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