Chapter 8 Concepts of Chemical Bonding

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Chapter 8Concepts of Chemical Bonding
Chemistry, The Central Science, 10th
edition Theodore L. Brown H. Eugene LeMay, Jr.
and Bruce E. Bursten
John D. Bookstaver St. Charles Community
College St. Peters, MO ? 2006, Prentice Hall, Inc.
2
Introduction

• Chapter 6 atoms and atomic structure
• Chapter 7 trends in the periodic table
• Ionization energy
• Electron affinity
• Chapter 8 Bonding how are atoms held together
  in molecules? (or how are ions held together in
  an ionic compound?)

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Chemical Bonds

• Three basic types of bonds
• Ionic (ionic compounds)
• Electrostatic
• attraction between oppositely charged ions
• repulsion bewteen similarly charged ions
• Covalent (molecular)
• Sharing of electrons
• Metallic
• Metal atoms bonded to several other atoms

e.g. NaCl-
Well focus on these two
Each line 2 shared electrons
e.g.
e.g.
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Ionic Bonding
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The Formation of NaCl(s) from Na(s) and Cl2(g)

• The formation reaction for NaCl(s) is written as
• Na(s) ½Cl2(g) ? NaCl(s)
• DHof,NaCl(s) -410.9 kJ/mol
• (from Appendix C)
• Step-by-step, how would Na be transformed into
  Na and Cl2 be transformed into Cl-?
• Two important things (among others) would need to
  happen
• Na ? Na and
• Cl ? Cl-

Exothermic
Na(s) ½ Cl2(g)
Na(g) Cl-(g)
NaCl(s)
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Energetics of Ionic Bonding

• As we saw in the last chapter, it takes 495
  kJ/mol to remove electrons from sodium.

Na(g) ? Na(g) e- I1 495 kJ/mol
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Energetics of Ionic Bonding

• We get 349 kJ/mol back by giving electrons to
  chlorine.

Cl(g) e- ? Cl-(g) E1 -349kJ/mol
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Considering These Two Energies

• Na(g) ? Na(g) e- DH 495kJ/mol
• Cl(g) e- ? Cl-(g) DH -349kJ/mol
• Na(g) Cl(g) ? Na(g) Cl-(g) DH 146kJ/mol

Endothermic
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Energetics of Ionic Bonding

• But these numbers dont explain why the reaction
  of sodium metal and chlorine gas to form sodium
  chloride is so exothermic!

Na(s) ½ Cl2(g)
endo?
Na(g) Cl-(g)
Na(s) Cl2(g) ? NaCl(s)
NaCl(s)
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Energetics of Ionic Bonding
net charge 1

• There must be another process (at least) which
  influences DH for this process
• What we havent yet unaccounted for is the
  electrostatic attraction between the newly formed
  sodium cation and chloride anion (in a
  lattice-like structure)

net charge -1
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NaCl Lattice
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Lattice Energy

• Oppositely charged ions are attracted to each
  other. When two oppositely charged ions are
  brought together, energy decreases
• The energy of attraction that holds together
  oppositely charged ions in an ionic lattice is
  called the lattice energy, Eel
• Lattice energy, Eel The energy required to
  completely separate a mole of a solid ionic
  compound into its gaseous ions (thus defined here
  as a positive energy)
• The energy associated with electrostatic
  interactions is governed by Coulombs law

d is the inter-nuclear separation
Q1 and Q2 are the charges (C) of the cation and
anion (absolute values)
k 8.99 x 109 J.m/C2
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Lattice Energy

• Lattice energy, then, increases with the charge
  on the ions.

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Lattice Energy

• Lattice energy, then, increases with the charge
  on the ions.

• Compare
• LiF-
• Mg2O2-
• Sc3N3-

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Lattice Energy

• It also increases with decreasing size of ions.

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Lattice Energy

• It also increases with decreasing size of ions

• Compare
• LiF
• LiCl
• LiI

F-
Cl-
I-
Remember top-to-bottom in a group, size increases
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Covalent Bonding
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Covalent Bonding (e.g. H2)
H H ? H2

• In these bonds atoms share electrons.
• There are several electrostatic interactions in
  these bonds
• Attractions between electrons and nuclei
• Repulsions between electrons
• Repulsions between nuclei

Stable arrangement of electrons, nuclei
H-H
H

H
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Polar Covalent Bonds
Consider the following diatomic (2-atom) molecules
H-F
F-F

• Although atoms often form compounds by sharing
  electrons, the electrons are not always shared
  equally.

homonuclear
heteronuclear
Polar covalent bond or polar bond
Non-polar covalent bond

• Fluorine pulls at the bonding electrons more than
  hydrogen than hydrogen does in HF.
• Therefore, the fluorine end of the molecule has
  more electron density than the hydrogen end.

