Chapter one gen. obc

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Chemistry 281(01) Winter 2003-2004
Instructor Dr. Upali Siriwardane Class Meetings
930-1045 a.m. Labs Tu 200 - 615 p.m.
e-mail upali_at_chem.latech.edu Office CTH 311
Phone [phone] Office Hours M, W, F 800-900,
1100-1200a.m. Tu, Th 1000 a.m.-1200 noon
January 7,  2004(Test 1)Chapter
1,2,3,4 February 2,      2004 (Test 2) Chapter
5,6,7February 18,    2004 (Test 3) Chapter
,8,9,10March 1,    2004 (Make Up Test)
930-1045 a.m. CTH 322  
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Chapter 3. Covalent Bonding

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

• Octet Rule All elements except hydrogen (
  hydrogen have a duet of electrons) have octet of
  electrons once they from ions and covalent
  compounds.

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Noble gas configuration

• The noble gases are noted for
• their chemical stability and
• existence as monatomic
• molecules.
• Except for helium,
• they share a common
• electron configuration
• that is very stable.
• This configuration has 8 valence-shell electrons.
  
• All other elements reacts to achieve Noble Gas
  Electron Configurations.

valence e- He
2 Ne 8 Ar 8 Kr 8 Xe
8 Rn 8
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The octet rule

• Atoms are most stable if they have a filled or
  empty outer layer of electrons.
• Except for H and He, a filled layer contains 8
  electrons - an octet.
• Two atoms will
• gain or lose (ionic compounds)
• share (covalent compounds)
• Many atoms will
• share (metallic compounds)
• electrons to make a filled or empty outer layer.

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What changes take place during this process of
achieving closed shells?

• a) sharing leads to covalent bonds and
  molecules
• b) gain/loss of electrons lead to ionic bond
• c) Sharing with many atoms lead to metallic
  bonds

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Lewis Electron Dot symbols
Basic rules Draw the atomic symbol. Treat each
side as a box that can hold up to two
electrons. Count the electrons in the valence
shell. Start filling box - dont make pairs
unless you need to.
X
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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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What is the Lewis Structure?

• CO2
• NH3 (PH3)
  
  
• PCl3 (PF3, NCl3)

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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 electrons - 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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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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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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An example SO42-

• 1. Write a possible
• arrangement.
• 2. Total the electrons.
• 6 from S, 4 x 6 from O
• add 2 for charge
• total 32
• 3. Spread the electrons
• around.

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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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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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Which Molecules are Polar
H2O Bent or angular, polar-covalent bonds,
asymmetric molecule-polar NH3 Pyramidal,
polar-covalent bonds, asymmetric
molecule-polar CO2 Linear, polar-covalent bonds,
symmetric molecule-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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What are p and s 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 atomic orbitals

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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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Electronegativity

• The ability of an atom that is bonded to another
  atom or atoms to attract electrons to itself.
• It is related to ionization energy and electron
  affinity.
• It cannot be directly measured.
• The values are unitless since they are relative
  to each other.
• The values vary slightly from compound to
  compound but still provide useful qualitative
  predictions.

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Electronegativities
Electronegativity is a periodic property.
Electronegativity
Atomic number
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Electronegativity

• Relative ability of atoms to attract electrons of
  bond.

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Electronegativity

• The greater the difference in electronegativity
  between two bonded atoms, the more polar the
  bond.
• If the difference is great enough, electrons are
  transferred from the less electronegative atom to
  the more electronegative one.
• - Ionic bond.
• Only if the two atoms have exactly the same
  electronegativity will the bond be nonpolar.

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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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Electronegativity

• Determine the difference in electronegativity
  between the bonded atoms in the following
  compounds.
• KCl ENK 0.9 ENCl 2.8 D 1.9
• H2O ENH 2.2 ENo 3.5 D 1.3
• CH4 ENC 2.5 ENH 2.2 D 0.3
• NO2 ENN 3.1 ENO 3.5 D 0.4

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What is Ionic Bond?

• a) Bonding between elements with higher
  electronegtivity difference( gt 1.5)
• b) The electrostatic attraction of ions
• c) Ions attract to many other ions
• d) leads to stronger attractions
• e) leads to collections of ions or ionic
  solids
• f) To understand ionic bond one has to study
  factors that affecting the lattice formation

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What is Covalent Bond?

• Forces of attraction created by sharing a pair
  of electrons between two atoms
• A covalent bond creates a molecule
• What is a Coordinate Covalent Bond?When the
  shared electron pair is donated byone atom

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What are the differences betweenMolecular
Compounds and Ionic Compounds?

• Ionicsolids at room temperaturehave higher
  b.p., and m.p.Strong intermolecular forces
• Moleculargases and liquids at room
  temperaturehave lower b.p., and m.p.weak
  intermolecular forces