Chapter 8
Bonding
What is a Bond?
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A force that holds atoms together.
The properties of substances are largely
determined by the type of bonds that
hold them together.
There are three general types of bonds:
Ionic, covalent, and metallic
Metallic Bonds
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The easiest type to recognize!
Found in metals (Cu, Mg)
In metals, each atom is bonded to
several neighboring atoms,
The electrons move easily and account
for the typical properties of metals (luster
and conductivity)
Called a “sea of mobile electrons”
Ionic Bonding

Generally, a metal bonds with a nonmetal
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The electron is transferred
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Opposite charges hold the atoms
together.
Covalent Bonding
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Sharing of electrons between atoms
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Nonmetal bonded to nonmetal
Lewis Symbols
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Represent the valence electrons only
Why are we interested in the valence
electrons?
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Those are the ones involved in bonding!
Lewis (1875-1946) devised a way of
“tracking” electrons in the course of bond
formation.
The Lewis symbol for an element is the
chemical symbol for the element, and a
dot for each valence electron
More Lewis symbols!
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Dots are placed on four sides of the
symbol, with each side holding up to two
dots for a total of eight dots (Octet rule)
(Noble gases)
The placement of the dots is arbitrary
Ex.
..
.
Be
B:
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Can you figure out the relationship
between dots and groups?
Octet Rule
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Atoms tend to lose or gain electrons to
achieve the same number of electrons
as the noble gas closest to them in the
periodic table.
Octet Rule- atoms tend to lose or gain
electrons until they are surrounded by
eight valence electrons.
Of course, there are exceptions
Ionic Bonding

Na(s) + ½ Cl2(g)  NaCl + 410.9 kJ

The production of sodium chloride
involves the formation of an ionic bond
between sodium and chlorine.
Highly exothermic Hf = -410.9 kj/mol
Why ½?
Why -?
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Draw it
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In the formation of this ionic bond, the
metal with a low ionization energy
transfers an electron to the nonmetal
with the high electron affinity!
..
..
Na· + · Cl :  Na+1 + [ : Cl : ]-1
..
..
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Each ion has a full octet!
Enthalpy (heat of formation) is highly
exothermic!
Exothermic for the formation of most
ionic substances.
Why?
Practice, Practice, Practice
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Try to draw the dot diagrams for the
formation of the following ionic
compounds:
LiF
MgCl2
K2S
BeO
Back to NaCl & Energies of
Bond Formation!
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Formation of NaCl = highly exothermic
Ionization of Na requires 496 kJ
Adding an electron to Cl releases 349
kJ/mol
496-349 = 147 kJ/mol This reaction
seemingly requires energy. Can you
account for the difference?
Lattice Energy
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Ionic compounds stay together because
of strong attractions between the ions.
The attraction draws the ions together,
releasing energy, and causing the ions to
form a solid array or “lattice”
Lattice Energy- the energy required to
completely separate a mole of solid ionic
compound into it’s ions
Lattice Energy
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NaCl(s)  Na+ + ClThe lattice energy is given as
ΔHlattice = +788 kJ/mol
Which means in order to break the lattice
you would need to add 788 kJ’s of
energy

In order to form NaCl (the opposite
process) you would be releasing the
same amount of energy

