Transcript CH.9 - coolchemistrystuff

MOLECULAR GEOMETRY AND
BONDING THEORIES (CH.9)
By: Maggie Dang
9.1 Molecular Shapes
 The overall shape of a molecule is determined by
its bond angles, the angles made by the lines
joining the nuclei of the atoms in the molecule
 Molecules with a central atom A surrounded by n
atoms B, denoted ABn, adopt a number of
different geometric shapes, depending o n the
value of n and on the particular atoms involved
9.2 The VSEPR Model
 Valence-shell electron-pair repulsion (VSEPR) model
rationalizes molecular geometries in terms of the
repulsions between electron domains, which are regions
about a central atom in which electrons are likely to be
found.
 Bonding pairs of electrons are involved in making bonds
 Nonbonding pairs of electrons, also called lone pairs,
both create electron domains around an atom
Electron Domains
 Based on the VSEPR model, electron domains orient
themselves to minimize electrostatic repulsions and
to remain as far apart as possible
 Electron domains from nonbonding pairs exert slightly
greater repulsions than those from bonding pairs
 Electron domains from multiple bonds exert greater
repulsions than those from single bonds
 Electron-Domain Geometry: arrangement of electron
domains around a central atom
 Molecular Geometry: arrangement of atoms
Steps to Predicting Molecular
Geometries with the VSEPR Model
Sketch the Lewis structure of the molecule or ion
Count the total number of electron domains around the central
atom, and arrange them in the way that minimizes the repulsions
among them
Describe the molecular geometry in terms of the angular
arrangement of the bonded atoms
A double or triple bond is counted as one electron domain when
predicting geometry.
1.
2.
3.
4.




Ex: CO2 has C=O double bonds
When we apply the VSEPR model to CO2, each double bond counts as one
electron domain. The VSEPR model predicts that CO2 is linear.
Because multiple bonds count as one electron domain, the number of electron
domain can be counted as (# of electron domains)= (# of atoms bonded to the
central atom) + (# of nonbonding pairs on the central atom)
Refer to pgs 207-209 for molecular geometry tables
Example
 Using the VSEPR model, predict the
molecular geometries of O3.
The Effect of Nonbonding Electrons and
Multiple Bonds on Bond Angles
 Electron domains for non-bonding electron pairs exert
greater repulsive forces on adjacent electron domains
and thus tend to compress the bond angles
 Electron domains for multiple bonds exert a greater
repulsive force on adjacent
electron domains than do single bonds.
Molecules with Expanded Valence Shells
 These shapes generally contain axial and equatorial
positions
 When pointing toward an axial position, an electron domain is
situated 90° from three equatorial positions
 In equatorial position an electron domain is situated 120° from
the other two equatorial positions and 90° from the two axial
positions
 Repulsions between domains are much greater when
they are situated 90° from each other than when they are
at 120°.
 Variations of the trigonal bipyramidal shape show lone
electron pairs in the equatorial position
 Variations of the octahedral shape show lone electron
pairs in the axial positions
Molecules with More than
One Central Atom
 The VSEPR theory can be used for molecules with
more than one central atom
9.3 Polarity of Polyatomic
Molecules
 The dipole moment of a polyatomic molecule depends on the vector
sum of the dipole moment due to each individual bond, called the
bond dipole.
 Certain molecular shapes, such as linear AB2 and trigonal planar AB3,
assure that the bond dipoles cancel, leading to a dipole moment of
zero for the molecule.
 In other shapes such as bent AB2 and trigonal pyramidal AB3, the
bond dipoles do not cancel and the molecule will have a nonzero
dipole moment called a polar molecule
 One with a zero dipole moment is called nonpolar.
 Polarity is also used when talking about covalently bonded
molecules.
 If the molecule has only 2 different atoms, such as, HF or CCl4 you
can calculate the electronegativity difference and determine the
type of covalent bond (polar or non-polar).
Example
 Predict whether BrCl is polar or nonpolar.
 Chlorine is more electronegative than bromine.
Consequently, BrCl will be polar with chlorine
carrying the partial negative charge.
Polarity and Bond Type
Electronegativity Difference
Bonding Type
 <0.5
 Non-polar covalent
 0.5 – 1.9
 Polar covalent
 > 2.0
 ionic
9.4 Covalent Bonding and
Orbital Overlap
 Valence-bond theory is an extension of Lewis’s notion of
electron-pair bonds. In valence-bond theory, covalent
bonds are formed when atomic orbitals on neighboring
atoms overlap one another.
 The overlap region is a favorable one for the two
electrons because of their attraction to two nuclei.
 The greater the overlap between two orbitals, the
stronger the bond that is formed.
9.5 Hybrid Orbitals
 To extend the ideas of valence-bond theory to
polyatomic molecules, it is useful to envision the mixing
of s,p, and sometimes d orbitals to form hybrid orbitals.
 Hybrid orbitals can overlap with orbitals on other atoms
to make bonds, or they can accommodate nonbonding
pairs.
 The process of hybridization leads to hybrid orbitals that
are directed along certain definite directions
sp, sp2, and sp3 Hybrid
Orbitals
 One s orbital and one p orbital can hybridize to form two
equivalent sp hybrid orbitals
 For sp2, using BF₃ as an example, a 2s electron on the B atom
can be promoted to a vacant 2p orbital. Mixing the 2s and two
of the 2p orbitals yields three equivalent sp2 hybrid orbitals
 For sp3, using CH₄ as an example, it forms four equivalent
bonds with the four hydrogen atoms. This process results
from the mixing of the 2s and all three 2p atomic orbitals of
carbon to creat four equivalent sp3 hybrid orbitals.

