Chapter 9 Notes- Molecular Geometry and Bonding Theories

Chapter 9 Notes- Molecular Geometry and Bonding Theories

9.1 Molecular Shapes [p.342]

1. Tetrahedral- central atom surrounded by 4 atoms equidistant apart

a. Bond angle is 109.5˚

b. Bond length same for all four bonds

c. Generic formula AB4

d. Example- CH4 and CCl4

2. AB2 Molecular Shapes

a. Linear

i. Bond angle 180˚

ii. Example CO2

b. Bent

i. Bond angle NOT = 180˚

ii. Example H2O

3. AB3 Molecular Shapes

a. Trigonal Planar

i. Central atom lies in the same plane as the three atoms attached

ii. Example SO3

b. Trigonal Pyramidal

i. Central atom sits above three atoms that form a triangle

ii. Example NH3

c. Trigonal Bipyramidal

i. Central atom surrounded by 5 atoms, 3- 120˚ apart at “equator” and 1 located at the “north and south poles”.

ii. Example PCL5

d. Octahedral

i. Central atom surrounded by 6 atoms, 4- 90˚ apart at “equator” and 1 located at the “north and south poles”.

4. VSEPR model

a. Based on the fact that electrons repel each other

b. Best arrangement of electrons to minimize repulsive forces

9.2 VSEPR Model [p.344]

1. General Information

a. Electron Domains

i. Bonding Pair- a pair of electrons that defines a region in which electrons is most likely to be found.

ii. Nonbonding pair (lone pair) – electron domain located principally on one atom

iii. Multiple Bonds = One electron domain

b. Electron –Domain Geometry –

i. arrangement of electron domains about the central atom of a molecule or an ion that reduces repulsion

ii. Shape is dependent on the number of electron domains surrounding the central atom

iii. relates to VSEPR

c. Molecular Geometry –

i. arrangement of ONLY THE ATOMS in a molecule or ion (nonbonding pairs are not part of the description)

ii. Use electron-domain geometry to predict molecular geometry

d. Steps for predicting shape

i. Draw Lewis Structure

ii. Determine electron-domain geometry by arranging electron domains so that repulsion is minimized

iii. Use arrangement of bonded atoms to determine the molecular geometry

2. The Effect of Nonbonding Electrons and Multiple Bonds on Bond Angles

a. Electron domains for nonbonding electron pairs exert greater repulsive forces on adjacent electron domains and tend to compress the bond angles

b. Electrons domains for multiple bonds exert greater repulsive force on adjacent electron domains for single bonds

3. Molecules with Expanded Valence Shells

a. Molecules that have more than an octet are molecules that have a central atom from the third period or larger

b. Two electron domains

i. Trigonal Bipyramidal – (3- equatorial and 2-axial) produces 4 molecular geometries

a. Trigonal Bipyramidal ( 5- bonding domains and 0 nonbonding domains)

b. Seesaw ( 4-bonding domains and 1 nonbonding domains)

c. T-shaped (3-bonding domains and 2 nonbonding domains)

d. Linear (2-bonding domains and 3 nonbonding domains)

ii. Octahedral – (4- equatorial and 2-axial) produces 3 molecular geometries

a. Octahedral (6 bonding)

b. Square Pyramidal ( 5- bonding and 1 nonbonding)

c. Square Planar ( 4- bonding and 2 nonbonding)

4. Shapes of Larger Molecules

use same rules just apply individually to each central atom
see example on p.352 of Brown

9.3 Molecular Shape and Molecular Polarity [p. 353]

Bond polarity caused by a difference in electronegativites
Dipole moment = measure of the amount of charge separation in a molecule

a. Dependent on both the polarities of the individual atoms and the geometry of the molecule

b. Bond Dipole – dipole moment due to the two atoms in a bond

i. Vector quantities

ii. Overall dipole moment for molecule =’s the vector sum of the bond dipoles

1. 0= nonpolar molecule

2. Nonzero overall dipole moments - polar molecule

9.4 Covalent Bonding and Orbital Overlap [p.355]

1. Valance- bond Theory –marriage of Lewis’ notion of electron pairs and atomic orbitals

Outer orbitals of electron configuration overlap to form bond
At far distances the bonding energy = 0, as the atoms approach each other the energy in creases, until it gets to a minimum PE which corresponds to the bond strength, then the energy becomes repulsive

9.5 Hybrid Orbitals [p.357]

General Information

a. Hybrid orbitals- shape of original orbitals resulting from a mixing of atomic orbitals (Hybridization)

b. Total number of atomic orbitals remains constant, number of hybrid orbitals is dependent on the mixing of orbitals.