ΔHf = -788 kJ/mol (highly exothermic)
Lattice Energy
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E= k(Q1Q2)/d
Q is the charges of the particles.
d is the distance between the centers.
k=8.99 x 109 J-m/C2
Lattice Energy
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Basically, the lattice energy depends on
the size of the ions, and their charge.
Lattice energy increases as the charge
on the ions increase, and their radii
decrease
Electron Configuration of
representative elements
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Group 1= +1
Group 2= +2
Group 3= +3
Group 5= -3
Group 6= -2
Group 7= -1
Group 8= 0
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You would never find a Na+2 ion. Even
though lattice energy increases with
increasing charge, the energy needed to
break past the valence shell is much to
high for that to occur
What happened to group 4?
We rarely find ionic compounds from
group 4, usually covalent.
Transition metal ions
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The lattice energies of ionic compounds
is usually large enough to compensate
for the loss of up to 3 electrons
Most transition metals have more than 3
electrons beyond a noble gas
configuration so achieving a noble gas
configuration is not usually possible
Transition metal ions
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Metals of group 1B usually form +1 ion (Cu,
Ag, Au)
Generally transition metals do not form ions
with noble gas configurations, and do not
satisfy the octet rule
In forming ions, transition metals lose the
valence shell s electrons first, then as many d
electrons as are required to reach the charge
of the ion
Transition metal ions
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Valence shell s are those that have the
largest value of n
According to our drawing filling order is
as follows: 1s 2s 2p 3s 3p 4s 3d 4p 5s
4d 5p 6s 4f etc.
Ag = [Kr] 5s1 4d10 (why written like this?)
Loses the 5s electron first to form Ag+1
Sizes of ions
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Size of ions is crucial in determining the
stability of an ionic solid.
Size of an ion depends on its nuclear
charge, the number of electrons it
possesses, and the orbitals where the
outer electrons are
Sizes of ions
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Cations are smaller than their parent atomspositive ions are formed by removing an
electron from the most spatially extended
orbitals, decreasing the total electron-electron
repulsions
Anions are larger than their parent atomswhen an electron is added to form an anion,
the increased electron- electron repulsion
causes the electrons to spread out
Ions of the same charge
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Size increases as you go down a group
for both the parent atom, and the ion
across a row they get smaller, and then
suddenly larger.
First half are cations.
Second half are anions.
Size of Isoelectronic ions
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Iso - same
Iso electronic ions have the same # of
electrons
Al+3 Mg+2 Na+1 Ne F-1 O-2 and N-3
All have 10 electrons.
All have the configuration 1s22s22p6
Size of Isoelectronic ions
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Positive ions have more protons so they
are smaller.
Al+3
Na+1
Mg+2
Ne
F-1
O-2
N-3
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Notice that Nitrogen is the largest ion,
and has the smallest atomic number
Aluminum is the smallest ion, and has
the largest atomic number.
Characteristics of ionic solids
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Brittle
High melting points
Crystalline
Cleaved (break apart along smooth, flat
surfaces)
In a rigid three dimensional arrangement
Gas, water, plastic
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These do not exhibit ionic properties
Lewis proposed that these atoms might
achieve a noble gas configuration by
sharing electrons
H2 is the simplest explanation of
covalent bonding
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When two hydrogen atoms are close
together, electrostatic interactions occur
between them. The attractions between
the electrons and the nucleus cause the
electron density to concentrate between
the nuclei
Essentially, the shared pair of electrons
act as a glue to bind atoms together
Lewis Dot Diagrams
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Writing Lewis structures, show each
electron pair between atoms as a line,
and unshared electron pairs as dots.
For the nonmetals, the number of
valence electrons is the same as the
group number
(Reminder covalent bonds occur
between two non-metals)
Multiple Bonds
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In many molecules atoms fill their octets
by sharing more than one pair of
electrons between them
When two electron pairs are shared =
double bond
When three electron pairs are shared=
triple bond
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The average distance between bonded
atoms differs with the number of shared
electron pairs. The more shared pairs,
the shorter the bond.
Common molecules that contain multiple
bonds: CO2, N2, O2,
Covalent Bonding
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Electrons are shared by atoms.
There are two ways of sharing: equally &
unequally
In polar covalent bonds, the electrons
are not shared evenly, one end is slightly
positive, the other slightly negative.
These partial charges are indicated
using small delta d.
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These partial charges come when one
atom exerts a greater attraction for
electrons than the other.
If the difference in ability to attract
electrons is large enough the one atom
will completely “steal” the electron and
now we are back to ionic bonding
Non-Polar Covalent Bonds
Equal sharing
Most common= are the diatomics
For example, each Hydrogen atom in H2
has the same desire for electrons, so the
atoms are shared equally.
Electronegativity
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We can use electronegativity to predict
whether a bond will be ionic, polar covalent, or
non-polar covalent
Electronegativity- the ability of an atom in a
molecule to attract electrons to itself.
Related to electron affinity, and ionization
energy
Atoms with very negative electron affinities,
and high ionization energies will be highly
electronegative
Linus Pauling
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(1901-1994)
Developed an electronegativity scale
with no units
Fluorine has the greatest
electronegativity (4.0)
Cesium has the lowest (0.7)
Electronegativity Trends
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(For the Representative Elements)
Across a period, electronegativity
increases
Down a group elecronegativity
decreases
Predicting Bonds
Compound
F2