See pgs 318-320 for diagrams
Hybridization Involving d
Orbitals
 Atoms in the third period and beyond can use d orbitals to form
hybrid orbitals. Mixing one s orbital, three p orbitals, and one d
orbital leads to five sp3d hybrid orbitals. These hybrid orbitals are
directed toward the vertices of a trigonal bipyramid.
 Similarly, mixing one s orbital, three p orbitals, and two d orbitals
gives six sp3d2 hybrid orbitals, which are directed toward the
vertices of an octahedron.
 The use of d orbitals in constructing hybrid orbitals corresponds to
the notion of an expanded valence shell.
Example
 Predict the hybridization of SF4

- There are five electron
domains around S, giving rise to the
trigonal bipyramidal electron-domain geometry. With an
expanded octet of 10 electrons, the use of a d orbital on the sulfur
is required. The trigonal bipyramidal electron-domain geometry
corresponds to SP3 d hybridization. One of the hybrid orbitals
that points in an equatorial direction contains a nonbonding pair
of electrons; the other four are used in forming the S-F bonds.
Steps to Predict Hybrid
Orbitals
Draw the Lewis structure for the molecule or ion.
2. Determine the electron-domain geometry using the
VSEPR model.
3. Specify the hybrid orbitals needed to accommodate the
electron pairs based on their geometric arrangement.
1.
9.6 Multiple Bonds
 Sigma bonds (σ): covalent bonds in which the electron density
lies along the line connecting the atoms
 Pi bonds (π): formed from the overlap of p orbitals that are
oriented perpendicular to the internuclear axis
 A double bond, such as that in C₂H₄, consists of one sigma
bond and one pi bond.
 A triple bond, such as that in C₂H₂, is composed of one sigma
bond and two pi bonds.
 Ex: H-H has one sigma bond, N₂ has one sigma plus two pi
bond
Delocalized Pi Bonding
 Every pair of bonded atoms shares one or more pairs of electrons.
In every bond at least one pair of electrons is localized in the space
between the atoms, in a sigma bond.
 The electrons in sigma bonds are localized in the region between
two bonded atoms and do not make a significant contribution to
the bonding between any other two atoms.
 When atoms share more than one pair of electrons, the additional
pairs are in pi bonds. The centers of charge density in a pi bond lie
above and below the bond axis
 Molecules with two or more resonance structures can have pi
bonds that extend over more than two bonded atoms. Electrons
in pi bonds that extend over more than two atoms are said to be
delocalized.
 Delocalized: Pi bonds are spread among several atoms
9.7 Molecular Orbitals
 Molecular orbital theory: another model used to describe
the bonding in molecules. In this model, the electrons
exist in allowed energy states call molecular orbitals
(MOs).
 A molecular orbital can be spread among all the atoms of
a molecule, can have a definite energy, and can hold two
electrons of opposite spin.
The Hydrogen Molecule

Whenever two atomic orbitals overlap, two moleular orbitals form. Thus, the overlap
of the 1s orbitals of two hydrogen atoms to form H₂ produces two MOs .

