Zumdahl Chapter 8

1
Zumdahl Chapter 8

• Chemical Bonding
• Stealing and Sharing
• of electrons

2
Chapter Contents

• The Covalent Bond
• Bond Energies
• Local Global Model
• G.N. Lewis Structures
• Octet or not Octet
• Resonance
• Valence Shell Electron Pair Repulsion Theory

• Types of Bonds
• Electronegativity
• Polarity Dipoles
• Ions
• Binary Ionic Compounds The Born-Haber Cycle
• Polar Covalency

3
Types of Chemical Bonds

• Ionic Bonds
• Large differences in electronegativity hold atoms
  together by Coulombic potentials.
• Radically polar poorly-shared binding electrons
• Covalent Bonds
• Near equally-shared, weakly polar binding pair
• Dative Bonds (donor-receptor model)
• Metallic Bonding (delocalized electrons)

4
Bond Energy The energy required to break the
bond

• Ionic Bonds
• Closely packed, oppositely charged ions.
• Energy of interaction between two ions

The calculation uses Coulombs Law in the form
(Q1 Q2)
E (2.31 x 10-19 J nm)
r
For NaCl
(1) (-1 )
-8.37 x 10-19 J
E (2.31 x 10-19 J nm)
0.276 nm
5
Electronegativity, ?

• Measure of ionization potential and electron
  affinity. The power to hold and attract.
• Empirically, deviation of bond energies, DHX,
  from the geometrical average of DHH and DXX.
• ?? ?A ?B determines ionicity of AB
• ?? lt 1.7 covalent while ?? gt 1.7 ionic
• Polar covalent for ?? between 1 and 2.
• Not yet ionic, but still significantly polar.

6
Paulings Values
7
Dipole Moment, ?, Polarity

• ?? ? 0 implies charge separation in bonds.
• Charge separation means bond polarity.
• Molecular dipole is vector sum of bond dipoles.
• Symmetries cause vector cancellations!
• Bond polarity is necessary but not sufficient
  cause of molecular polarity.

8
Ions

• Electron configuration follows removal of
  electrons with highest n from parent atom.
• Zeff governs size of ion
• Cations are much smaller than parent atom.
• Anions are much larger but
• High charge anions are unstable except in
  compounds because 2nd extra electron is repelled
  by 1st anions negative charge.

9
Recipe for a Crystal

• Gasify both the electropositive and the
  electronegative elements. ( ?Hphase )
• Dissociate both to atoms. ( DXX )
• Make the cation and anion. ( IP EA )
• Condense the crystal from the infinitely remote
  ions. ( Ulattice )
• Add and components ? ?Hformation

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The Born-Haber Cycle
?H1 ?H2 ?H3 ?H4 ?H5 ?Hof
The Lattice Energy is defined as the reverse of
?H5
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Born-Haber Processes
?H1
Sublimation
M (s) ?H1 M (g)
?H2
Dissociation
X2 (g) ?H2 2 X (g)
?H3
Ionization
M (g) ?H3 M (g) e-
?H4
Electron Affinity
2 X (g) 2 e- 2 X- (g) ?H4
?H5
Formation of Ionic Solid
M (g) X- (g) MX (s) ?H5
?Hof
Enthalpy of formation
M (s) 1/2 X2 (g) MX (s)
12
Polar Covalency

• Most compounds have heteronuclear bonds, so ?? ?
  0 and the bond is polar.
• But few will have ?? so high (gt2) that their
  bonds are truly ionic.
• Slight inequality in electron sharing makes them
  polar covalent.

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Covalent Bonds

• Electron pairs shared (more or less equally)
  between bonding partners.
• Build-up of e density between nuclei.
• Counteracts nuclear (proton) repulsion
• Increases attractions (to neighboring protons)
• Reduces e density near nucleus, reducing the
  ee repulsions there.

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Attractive and repulsive forces in hydrogen
Red Repulsive
Blue Attractive
15
Hydrogen Bond Formation
0
- 436
16
Lennard-Jones PE Plot
Equilibrium Internuclear distance
(Bond Length)
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Definition of the Chemical Bond
When the forces acting between atoms are such as
to lead to the formation of an aggregate with a
lower energy than the separate atoms from which
it is formed.
-- The Nature of the Chemical Bond (1939)
Linus Pauling
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Bond Energies

• While it takes a definite energy to break a
  particular bond in a unique molecule, bond
  energies refers to some weighted average of a
  particular bond in all the molecules in which it
  is found.
• As such, they are correct for no molecule.
• Still energies can be estimated from ?D, for
  bonds changed during a reaction.