sp Hybrid Orbital

a. formed by hybridizing of an s and p orbital

b. produces 2 sp hybrid orbitals that have a uneven dumbbell shape

c. linear arrangement of electron domains implies sp hybridization

sp2 and sp3 Hybrid Orbitals

a. sp2 Hybrid

made from 1s and 2p orbitals
produces three equivalent orbitals
lie in the same plane 120 ˚ apart
unfilled 2p orbitals do not hybridize
example BF3

b. sp3 Hybrid

s orbital mixes with 3 p’s in the same subshell
produces 4 equivalent orbitals
3-D with bond angles of 109.5 ˚
Large lobe of orbital is positioned on the vertex of the tetrahedron
Hybridization Involving d Orbitals

a. Occurs in molecules containing more than 8 electrons around the central atom

b. sp3d

Forms a sp3d hybrid orbital from a s, p ,and d orbital in the same principal energy level
Larger lobe points toward the vertex of a Trigonal pyramid

c. sp3d2

Forms a sp3d2 hybrid orbital from a s, p ,and d orbital in the same principal energy level
Larger lobe points toward the vertex of a Octahedral
Hybrid Orbital Summary

a. Draw Lewis Structure

b. Determine electron-domain geometry

c. Specify hybrid orbitals needed

9.6 Multiple Bonds [p. 319]

Sigma (σ) bonds- Covalent bonds concentrate the electron density on the internuclear axis (line joining the two nuclei through the middle of the overlap region

a. Overlap of two s orbitals (H2)

b. Overlap of an s and p orbital (HCl)

c. Overlap of two p orbitals (Cl2)

d. Overlap of p and sp hybrid orbital (BeF2)

2. Pi (π) Bonds- overlap of two p orbitals oriented perpendicularly to the internuclear axis

a. Above and below the internuclear axis

b. No probability of finding electrons on the internuclear axis

c. Occurs less than σ bonds

d. Weaker than σ bonds

3. General rules

a. All single bonds are σ bonds

b. Double bonds are one σ bonds and one π bond

c. Triple bonds are one σ bonds and two π bonds

Resonance Structures, Delocalization and π Bonding

a. σ bonds and π bonds deal with localized bonding electrons (electrons associated totally with two atoms that form bonds)

b. delocalized π bonds explain the bonding in resonance structures in which the electrons are not associated with two bonding atoms

causes color changes in many organic molecules
must lie in same plane for optimal overlap (lends a certain rigidity to the molecule)
General Conclusions

a. Every pair of bonded atoms shares one or more pairs of electrons

b. σ bonds are localized

c. when multiple bonds occur 1- σ bonds and all additional bonds are π bonds

d. Molecules with two or more resonance structures can have π bonds that extend over more than two bonded atoms. (delocalized electrons)

9.7 Molecular Orbitals [p. 368]

1. Molecular Orbital Theory

a. Describes the electrons in molecules by using specific wave functions called Molecular orbitals (MO)

b. Molecular Orbitals

i. hold 2 electrons with opposite spin

ii. has a definite energy

iii. helps visualize the electron-density distribution of a molecule

iv. associated with the entire molecule, not single atom

2. The Hydrogen Molecule

a. Molecular orbitals - When have two atomic orbitals overlap, one high energy molecular orbital and one lower energy molecular energy form

i. Bonding Molecular Orbitals low energy molecular orbital

1. sigma (σ) molecular orbital (σ1s)

2. Result of the adding of two atomic orbitals

3. Concentrates electron density between the two nuclei

4. Electron is stable because it is attracted to both nuclei

ii. Anti-bonding molecular orbital high energy molecular orbital

1. Greatest density of electrons outside the region between nuclei

2. Repelled from bonding region

3. less stable

4. sigma (σ) molecular orbital (σ1s*)

iii. Energy-level diagrams – show the interaction between two 1s atomic orbitals. (see p.369)

3. Bond Order

a. Related to the stability of a bond

b. Equation

Bond order = ½ (no. of bonding electrons – no. of antibonding electrons)

c. bond order =1 single bond

d. bond order = 2 double

e. bond order = 0 no bond

f. bond order of ½ , 3/2, 5/2 corresponds to molecules with an odd number of valence electrons.

9.8 Second Row Diatomic Molecules [p.371]

1. General rules for formation of MO’s

a. Number of MOs formed equals the number of atomic orbitals combined

b. Atomic Orbitals combine best with orbitals in similar energy levels

c. Effective of combining is proportional to their overlap

d. Each MO can accommodate only two electrons (Pauli Exclusion)

e. MO’s must ½ orbitals first (Hund’s)

2. Molecular Orbitals for Li2 and Be2

a. Set-up energy level diagram similar to H2 just make to energy levels showing that 2s is higher in energy

b. General rule is that core electrons do not contribute significantly in molecule formation

c. Be2 energy level diagram indicates a bond order of 0 which matches the fact that Be2 does not exist

3. Molecular Orbitals from 2p Atomic Orbitals

a. Generates 2 pi (π) molecular orbitals for each axis

b. One p orbital overlaps “end-on” forming a σ molecular orbital which is lower in energy (more stable) and σ* MO is higher in energy (less stable) than π* MO

c. Other 2 axes overlap “sidewasys”

4. Electron Configurations for B2 through Ne2

a. The 2s orbitals are lower in energy than the 2p with the σ2s lower than the σ2s*

b. The “sideway” overlap of the 2p orbital forms a σ2P lower in energy than σ2P*

c. Both π2p and π2p* are doubly degenerate and sit in between the σ2s and σ2s*

5. Electron Configurations and Molecular Properties

a. Paramagnetism occurs in molecules that have one or more unpaired electrons that are attracted to magnetic fields

b. Substances with no unpaired electrons weakly repel magnetic fields (diamagnetism)

c. Way to determine which property a molecule has is by weighing it the presence of a magnetic field. If object is heavier, it is parmagnetic.

d. Multiple bonds have shorter bond lengths, greater bond enthalpies, and greater bond orders (very stable/ less reactive)

6. Heteronuclear Diatomic Molecules – set up similar to homonuclear diatomic molecules but an MO will have a greater contribution from the atomic orbital to which it is closer in energy