The lower-energy MO of H₂ concentrates electron density between the two hydrogen
nuclei and is called the bonding molecular orbital.

The higher-energy MO has very little electron density between the nuclei and is called
the antibonding molecular orbital.

The electron density in both the bonding and the antibonding molecular orbitals of H₂
is centered about the internuclear axis. MOs of this type are called sigma molecular
orbitals.

The bonding sigma molecular orbital of H₂ is labeled σ1s, indicating that the MO is formed
from two 1s orbitals.

The antibonding sigma molecular orbital of H₂ is labeled σ1s*, the asterisk denoting that
MO is antibonding.

The interaction between two 1s orbitals to form σ1s and σ1s* molecular orbitals can be represented
by an energy-level diagram(molecular orbital diagram). It shows the interacting atomic orbitals in
the left and right columns and the MOs in the middle column.
Bond Order
 The stability of a covalent bond is related to its bond order.

Bond order= ½ (# of bonding electrons - # of antibonding electrons)
 A bond order of 1 represents a single bond, a bond order of 2
represents a double bond, and a bond order of 3 represents a
triple bond.
 Because MO theory also treats molecules with an odd number
of electrons, bond orders of ½, 3/2, or 5/2 are possible.

Ex: H₂ has 2 bonding electrons and no antibonding ones. It has a bond
order of ½(2-0)=1.
Example
 What is the bond order of the O2+ ion?
 The O2+ ion has eight bonding electrons and
three antibonding ones. Thus, its bond order
is
 Bond order= ½ (8-3) = 2.5
9.8 Second-Row Diatomic
Molecules
 Second-row atoms have more than one atomic orbital
 The way we place electrons in the orbitals:
1.
The number of Mos formed equals the number of atomic
2.
3.
4.
5.
orbitals combined
Atomic orbitals combine most effectively with other atomic
orbitals of similar energy.
The effectiveness with which 2 atomic orbitals combine is
proportional to their overlap with one another. As the overlap
increases, the bonding MO is lowered in energy, and the
antibonding MO is raised in energy.
Each molecular orbital can accommodate, at most, two electrons,
with their spins paired (Pauli exclusion principle)
When Mos have the same energy, one electron enters each orbital
( with the same spin) before spin pairing occurs (Hund’s rule)
Molecular Orbitals for Li₂ and
Be₂
 Core electrons usually do not contribute
significantly to bonding in molecule
formation.
Molecular Orbitals from 2p Atomic
Orbitals
 The p orbitals that point directly at one another can form sigma
bonding and sigma antibonding MOs.
 The p orbitals that are oriented perpendicular to the internuclear
axis combine to form pi molecular orbitals.
 In diatomic molecules, the pi molecular orbitals occur as pairs of
degenerate (same energy) bonding and antibonding MOs.
 The σ2p bonding MO is expected to be lower in energy (more stable)
than the π2p bonding Mos because of larger orbital overlap. This
ordering is reversed in B₂, C₂, and N₂ because of interaction between
the 2s and 2p atomic orbitals.
Electron Configurations for B₂
Through Ne₂
 For B₂, C₂, and N₂, the σ2p MO is above the π2p
molecular orbitals in energy. For O₂, F₂, and Ne₂,
the σ2p MO is below the π2p molecular orbitals.
Electron Configurations and
Molecular Properties
 Molecules with 1 or more unpaired electrons are
attracted into a magnetic field. The more unpaired
electrons in a species, the stronger the force of
attraction. This type of magnetic behavior is called
paramagnetism.
 Substances with no unpaired electrons are weakly
repelled from a magnetic filed. This property is called
diamagnetism. It is a much weaker effect than
paramagnetism.