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Bond Enthalpy

• ?H ?D since ?H measures differences in the
  formation of compounds from elements while ?D
  shows the differences in the destruction of
  compounds. Opposites.
• So ?H ? D broken ? D formed
• If ? D formed gt ? D broken products are more
  stable than reactants, and we expect exothermic
  reaction, heat evolved as potential lowers.

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Covalency is Local

• In the Local Electron model, electrons are
• Shared in pairs between adjacent atoms,
• unshared entirely as lone pairs,
• or hidden deep and uninvolved as core.
• This ignores the reality of global electrons,
  free to roam over more than one atom pair.
• G.N. Lewis had a patch for this difficulty with
  the LE model, but we need a better model.

21
Gilbert Newton Lewis

• Developed a method to manipulate valence
  electrons to satisfy local atomic needs.
• Devised Lewis Structures and rules for the
  placement of electron pairs about them to
• Put at least one pair between all bonded atoms.
• Complete rare gas electronic configurations
  about all the molecules atoms.
• Arrange multiple bonds (multi-electron pairs)
  between atoms to minimize formal charges.

22
Lewis Structure Rules

• Sum all valence and ionic electrons N.
• Pick an atomic skeletal structure.
• Place two of N between all bonded atoms.
• Distribute remainder as lone pairs to achieve an
  octet (only duet for hydrogen).
• Minimize formal charge by stealing lone pairs to
  make additional (multiple) bond pairs. If FC is
  dumb, pick another skeleton.

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Formal Charges

• Shared electrons count toward BOTH atoms octets.
• But shared electrons are divided equally between
  their bonded atoms for FC.
• Lone pairs count fully in FC of an atom.
• FC valence electrons sum of above.
• Best value is FC0 for all atoms.
• But ? FC ions charge, so some wont be 0.
• Negative FC goes to highest electronegativity.

24
Egregious Example, NOCl

• Nitrosyl chloride, NOCl, is a horrific
  non-aqueous solvent for some food processing
  applications ONCl seems more likely.
• Try to find a good NOCl Lewis Structure
• 567 18 valence electrons
• NO Cl uses 4, so use 7 lone pairs (14 e)
• NOCl fails to deliver a N octet
• NOCl gives everyone an octet OK?

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NOCl Bombs in Formal Charge!

• Since N, O, and Cl are expecting 5, 6, and 7
  valence electrons, respectively,
• NOCl then shows formal charges of
• 1 1 0
• putting the positive charge on the most
  electronegative atom?!?
• Its more likely we had the skeleton wrong, so
  lets try ONCl.

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ONCl, a Better Lewis Structure

• ONCl still needs 7 lone pairs
• ONCl still needs N octet help
• ONCl satisfies octets gives FC of
• 0 0 0 Perfect!
• So Lewis Structures permit us to correct an
  incorrect molecular formula into one truly
  reflecting the geometry, ClNO.

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Octet Trumps Formal Charge

• If you cant get both, sacrifice good formal
  charges to securing an octet. For example,
• CO has 4610 valence electrons
• CO needs 4 more lone pairs
• CO gives nobody an octet
• C ? O isoelectronic with N?N, but FC are
• 1 1 not what one would expect at all!
• but that inobvious polarity is correct.

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Breaking the Octet

• Rare Gas Configuration is 8 valence electrons
  (save for He) ending in ns2 np6.
• Atoms beyond row 2 have d orbitals which empower
  them to adopt more than 8.
• Often central, such atoms can surround
  themselves with typically 12 but sometimes 14 or
  more valence electrons in their molecules.