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Electronegativity, EN

• The ability of atoms in a molecule to attract
  bonding electrons toward itself and away from the
  other atom involved in the bond.
• On the periodic chart, electronegativity
  increases as you go
• from left to right across a row.
• from the bottom to the top of a column.

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Polar Covalent Bonds

• The greater the difference in electronegativity,
  the more polar the bond.

a positive side
a negative side
H-F
H-Cl
H-Br
H-I
DEN 0 to 0.4 non-polar covalent bond DEN 0.4
lt x lt 1.9 polar covalent bond DEN x gt 1.9
ionic bond
For LiF DEN 3.0 (ionic) Lithium fluoride is
written as LiF-
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Polar Bonds

• When bonding electrons are unequally shared, a
  bond dipole results
• One atom of the bond possesses more negative
  charge than it would given equal electron sharing
  and the other, less negative charge
• For example, in the following hypothetical
  molecule (H-X, X more electronegative than H),
  there are 0.5 e- more charge on X in this
  structure than there would be if the electrons of
  the H-X bond were equally shared

This atom has an effective charge of -0.5 e
This also means H has 0.5 e- less than it
would in a H-X molecule where the electrons of
the bond are equally shared
H
X
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Polar Covalent Bonds

• Just how unequally electrons in the bond are
  shared is reported as the dipole moment, m.
• The dipole moment produced by two equal but
  opposite charges separated by a distance, r, is
  calculated
• ? Qr
• It is reported in debyes (D).

1D 3.34x10-30 C.m
Charge on 1e- 1.602 x 10-19 C
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Dipole Moment

• For HCl, which has a dipole moment of 1.08 D and
  a bond length of 1.27 Å, calculate the effective
  charges on H and Cl in units of electronic
  charge, e.
• ? Qr
• We are given m (1.08 D) and r (1.27 Å)
• We want to solve for Q, the effective charge

C coulomb, SI unit of charge
1D 3.34x10-30 C.m
1Å 1 x 10-10 m
Charge on one electron 1.602 x 10-19 C
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Lewis Structures
hydrogen, H2
chlorine, Cl2
non-bonding electrons (valence) also called lone
pairs
covalent bond

• Lewis structures are representations of
  molecules showing all electrons, bonding and
  nonbonding.

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Writing Lewis Structures

• Find the sum of valence electrons of all atoms in
  the polyatomic ion or molecule.
• If it is an anion, add one electron for each
  negative charge.
• If it is a cation, subtract one electron for each
  positive charge.

• PCl3

5 3(7) 26
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Writing Lewis Structures

• The central atom is the least electronegative
  element that isnt hydrogen. Connect the outer
  atoms to it by single bonds.

Each bond costs 2 e-
Keep track of the electrons 26 ? 6 20
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Writing Lewis Structures

• Surround each of the outer atoms with enough
  electrons to give each outer atom 8 surrounding
  electrons (this doesnt apply for hydrogen H
  can only hold two electrons)

Keep track of the electrons 26 ? 6 20 ? 18 2
Non-metals tend to be surrounded by 8 electrons
in Lewis diagrams octet rule (2nd row
elements)
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Writing Lewis Structures

• After completing the octets of the outer atoms,
  use any remaining electrons to complete the octet
  of the central atom

Keep track of the electrons 26 ? 6 20 ? 18
2 ? 2 0
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Writing Lewis Structures

• If you run out of electrons before the central
  atom has an octet
• form multiple bonds until it does.

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Try these out

• Give the Lewis structures for the following
• CH4
• NH4
• H2O
• CO2
• HCO2-
• C2H2

• Rules
• Count valence electrons
• Put least EN atom in center,
• connect with bonds
• Fill octets of outer atoms
• If any electrons left over,
• put them on the central atom
• 5) Use multiple bonds if necessary

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Formal Charges

• We could arrive at any of the following Lewis
  structures by following the rules discussed
  earlier
• These structures differ in the location of the
  electron pairs (as lone pairs or in bonds)

1 2 3
Double-headed arrows indicate that
these structures only differ in the location of
their electron pairs
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Writing Lewis Structures

• Assigning formal charges
• For each atom, count the electrons in lone pairs
  and half the electrons it shares with other atoms
  (in bonds)
• Subtract that from the number of valence
  electrons for that atom The difference is its
  formal charge.

same as group (e.g. C is group 4A)
Formal charge on an atom group - of
lone of bonds
electrons
number
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Writing Lewis Structures

• The best Lewis structure
• is the one with the fewest charges.
• puts a negative charge on the most
  electronegative atom.

Draw all Lewis structures for cyanate ion
(CNO-), and determine which is most correct
-2 0 1 -1 0 0
0 0 -1
Notice that the formal charges always sum to
equal the charge of the ion
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Resonance

• This is the Lewis structure we would draw for
  ozone, O3.

