Zumdahl’s Chapter 8
Chemical Bonding:
Stealing and Sharing
of electrons
Chapter Contents
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Types of Bonds
Electronegativity
Polarity & Dipoles
Ions
Binary Ionic
Compounds
• Polar Covalency
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The Covalent Bond
Bond Energies
Local & Global Model
G.N. Lewis Structures
Octet or not Octet
Resonance fudge
Valence Shell Electron
Pair Repulsion Theory
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)
• Metal Bonding (macroscopic wavefunction)
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
•  < ~1 “covalent” while  > ~2 “ionic”
• “Polar covalent” for  between 1 and 2.
• Not yet ionic, but still significantly polar.
Pauling’s Definition
• DAB = [ DAADBB ]½ + 96.5(  )²
H
2.1
Pauling
Electronegativities
He
Li
1.0
Be
1.5
B
2.0
C
2.5
N
3.0
O
3.5
F
4.0
Ne
Na
0.9
Mg
1.2
Al
1.5
Si
1.8
P
2.1
S
2.5
Cl
3.0
Ar
soon
K
0.8
Ca
1.0
Ga
1.6
Ge
1.8
As
2.0
Se
2.4
Br
2.8
Kr
3.0
Rb
0.8
Sr
1.0
In
1.7
Sn
1.8
Sb
1.9
Te
2.1
I
2.5
Xe
2.7
Cs
0.7
Ba
0.9
Tl
1.8
Pb
1.9
Bi
1.9
Po
2.0
At
2.2
Rn
Dipole Moment, , & Polarity
•   0 implies charge separation in bonds.
• Charge separation means bond polarity:
• Product of Q times RAB, bond distance.
• Molecular dipole is vector sum of bond
dipoles.
• Symmetries cause vector cancellations!
• Bond polarity is necessary but not sufficient
cause of molecular polarity.
Ions
• Electron configuration follows removal of
electrons with highest n from parent atom.
• Effective charge 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 anion’s negative charge.
Recipe for a Crystal
• Gasify both the electropositive and the
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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
Madlung Constants
• Lattice energies can be calculated by crystal
geometries. Nearest neighbors attract, next
nearest neighbors repel, and so on forever.
• Madlung constants (k = 1.5– 4.2) sum those all
up as a scale factor for binary ionic potentials.
• While O2– can’t exist by itself (O– has a positive electron
affinity), it is stabilized in Mg2+O2–, say, by its very
negative lattice energy, 4 that of NaCl due to Q1Q2,
with which it shares a crystal structure (same k).
Polar Covalency
• Most compounds have heteronuclear bonds,
so   0 and the bond is polar.
• But few will have  so high (>~2) that their
bonds are truly ionic.
• Slight inequality in electron sharing makes
them “polar covalent.”
• % ionic = 100% [ measured / ionic ]
• Not to be confused with ionic % in .
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)
• Stretches ; helps minimize e– kinetic energy
• Reduces e– density near nucleus, reducing the
e––e– repulsions there.
Two H atoms NOT Bonding
Notice the
extent of
overlap.
H2 with Bonding turned ON
Deeper
Overlap
“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.
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 >  D broken , products are more
stable than reactants, and we expect exothermic
reaction, heat evolved as potential lowers.
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.
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 molecule’s atoms.
• Arrange multiple bonds (multi-electron pairs)
between atoms to minimize “formal charges.”
Lewis Structure Rules
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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.
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 FC=0 for all atoms.
• But  FC = ion’s charge, so some won’t be 0.
• Negative FC goes to highest electronegativity.
Egregious Example, NOCl
• Nitrosyl chloride, “NOCl,” is a horrific nonaqueous solvent for some food processing
applications; O=N–Cl seems more likely.
• Try to find a good NOCl Lewis Structure:
• # = 5+6+7 = 18 valence electrons
• N–O –Cl uses 4, so use 7 lone pairs (14 e–)
• :N:–:O:–:Cl:: fails to deliver a N octet
• :N:=O:–:Cl:: gives everyone an octet; OK?
NOCl Bombs in Formal Charge!
• Since N, O, and Cl are expecting 5, 6, and 7
valence electrons, respectively,
• :N:=O:–:Cl:: then shows “formal charges” of
• –1 +1 0
• putting the positive charge on the most
electronegative atom?!?
• It’s more likely we had the skeleton wrong,
so let’s try ONCl.
ONCl, a Better Lewisite
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O–N–Cl still needs 7 lone pairs
::O:–N:–:Cl:: still needs N octet help
::O=N:–:Cl:: 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.
Octet Trumps Formal Charge
• If you can’t get both, sacrifice good formal
charges to securing an octet. For example,
• CO has 4+6=10 valence electrons
• C–O needs 4 more lone pairs
• :C:–:O: 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.
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.
Meta Arsenic Acid
• (HO)AsO2 has As as its central atom.
• # = 5+3(6)+1 = 24 valence electrons
• ::O = As = O:: distributes all 24
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and puts octets on all oxygens
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::O – H
(and the duet on hydrogen)
• with zero formal charges everywhere,
• but requires 10 not 8 electrons on As.
• HNO3 exists too but has a N octet and FCs!
The Impoverished
nd
2
Row
• But perchloric acid, (HO)ClO3
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::O – H
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• ::O = Cl = O::
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:O:
• has no fluorine (octet) analogue.
• No 2d orbitals & F can’t take +3 formal charge!
Newest Rare Gas Compound
• HArF (sounds like a hairball), recently
synthesized, leaves only He and Ne as truly
noble gases.
• But neither has d orbitals.
• It’s Ar’s d orbitals that permit H–Ar:::–F:::
with its 10 electrons around Ar.
• Don’t hold your breath for Ne compounds.
Resonance
• Often several different legitimate Lewis
Structures are possible. NO3– ion is all of
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: O:: _
:O: _
: O:: _
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and
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and
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• ::O=N–O:::
:::O–N–O:::
:::O–N=O::
• with FCs of +1 (N), –1 (O–), and 0 (O=)
• representing delocalized electrons, greater
stability, and uniform NO bond orders of 1.33!
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 e––e– 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 e–s point.
# of directional electron pairs:
• 1: linear A-B molecule • 4: tetrahedral
• 109.47°
• 2: collinear
• 180°
• 3: trigonal
• 120°
• 5: trigonal bipyramid
• 120° and 90°
• 6: octahedral
• 90°
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 aren’t involved
in describing molecular shapes.
• A Xn Em used to count bonding directions
(n) and lone pairs (m). A is central atom.
Examples of 5 Directional Pairs
• AX5
• Trigonal bipyramid
• PCl5
• AX4E
• See Saw
• SF4
• AX3E2
• T-shaped
• BrF3
• AX2E3
• Linear
• I3–
Central Atoms
• … aren’t terminal atoms in the molecule.
• Molecules may have several central atoms.
• Geometry is determined at each by VSEPR.
Acetic
acid
Mint
Histamine
Nitroglycerin