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Exceptions to the Octet Rule
We have already seen one of these cases
Molecules with an expanded octet
e.g. PF5
The other exceptions are
Incomplete octets
Such as BeF2
The small Be atom cannot accommodate a large
formal charge.
Other Group 2A elements form ionic bonds.
30
Incomplete Octets
Another common Incomplete Octet are the elements
of Group IIIA
e.g. BF3
Other Group IIIA elements also exhibit this
incomplete octet.
31
Odd-Electron Species
Some molecules have an odd number of electrons
e.g. NO
And NO2
32
The Concept of Resonance
Consider the Lewis Structure of ozone, O3
36 18 electrons 9 pairs
The best Lewis structure has one double bond and
one single bond
1
-1
0
This would lead one to expect that ozone contains
two different bond lengths a shorter (stronger)
double bond and a longer (weaker) single bond.
33
Measured Bond Lengths in O3
Experimental measurement of bond lengths shows
that the bonds in ozone are identical
34
Resonance of O3
The experimental results suggest that the double
bond is shared equally in the two positions
AND
It therefore takes both Lewis structures to
adequately represent the O3 molecule
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Resonance of O3
The different structures are called resonance
structures and are generally shown with a
double-headed arrow between them
The important points to remember about resonance
forms are .
The molecule is not rapidly oscillating between
two different forms
There is only one form of the ozone molecule.The
bond lengths between the oxygens are intermediate
between characteristic single and double bond
lengths between a pair of oxygens.
We draw two Lewis structures (in this case)
because a single structure is insufficient to
describe the real structure.
36
Lewis Structure of the NO3- ion
Consider the Lewis Structure of nitrate, NO3-
36 5 1 24 electrons 12 pairs
The best Lewis structure has one double bond and
two single bonds
Experimental measurement of bond lengths shows
that the bonds in the nitrate ion are identical
37
Resonance of the NO3- ion
It takes three resonance structures to adequately
represent the nitrate ion
38
Resonance in Benzene C6H6
Measurements show all benzene bonds to be
identical.
39
Two Simple Theories of Covalent Bonding

• Valence Shell Electron Pair Repulsion Theory
• Commonly designated as VSEPR
• R. J. Gillespie in the 1950s
• Valence Bond Theory
• Involves the use of hybridized atomic orbitals
• L. Pauling in the 1930s 40s

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Valence Shell Electron Pair Repulsion Theory
(VSEPR)

• Lewis Structures show e pairs on between
  atoms.
• VSEPR uses those to predict shapes as a
  consequence of ee repulsions such that
• Lone pairs repel more strongly than bond pairs.
• All pairs seek geometries that distance them.
• Pairs in multiple bonds are treated as if they
  are a single bond pair for directional purposes.
• Nuclei follow blindly where bonding es point.

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Valence Shell Electron Pair Repulsion Theory

• Describes the geometric arrangement of electron
  pairs around a central atom in terms of electron
  repulsions.

Molecule will assume the geometry that has the
minimum repulsion
42
Molecular Shapes

• Determined only by the direction of the bonding
  pairs since only they terminate in atoms.
• Lone pairs still dictate where bonding pairs go,
  but lone pair directions arent involved in
  describing molecular shapes.
• A Xn Em used to count bonding directions (n) and
  lone pairs (m). A is central atom.

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VSEPR Theory

• Two bonding pairs around the central atom.

LINEAR
AB2
Example BeCl2
44
VSEPR Theory

• Three bonding pairs around the central atom.

TRIGONAL PLANAR
AB3
Example BF3
45
VSEPR Theory

• Four bonding pairs around the central atom.

TETRAHEDRAL
AB4
Examples CH4 CCl4 NH4
46
VSEPR Theory

• Five bonding pairs around the central atom.

Trigonal bipyramidal
AB5
Example PCl5
47
VSEPR Theory

• Six bonding pairs around the central atom.

OCTAHEDRAL
AB6
Example SF6
48
The six electronic arrangements
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Effect of Lone Pairs
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AB3E
AB2E2
..
AB4
O
H
..
104.5o
H
CH4, methane
NH3, ammonia
H2O, water
..
AB2E
lone pair electrons
119.5o
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Molecular Shapes
AB2 Linear
AB3 Trigonal planar
AB2E Angular or Bent
AB4 Tetrahedral
AB3E Trigonal pyramidal
AB2E2 Angular or Bent
AB2E3 Linear
AB5 Trigonal bipyramidal
AB4E Irregular tetrahedral (see saw)
AB3E2 T-shaped
AB5E Square pyramidal
AB4E2 Square planar
AB6 Octahedral