Post on 21-Jan-2016
description
PRINCIPLES OF CHEMISTRY I
CHEM 1211
CHAPTER 10
DR. AUGUSTINE OFORI AGYEMANAssistant professor of chemistryDepartment of natural sciences
Clayton state university
CHAPTER 10
MOLECULAR STRUCTURE AND
BONDING THEORIES
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
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
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
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
Five Electron Pairs (5 Electron Domains)
- Separated by 90o and 120o
- This electron pair arrangement is said to be trigonal bipyramidal
ELECTRON PAIRS
Six Electron Pairs (6 Electron Domains)
- Separated by 90o
- This electron pair arrangement is said to be octahedral
ELECTRON PAIRS
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
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
Molecules with Two VSEPR Electron Groups
- These molecules are linear
ExamplesCO2 (carbon dioxide)
HCN (hydrogen cyanide)BeCl2 (beryllium chloride)
VSEPR MODEL
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
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
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
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
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
- 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
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
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
MOLECULAR POLARITY
O C OCO2
Linear, symmetrical and nonpolar
H2O O
H H
Nonlinear and polar
HCN H C N
Linear but polar
- 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
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
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
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
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
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
- 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
- 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
- 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
- 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
- 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)
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
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)
- 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)
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
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