-
Average Bond Lengths
double bonds are shorter than single bonds
1.48Å 1.21Å
1Å 1 x 10-10 m
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Resonance

• But this is at odds with the true, observed
  structure of ozone, in which both OO bonds are
  the same length.

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Resonance

• One Lewis structure cannot accurately depict a
  molecule such as ozone.
• We use multiple structures, resonance structures,
  to describe the molecule.

1
2
3
1
2
3
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Resonance

• Just as green is a synthesis of blue and yellow
• the true Lewis structure for ozone would be a
  combination of the two shown on the left

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Resonance

• The electrons that form the second CO bond in
  the double bonds below do not always sit between
  that C and that O, but can exist in the other
  bond instead.
• They are not localized, but rather are
  delocalized.

formate ion HCO2-
means resonance structures
delocalized spread out
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Resonance

• The organic compound benzene, C6H6, has two
  resonance structures.
• It is commonly depicted as a hexagon with a
  circle inside to signify the delocalized
  electrons in the ring.

combined format
Many organic structures include benzene rings,
and rather than drawing all resonance forms, they
are often drawn with the combined structure
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Exceptions to the Octet Rule

• There are three cases (ions or molecules) where
  the octet rule is not followed
• Ions or molecules with an odd number of
  electrons.
• Ions or molecules with less than an octet.
• Ions or molecules with more than eight electrons
  on central atom (an expanded octet).

central atom is group 2 or group 3
For central atoms from 3rd row (or later)
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Odd Number of Electrons

• Though relatively rare and usually quite
  unstable and reactive, there are ions and
  molecules with an odd number of electrons.

e.g. NO (11 e-)
e.g. NO2 (17 e-)
These will often involve a group 5A element (e.g.
N)
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Fewer Than Eight Electrons
This case usually applies for formulas involving
group 2A and 3A elements

• Consider BF3
• Giving boron a filled octet places a negative
  charge on the boron and a positive charge on
  fluorine.
• This would not be an accurate picture of the
  distribution of electrons in BF3.

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Fewer Than Eight Electrons

• Therefore, structures that put a double bond
  between boron and fluorine are much less
  important than the one that leaves boron with
  only 6 valence electrons.

Experimental bond length found to be slightly
less than average B-F single bond length
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More Than Eight Electrons
This rule applies for heavier central atoms (3rd
period and lower)

• The only way PCl5 can exist is if phosphorus has
  10 electrons around it.
• For central atoms from the third row or later,
  expanded octets are possible.
• Presumably d orbitals in these atoms participate
  in bonding.

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More Than Eight Electrons

• Even though we can draw a Lewis structure for the
  phosphate ion that has only 8 electrons around
  the central phosphorus, the better structure (the
  one that agrees better with experimental
  evidence) puts a double bond between the
  phosphorus and one of the oxygens.

For phosphate, PO43-
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More Than Eight Electrons

• This eliminates the formal charge on the
  phosphorus and the formal charge on one of the
  oxygens.
• When the central atom is from the 3rd row (or
  below) of the periodic table, and expanding its
  octet eliminates some formal charges, do so.

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Covalent Bond Strength

• The strength of a bond is measured by determining
  how much energy is required to break the bond.
• This is the bond enthalpy.
• The bond enthalpy for a ClCl bond,
• D(ClCl), is measured to be 242 kJ/mol.

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Average Bond Enthalpies

• This table (8.4 in your text) lists the average
  bond enthalpies for many different types of
  bonds.
• Average bond enthalpies are positive, because
  bond breaking is an endothermic process.

e.g. CH4(g) ? C(g) 4H(g) DH 1660 kJ
and thus bond forming is exothermic
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DH for a reaction

• We have seen a few ways to calculate DH for some
  chemical reaction
• DHreaction SDHof,products SDHof,reactants
• DH DG TDS
• DH q/mol (from calorimetry)

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Enthalpies of Reaction
CH4(g) Cl2(g) ? CH3Cl(g) HCl(g)

• Yet another way to estimate ?H for a reaction is
  to compare the bond enthalpies of bonds broken to
  the bond enthalpies of the new bonds formed.

To calculate DHrxn from bond enthalpies

• ?Hrxn ?(bond enthalpies bonds broken) ??(bond
  enthalpies bonds formed)

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Enthalpies of Reaction

• CH4(g) Cl2(g) ???
• CH3Cl(g) HCl(g)
• In this example, one
• CH bond and one
• ClCl bond are broken one CCl and one HCl bond
  are formed.

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Enthalpies of Reaction
bonds formed
bonds broken

• So,
• ?Hrxn D(CH) D(ClCl) ? D(CCl) D(HCl)
• (413 kJ/mol) (242 kJ/mol) ? (328 kJ/mol)
  (431 kJ/mol)
• (655 kJ/mol) ? (759 kJ/mol)
• ?104 kJ/mol

Compare with DH -101 kJ/mol for formation
enthalpies