Chapter one - gen. obc

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Chemistry 481(01) Winter 2003
Instructor Dr. Upali Siriwardane e-mail
upali_at_chem.latech.edu Office CTH 311 Phone
257-4941 Office Hours 800-900 a.m.
1100-1200 a.m. M, W 800-1000 a.m. Tu, Th,
F. April 8, 2003 Test 1 (Chapters 4,  5   6)
May 1, 2003 Test 2 (Chapters  7, 8-12  13)
May 22, 2003 Test 3 (Chapters. 14, 16 17)
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Chapter 3. Molecular Structure and Bonding

• Bonding Theories
• 1. Lewis Theory
• 2. VSEPR Theory
• 3. Valence Bond theory (with hybridization)
• 4. Molecular Orbital Theory (molecualr
  orbitals)

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Lewis symbols
Lewis symbols of second period elements
Li Be B C N O F Ne
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What is a Lewis Structure (electron-dot formula)
of a Molecule?

• A molecular formulas with dots around atomic
  symbols representing the valence electrons
• All atoms will have eight (octet) of electrons
  (duet for H) if the molecule is to be stable.

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Single covalent bonds
H
H
H
C
H
H
H
Do atoms (except H) have octets?
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Lewis structures

• This is a simple system to help keep track of
  electrons around atoms, ions and molecules -
  invented by G.N. Lewis.
• If you know the number of electrons in the
  valence-shell of an atom, writing Lewis
  structures is easy.
• Lewis structures are used primarily for s- and
  p-block elements.

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How do you get the Lewis Structure from Molecular
formula?
Add all valence electrons and get valence
electron pairs Pick the central atom Largest
atom normally or atom forming most bonds Connect
central atom to terminal atoms Fill octet to all
atoms (duet to hydrogen)
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Types of electrons

• Bonding pairs
• Two electrons that are shared between two atoms.
  A covalent bond.
• Unshared (nonbonding ) pairs
• A pair of electrons that are not shared between
  two atoms. Lone pairs or nonbonding electrons.

Unshared pair
oo
H Cl
oo
oo
oo
Bonding pair
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2 bond pairs 2 x 2 4
2 lone pairs 2 x 2 4

Total 8 4 pairs Bond pairs an electron
pair shared by two atom in a bond. E.g. two pairs
between O--H in water. Lone pair an electron
pair found solely on a single atom. E.g. two
pairs found on the O atom at the top and the
bottom.
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Lewis structure and multiple bonds
This arrangement needs too many electrons.
O C O
How about making some double bonds?
That works!
is a double bond, the same as 4 electrons
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Multiple bonds

• So how do we know that multiple bonds really
  exist?
• The bond energies and lengths differ!
• Bond Bond Length Bond energy
• type order pm kJ/mol
• C C 1 154 347
• C C 2 134 615
• C C 3 120 812

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Formal Charges

• Formal charge
• valence electrons - ½ bonding -non bonding
  electrons
• There are two possible Lewis structures for a
  molecule. Each has the same number of bonds. We
  can determine which is better by determining
  which has the least formal charge. It takes
  energy to get a separation of charge in the
  molecule
• (as indicated by the formal charge) so the
  structure with the least formal charge should be
  lower in energy and thereby be the better Lewis
  structure

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Which Lewis Structure?
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What is Resonance Structures?

• Several Lewis structures that need to be drawn
  for molecules with double bonds
• One Lewis structure alone would not describe the
  bond lengths of the real molecule.
• E.g. CO32-, NO3-, NO2-, SO3

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Resonance structures

• Sometimes we can have two or more equivalent
  Lewis structures for a molecule.
• O - S O O S - O
• They both - satisfy the octet rule
• - have the same number of bonds
• - have the same types of bonds
• Which is right?

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Resonance structures

• They both are!
• O - S O O S - O
• O S O
• This results in an average of 1.5 bonds between
  each S and O.