Electronegativity
difference
Bond Type
HF
LiF
4.0
4.0
4.0
- 4.0
- 2.1
-1.0
0
1.9
3.0
Non- polar Polar
Ionic
Covalent Covalent
d+ d-
H-F
Zero
Intermediate
Large
Non polar
Polar
Ionic
Covalent Character
decreases
Ionic Character increases
Electronegativity Bond
difference
Type
Dipole Moments
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A molecule with a center of negative
charge and a center of positive charge is
dipolar (two poles),
or has a dipole moment.
Center of charge doesn’t have to be on
an atom.
Will line up in the presence of an electric
field.
How It is drawn
H-F
Drawing Lewis Structures
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Single lines represent 2 electrons
Everything must satisfy the octet rule
except for Hydrogen
Try : Phosphorous trichloride
CH2Cl2
HCN
The rules may help!
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You can start by adding the valence electrons
from all the atoms
Write the symbols for the atoms, and connect
them with a single bond
Complete the octets of the atoms bonded to
the central atom first
Place leftover electrons on the central atom
Try multiple bonds
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BrO3-1
PO4-3
CCl4
Formal Charge
 The
difference between the number of
valence electrons on the free atom and that
assigned in the molecule.
 We count half the electrons in each bond as
“belonging” to the atom.
 Molecules try to achieve as low a formal
charge as possible.
Examples
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[:C:::N:]First, count the # of electrons on C.
Total of 5. The formal charge on C is 45=-1
For N the formal charge is 5-5=0
Thus the total formal charges add up to 1 which is the charge of the ion
But how does it help?
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What about CO2
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O::C::O
6 4 6
-6 4 6
0 0 0
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or
O:C:::O
6 4 6
-7 4 5
-1 0 +1
Go for most stable
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When there are several possible Lewis
Structures the most stable one will be
the one where the atoms bear the
smallest formal charges, and any
negative charge resides on the more
electronegative atom.
The first model would be the best
because the atoms carry no formal
charges
Practice
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There are three possible Lewis
Structures for the thiocyanate ion NCS-1
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[N:C:::S]-
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Which one is the best?
[N::C::S]-
[N:::C:S]-1
Resonance
Sometimes there is more than one valid
structure for an molecule or ion.
 NO3
 Use double arrows to indicate it is the “average”
of the structures.
 It doesn’t switch between them.
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Drawing Resonance Structures
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The same atoms must be bonded to one
another in all structures
The only difference is the arrangement
of the electrons
Resonance in Benzene
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Important aromatic organic molecule
Millions of compounds contain the six
carbon ring
C6H6
Exceptions to the octet rule
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(Remember Lewis Diagrams didn’t work for all
transition elements? There are exceptions for
covalent models too!)
Three main exceptions
1) Molecules with an odd # of electrons
2) Molecules in which an atom has less than
an octet
3) Molecules in which an atom has more than
an octet.
Odd Number of electrons
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Rare
Example= NO, NO2, ClO2
NO N=5 O=6 5+6=11
Impossible to have complete pairing, so
an octet for each atom won’t occur
Less than an Octet
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Also rare, usually in molecular
compounds of Boron and Beryllium.
The resonance structures for Boron
doesn’t work. If you were to calculate
the formal charges the B would have the
negative and the F would have the
positive. Since this is not consistent with
Fluorine’s high electronegativity
More than an octet
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Biggest class of exceptions
There are more than eight electrons in
the valence shell of an atom
PCl5 , AsF6, ICl4
The corresponding models with a
second period atom do not exist (NCl5)
Only can have an “expanded” valence
shell if you are in period 3 and beyond
More than an octet
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Although third period elements (and
beyond) do usually satisfy the octet rule,
they often exceed the octet rule by using
their empty d orbitals to accommodate
extra electrons.
Expanded valence shells usually occur
when the central atom is bonded to the
smallest and most electronegative atoms
(Cl, F, & O)
Strength of covalent bonds
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The strength of the bond is determined
by the amount of energy required to
break that bond
D(bond type) represents bond
enthalpies
Bond enthalpy is always positive, you
always need to put in energy in order to
break a bond
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The greater the bond enthalpy the
stronger the bond
Molecules with strong chemical bonds
are not likely to undergo chemical
change.
Enthalpies of Reactions
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Endothermic (ΔH>0)
Exothermic (ΔH<0)
Reactions occur in 2 steps
1) Break bonds of reactants
2) Make bonds of products
ΔH(rxn)= (bond enthalpies of bonds
broken-bond enthalpies of bonds
formed)
Covalent Bond Energies
 We
made some simplifications in
describing the bond energy of CH4
 Each C-H bond has a different energy.
 CH4  CH3 + H
DH = 435 kJ/mol
 CH3  CH2 + H
DH = 453 kJ/mol
 CH2  CH + H
DH = 425 kJ/mol
 CH C + H
DH = 339 kJ/mol
 Each bond is sensitive to its environment.
Averages
 Have
made a table of the averages of
different types of bonds pg. 365
 single bond one pair of electrons is shared.
 double bond two pair of electrons are
shared.
 triple bond three pair of electrons are
shared.
 More bonds, shorter bond length.
Using Bond Energies
can find DH for a reaction.
 It takes energy to break bonds, and end up
with atoms (+).
 We get energy when we use atoms to form
bonds (-).
 If we add up the energy it took to break the
bonds, and subtract the energy we get from
forming the bonds we get the DH.
 Energy and Enthalpy are state functions.
 We
Find the energy for this
2 CH2 = CHCH3 + 2NH3 + O2
 2 CH2 = CHC  N + 6 H2O
C-H
C=C
N-H
C-C
413 kJ/mol
614kJ/mol
391 kJ/mol
347 kJ/mol
O-H 467 kJ/mol
O=O 495 kJ/mol
CN 891 kJ/mol
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As the number of bonds between two
atoms increase the strength of the bond
increases and the length of the bond
decreases