) H2CO (formaldehyde)
- angular/bent/V-shaped (one electron group is nonbonding) SO2 (sulfur dioxide)
VSEPR MODEL
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Molecules with Four VSEPR Electron Groups
These molecules are
- tetrahedral (all electron groups are bonding) CH4 (methane)
- trigonal pyramidal (one electron group is nonbonding) NH3 (ammonia)
- angular/bent/V-shaped (two electron groups are nonbonding) H2O (water)
VSEPR MODEL
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Molecules With Five VSEPR Electron Groups
These molecules are
- trigonal bipyramidal (all electron groups are bonding) PCl5
- seesaw (one electron group is nonbonding) SF4
- T-shaped (two electron groups are nonbonding) ClF3
- linear (three electron groups are nonbonding) XeF2
VSEPR MODEL
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Molecules With Six VSEPR Electron Groups
These molecules are
- octahedral (all electron groups are bonding) SF6
- square pyramidal (one electron group is nonbonding) BrF5
- square planar (two electron groups are nonbonding) XeF4
VSEPR MODEL
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Molecules with More Than One Central Atom
- Determined by considering each central atom separately and combining the results
C2H2 (acetylene) and H2O2 (hydrogen peroxide)
VSEPR MODEL
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- Bond angles decrease as the number of nonbonding electron pairs increases
- Nonbonding electron pairs tend to exert greater repulsive forces on adjacent electron domains and compress bond
angles
- Multiple bonds also decrease bond angles (greater repulsive forces)
BOND ANGLES
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MOLECULAR POLARITY
Nonpolar Molecule - There is a symmetrical distribution of electron charge
Polar Molecule - There is an unsymmetrical distribution of electron charge
- Molecular polarity depends on bond polarity and molecular geometry
- Symmetrical molecules cancel polar bond effects
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MOLECULAR POLARITY
Diatomic Molecule
- polar bond results in polar molecule
- nonpolar bond results in nonpolar molecule
Generally- Molecules with lone pair of electrons on the
central atom are polar
- Molecules without lone pairs and with identical atoms on the central atom are nonpolar
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MOLECULAR POLARITY
O C OCO2
Linear, symmetrical and nonpolar
H2O O
H H
Nonlinear and polar
HCN H C N
Linear but polar
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- The assumption that atomic orbitals on an atom mix to form new orbitals of different shapes
- The process is called hybridization
- The number of hybrid orbitals equals the number of atomic orbitals mixed
HYBRID ORBITALS
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sp Hybrid Orbitals (sp hybridization)
- Two hybrid orbitals arranged at 180o involving one s orbital and one p orbital
- Each hybrid orbital has two lobes (one small and one large)
- Results in a linear arrangement of electron domains
BF2, BeCl2, CO2
HYBRID ORBITALS
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sp2 Hybrid Orbitals (sp2 hybridization)
- Three identical hybrid orbitals involving one s orbital and two p orbitals (at 120o)
- Three large lobes point towards the corners of an equilateral triangle
- Results in trigonal planar geometry
BF3
HYBRID ORBITALS
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sp3 Hybrid Orbitals (sp3 hybridization)
- Four identical hybrid orbitals involving one s orbital and three p orbitals (at 109o)
- Four large lobes point towards the vertex of a tetrahedron
- Results in a tetrahedral arrangement of electron domains
CH4
HYBRID ORBITALS
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sp3d Hybrid Orbitals (sp3d hybridization)
- Five hybrid orbitals arranged at 90o and 120o involving one s orbital, three p orbitals, and one d orbital
- Large lobes point towards the vertices of a trigonal bipyramid
PF5, SF4
HYBRID ORBITALS
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sp3d2 Hybrid Orbitals (sp3d2 hybridization)
- Six hybrid orbitals arranged at 90o involving one s orbital, three p orbitals, and two d orbital
- Large lobes point towards the vertices of an octahedron
SF6, ClF5
HYBRID ORBITALS
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- The overlap of two orbitals (electron density) along the internuclear axis (line connecting nuclei)
- The overlap of two s orbitals (H2)
- The overlap of an s and a p orbital (HCl)
- The overlap of two p orbitals (Cl2)
- The overlap of a p orbital and an sp hybrid orbital (BeF2)
SIGMA (σ) BONDS
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- Sideways overlap between two p orbitals (perpendicular to the internuclear axis)
- The regions overlapping lie above and below the internuclear axis
- Weaker than σ bonds (less total overlap)
- Most common in atoms having sp or sp2 hybridization (small atoms in period 2: C, N, O)
PI (π) BONDS
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- Single bonds are σ bonds (H2)
- Double bonds are comprised of one σ and one π bonds (C2H4)
- Triple bonds are comprised of one σ and two π bonds (C2H2 , N2)
MULTIPLE BONDS
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- Observed in resonance structures with π bonds
- Results in greater stability
- Responsible for colors of many organic compounds
Benzene (C6H6)- Delocalized π bonds among the six carbon atoms
- Bond lengths are identical and are between the C — C single bonds and the C = C double bonds
DELOCALIZATION
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- Most characteristics are the same as atomic orbitals
- Can hold a maximum of two electrons with opposite spins
- Atomic orbitals are associated with a single atom
- Molecular orbitals are associated with the entire molecule
- The number of molecular orbitals formed is equal to the number of atomic orbitals combined
MOLECULAR ORBITALS (MO)
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MOLECULAR ORBITALS (MO)E
ner
gy 1s 1s
H atom H atom
H2 molecule- Molecular orbital diagram for H2 (electron configuration is σ1s
2)- Two atomic orbitals overlap to form two molecular orbitals
- Energy level of one MO is lower than the atomic orbitals (filled with the two 1s electrons and is called bonding molecular orbital (σ1s)
- Energy level of the other MO is higher than the atomic orbitals (empty and is called antibonding molecular orbital (σ1s*)
- Electrons occupy lower energy and explains why hydrogen is diatomic
σ1s*
σ1s
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En
ergy 1s 1s
He atom He atom
He2 molecule
- Molecular orbital diagram for He2 (electron configuration is σ1s2 σ*1s
2)- Bonding molecular orbital (σ1s) is filled
- Antibonding molecular orbital (σ1s*) is also filled- Energy decrease in σ1s is offset by energy increase in σ1s*
- He2 is therefore unstable
σ1s*
σ1s
MOLECULAR ORBITALS (MO)
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- Determines the stability of covalent bonds
BOND ORDER
2
electronsgantibondinofnumberelectronsbondingofnumberOrderBond
- Single bonds: bond order is 1 - Double bonds: bond order is 2- Triple bonds: bond order is 3
- Bond order is 1 for H2 and 0 for He2 (no bond exists)
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Paramagnetism- Molecules with unpaired electrons are attracted into a
magnetic field
- Force of attraction increases with increasing number of unpaired electrons
Diamagnetism- Molecules without unpaired electrons are weakly repelled
from a magnetic field
MOLECULAR PROPERTIES
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Experimental Determination
- Weigh samples in the presence and absence of a magnetic field
- Paramagnetic substances will weigh more in the magnetic field
- Diamagnetic substances will weigh less in the magnetic field
MOLECULAR PROPERTIES