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What is the Bond Order between P-O?
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CO32- ion
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NO3- ion
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SO3 Molecule
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NO2- ion
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Resonance structures

• Benzene, C6H6, is another example of a compound
  for which resonance structure must be written.
• All of the bonds are the same length.

or
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Exceptions to the octet rule

• Not all compounds obey the octet rule.
• Three types of exceptions
• Species with more than eight electrons around an
  atom.
• Species with fewer than eight electrons around an
  atom.
• Species with an odd total number of electrons.

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Atoms with more than eight electrons

• Except for species that contain hydrogen, this is
  the most common type of exception.
• For elements in the third period and beyond, the
  d orbitals can become involved in bonding.
• Examples
• 5 electron pairs around P in PF5
• 5 electron pairs around S in SF4
• 6 electron pairs around S in SF6

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Atoms with fewer than eight electrons

• Beryllium and boron will both form compounds
  where they have less than 8 electrons around them.

FBF F




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Atoms with fewer than eight electrons

• Electron deficient. Species other than hydrogen
  and helium that have fewer than 8 valence
  electrons.
• They are typically very reactive species.

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Molecules with Odd Number of Electrons

• NO2

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What is VSEPR Theory

• Valence Shell Electron Pair Repulsion
• This theory assumes that the molecular structure
  is determined by the lone pair and bond pair
  electron repulsion around the central atom

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What Geometry is Possible around Central Atom?

• What is Electronic or Basic Structure?
• Arrangement of electron pairs around the central
  atom is called the electronic or basic structure
• What is Molecular Structure?
• Arrangement of atoms around the central atom is
  called the molecular structure

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Possible Molecular Geometry

• Linear (180)
• Trigonal Planar (120)
• T-shape (90, 180)
• Tetrahedral (109)
• Square palnar ( 90, 180)
• Sea-saw (90, 120, 180)
• Trigonal bipyramid (90, 120, 180)
• Octahedral (90, 180)

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Molecular Structure from VSEPRTheory

• H2O
• Bent or angular
• NH3
• Pyramidal
• CO2
• Linear

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Molecular Structure from VSEPRTheory

• SF6
• Octahedral
• PCl5
• Trigonal bipyramidal
• XeF4
• Square planar

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What is a Polar Molecule?

• Molecules with unbalanced electrical charges
• Molecules with a dipole moment
• Molecules without a dipole moment are called
  non-polar molecules

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How do you Pick Polar Molecules?

• Get the molecular structure from VSEPR theory
• From c (electronegativity) difference of bonds
  see whether they are polar-covalent.
• If the molecule have polar-covalent bond, check
  whether they cancel from a symmetric arrangement.
• If not molecule is polar

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What is Valence Bond Theory

• Describes bonding in molecule using atomic
  orbital
• orbital of one atom occupy the same region with a
  orbital from another atom
• total number of electrons in both orbital is
  equal to two

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Orbital Hybridization
New orbitals are constructed from pre-existing s,
p, and d-orbitals hybrid orbitals
1. Hybridize the CENTRAL ATOM ONLY (others as
needed)
2. Only use valence shell electrons
3. The number of hybrid orbitals formed number
of atomic orbitals used
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sp3 Hybridization
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sp3 hybridization for H2O
Needed to form 2 sigma bonds and 2 lone pairs
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sp2 Hybridization
BF3 - trigonal planar according to VSEPR Theory
(incomplete octet exception)
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For BF3, we need 3 hybrid orbitals, so 3 atomic
orbitals are required as follows (s p p)
sp2
Needed to form 3 sigma bonds
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sp2 Hybridization
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sp Hybridization
BeH2 - linear according to VSEPR Theory
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For BeH2, we need 2 hybrid orbitals, so 2 atomic
orbitals are required as follows (s p) sp
Needed to form 2 sigma bonds
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sp Hybridization
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H - 1s orbitals
sp hybrids
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sp3d Hybridization
PF5 - trigonal bipyramidal according to VSEPR
Theory
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For PF5, we need 5 hybrid orbitals, so 5 atomic
orbitals are required as follows (s p p p
d) sp3d
Needed to form 5 sigma bonds
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sp3d2 Hybridization
SF6 - octahedral according to VSEPR Theory
Lewis Structure
Electron Pair Geometry
Molecular Geometry
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For SF6, we need 6 hybrid orbitals, so 6 atomic
orbitals are required as follows (s p p p
d d) sp3d2
3 unhybridized d-orbitals
Six sp3d2 hybridized orbitals for S-F bonds
Needed to form 6 sigma bonds
Isolated S atom
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Electron pairarrangement
GeometricFigure
Example
Linear2 pairs sp hybrids
Trig. Plan. 3 pairs sp2 hybrids
tetrahedral4 pairs sp3 hybrids
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Electron pairarrangement
GeometricFigure
Example
Trig. Bipyram.5 pairs sp3d hybrids
Octahedral3 pairs sp3d2 hybrids
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Multiple Bonds
Sigma (?) bonds end-to-end overlap
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Pi (?) bond side-by-side overlap
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Refinements to molecular geometry
lone pair- lone pair gt lone pair-bond pair
gt bond pair-bond pair
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What is hybridization?

• Mixing of atomic orbitals on the central atom
• a hybrid orbital could over lap with another
  atomic orbital or hybrid orbital of another atom
• possible hybridizations sp, sp2, sp3, sp3d,
  sp3d2

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How do you tell the hybridization of a central
atom?

• Get the Lewis structure of the molecule
• Look at the number of electron pairs on the
  central atom. Note double, triple bonds are
  counted as single electron pairs.
• Follow the following chart

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Kinds of hybrid orbitals

• Hybrid geometry of orbital
• sp linear 2
• sp2 trigonal planar 3
• sp3 tetrahedral 4
• sp3d trigonal bipyramid 5
• sp3d2 octahedral 6

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Hybridization involving d orbitals

• Co(NH3)63 ion Co3 Ar 3d6
• Co3 Ar 3d6 4s0 4p0
• Concentrating the 3d electrons in the dxy, dxz,
  and dyz orbitals in this subshell gives the
  following electron configuration hybridization is
  sp3d2

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Multiple Bonding

• Double bonds In the case of ethylene, HCCH, we
  have the Lewis structure with sp2 hybridization
  with each carbon having an unhybridized o orbital
• Triple bonds In the case of acetylene, HC?CH, we
  have the Lewis structure with sp3 hybridization
  with each carbon having an unhybridized o orbital
  

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What are p, s and d bonds

• s bonds
• single bond resulting from head to head overlap
  of atomic orbital
• p bond
• double and triple bond resulting from lateral or
  side way overlap of p atomic orbitals
• d bond
• double, triple and quadruple bond resulting from
  lateral or side way overlap of d atomic orbitals

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Consequences of Electronegativity in Chemical
Bonds

• 0-0.6 0.6 - 1.5 gt 1.5
  
• covalent Polar-covalent ionic
• bond bond bond
• Polarity of a molecular substance is measured as
  dipole moment.
• If molecule has a dipole moment it means may have
  polar covalent bonds
• If polar covalent bonds are symmetrical, they may
  lead to zero dipole moment

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Symmetry Adopted Linear Combinations (SALC) of
Atomic Orbitals
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Energies of Molecular Orbitals
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Molecular Orbital Theory

• Molecular orbitals are obtained by combining the
  atomic orbitals on the atoms in the molecule.

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Bonding and Anti-bobding Molecular Orbital
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Basic Rules of Molecular Orbital Theory

• The MO Theory has five basic rules
• The number of molecular orbitals the number of
  atomic orbitals combined
• Of the two MO's, one is a bonding orbital (lower
  energy) and one is an anti-bonding orbital
  (higher energy)
• Electrons enter the lowest orbital available
• The maximum of electrons in an orbital is 2
  (Pauli Exclusion Principle)
• Electrons spread out before pairing up (Hund's
  Rule)

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Bond Order

• Calculating Bond Order

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Homo Nuclear Diatomic Molecules

• Period 1 Diatomic Molecules H2 and He2

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Homo Nuclear Diatomic Molecules

• Period 2 Diatomic Molecules and Li2 and Be2

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Homo Nuclear Diatomic Molecules
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Molecualr Orbital diagram for O2, F2 and Ne2
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Molecualr Orbital diagram for B2, C2 and N2
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Homonuclear Diatomic Molecules 2nd Period
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Electronic Configuration of molecules

• When writing the electron configuration of an
  atom, we usually list the orbitals in the order
  in which they fill.
• Pb Xe 6s2 4f14 5d10 6p2
• We can write the electron configuration of a
  molecule by doing the same thing. Concentrating
  only on the valence orbitals, we write the
  electron configuration of O2 as follows.
• O2 (2s) 2(2s) 2 (2p) 4 (2p) 2

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Electronic Configuration and bond oder
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Hetero Nuclear Diatomic Molecules
HF molecule
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Hetero Nuclear Diatomic Molecules
Carbon monoxide CO
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Molecular Orbital Diagram of H2O
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Conjugated and aromatic molecules

• trans-1,3-Butadiene
• Allyl radical
• Cyclopropenium ion C3H3
• Cyclobutadiene
• Cyclopentadiene
• Benzene
• C7H7 (tropyllium) and C8H82

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trans-1,3-Butadiene
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Allyl radical
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Cyclopropenium ion C3H3
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Cyclopentadiene
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Benzene
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Aromatic Rings
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The Isolobal Analogy

• Different groups of atoms can give rise to
  similar shaped fragments.

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Metallic Bonding

• Metals are held together by delocalized bonds
  formed from the atomic orbitals of all the atoms
  in the lattice.
• The idea that the molecular orbitals of the band
  of energy levels are spread or delocalized over
  the atoms of the piece of metal accounts for
  bonding in metallic solids.

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Bonding Models for Metals

• Band Theory of Bonding in Solids
• Bonding in solids such as metals, insulators and
  semiconductors may be understood most effectively
  by an expansion of simple MO theory to
  assemblages of scores of atoms

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Linear Combination of Atomic Orbitals
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Linear Combination of Atomic Orbitals
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Types of Materials

• A conductor (which is usually a metal) is a solid
  with a partially full band
• An insulator is a solid with a full band and a
  large band gap
• A semiconductor is a solid with a full band and a
  small band gap
• Element Band Gap C 5.47 eVSi 1.12
  eVGe 0.66 eVSn 0 eV

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Superconductors

• When Onnes cooled mercury to 4.15K, the
  resistivity suddenly dropped to zero

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The Meissner Effect

• Superconductors show perfect diamagnetism.
• Meissner and Oschenfeld discovered that a
  superconducting material cooled below its
  critical temperature in a magnetic field excluded
  the magnetic flux.Results in levitation of the
  magnet in a magnetic field.

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Theory of Superconduction

• BCS theory was proposed by J. Bardeen, L. Cooper
  and J. R. Schrieffer. BCS suggests the formation
  of so-called 'Cooper pairs'

Cooper pair formation - electron-phonon
interaction the electron is attracted to the
positive charge density (red glow) created by
the first electron distorting the lattice around
itself.
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High Temperature Superconduction

• BCS theory predicted a theoretical maximum to Tc
  of around 30-40K. Above this, thermal energy
  would cause electron-phonon interactions of an
  energy too high to allow formation of or
  sustain Cooper pairs.
• 1986 saw the discovery of high temperature
  superconductors which broke this limit (the
  highest known today is in excess of 150K) - it is
  in debate as to what mechanism prevails at
  higher temperatures, as BCS cannot account for
  this.