Chemical Bonding

1
Chemical Bonding

• Chapter 8

2
Chemical Bonds

• Three basic types of bonds
• Ionic
• Electrostatic attraction between ions
• Covalent
• Sharing of electrons
• Metallic
• Metal atoms bonded to several other atoms

3
Ionic Bonding
4
Energetics of Ionic Bonding

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

5
Energetics of Ionic Bonding

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

6
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!

7
Energetics of Ionic Bonding

• There must be a third piece to the puzzle.
• What is as yet unaccounted for is the
  electrostatic attraction between the newly formed
  sodium cation and chloride anion.

8
Lattice Energy

• This third piece of the puzzle is the lattice
  energy
• The energy required to completely separate a mole
  of a solid ionic compound into its gaseous ions.
• The energy associated with electrostatic
  interactions is governed by Coulombs law

9
Lattice Energy

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

• It also increases with decreasing size of ions.

10
Energetics of Ionic Bonding

• By accounting for all three energies (ionization
  energy, electron affinity, and lattice energy),
  we can get a good idea of the energetics involved
  in such a process.

11
Energetics of Ionic Bonding

• These phenomena also helps explain the octet
  rule.

• Metals, for instance, tend to stop losing
  electrons once they attain a noble gas
  configuration because energy would be expended
  that cannot be overcome by lattice energies.

12
The Ionic Bond
He
Ne
1s22s1
1s22s22p5
1s2
1s22s22p6
9.2
13
Valence Electrons

• Valence electrons are the outer shell s and p
  electrons of an atom.
• The valence electrons are the electrons that
  participate in chemical bonding.
• We show the valence electrons in Lewis dot
  structures of atoms and molecules.

14
9.1
15
A covalent bond is a chemical bond in which two
or more electrons are shared by two atoms.
Lewis structure of F2
9.4
16
Lewis structure of water


Double bond two atoms share two pairs of
electrons
or
Triple bond two atoms share three pairs of
electrons
or
9.4
17
Lengths of Covalent Bonds
Bond Type Bond Length (pm)
C-C 154
C?C 133
C?C 120
C-N 143
C?N 138
C?N 116
Bond Lengths Triple lt Double lt Single
9.4
18
Comparison of Ionic and Covalent Compounds
9.4
19
Polar Covalent Bonds
Polar covalent bond or polar bond is a covalent
bond with uneven electron density
electron rich region
electron poor region
e- rich
e- poor
d
d-
9.5
20
Electronegativity
Electronegativity is the ability of an atom to
attract toward itself the electrons in a chemical
bond.
Electronegativity - relative, F is highest
9.5
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Electronegativity
9.5
22
Classification of bonds
Difference
Bond Type
0
Covalent
? 2
Ionic
0 lt and lt2
Polar Covalent
9.5
23
Polar Covalent Bonds

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

24
Solve The Following
Cs 0.7
Cl 3.0
3.0 0.7 2.3
Ionic
Polar Covalent
H 2.1
S 2.5
2.5 2.1 0.4
N 3.0
N 3.0
3.0 3.0 0
Covalent
9.5
25
Drawing Lewis Structures

• 1. Set aside the single bonders (H, F, Cl, Br,
  I).
• 2. Put together the 2, 3, and 4 bonders with
  single bonds (May require thinking).
• 3. Count the holes left
• A. If holes single bonders, put them on.
• B. If holes gt single bonders, make double or
  triple bonds, or a ring until A is satisfied.

26
Lewis Structures

• Examples
• draw a Lewis structure for hydrogen peroxide,
  H2O2
• draw a Lewis structure for methanol, CH3OH
• draw a Lewis structure for acetic acid, CH3COOH

27
Lewis Structures

• Draw a Lewis Structure and Make a model for
• C2H6O
• C2H4O
• BF3
• HNO3
• HCN
• C2H3Cl

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9.6
29
9.6
30
An atoms formal charge is the difference between
the number of valence electrons in an isolated
atom and the number of electrons assigned to that
atom in a Lewis structure.
The sum of the formal charges of the atoms in a
molecule or ion must equal the charge on the
molecule or ion.
9.7
31
Writing Lewis Structures

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

32
Formal Charge and Lewis Structures

1. Lewis structures with large formal charges are
   less plausible than those with small formal
   charges.
2. Among Lewis structures having similar
   distributions of formal charges, the most
   plausible structure is the one in which negative
   formal charges are placed on the more
   electronegative atoms.

9.7
33
Resonance Structures

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

-
34
Resonance Structures

• But this is at odds with the true, observed
  structure of ozone, in which
• both OO bonds are the same length.
• both outer oxygens have a charge of ?1/2.

35
Resonance Structures

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

36
Resonance Structures

• Just as green is a synthesis of blue and yellow
• ozone is a composite of these two resonance
  structures.

37
Resonance Structures

• In truth, the electrons that form the second CO
  bond in the double bonds below do not always sit
  between that C and that O, but rather can move
  among the two oxygens and the carbon.
• They are not localized, but rather are
  delocalized.

38
Resonance Structures

• 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.

39
Resonance Structures
A resonance structure is one of two or more Lewis
structures for a single molecule that cannot be
represented accurately by only one Lewis
structure.
O
O
40
Exceptions to the Octet Rule

• There are three types of ions or molecules that
  do not follow the octet rule
• Ions or molecules with an odd number of
  electrons.
• Ions or molecules with less than an octet.
• Ions or molecules with more than eight valence
  electrons (an expanded octet).

41
Odd Number of Electrons

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

42
Fewer Than Eight Electrons

• 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.

43
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.

44
Fewer Than Eight Electrons

• The lesson is If filling the octet of the
  central atom results in a negative charge on the
  central atom and a positive charge on the more
  electronegative outer atom, dont fill the octet
  of the central atom.

45
More Than Eight Electrons

• The only way PCl5 can exist is if phosphorus has
  10 electrons around it.
• It is allowed to expand the octet of atoms on the
  3rd row or below.
• The d orbitals in these atoms participate in
  bonding definitely!

46
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 puts
  a double bond between the phosphorus and one of
  the oxygens.

47
More Than Eight Electrons

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

48
Exceptions to the Octet Rule
The Incomplete Octet
BeH2
BF3
9.9
49
Exceptions to the Octet Rule
Odd-Electron Molecules
NO
The Expanded Octet (central atom with principal
quantum number n gt 2)
SF6
9.9
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Covalent Bond Strength

• Most simply, 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.

51
The enthalpy change required to break a
particular bond in one mole of gaseous molecules
is the bond energy.
Bond Energy
9.10
52
Average Bond Enthalpies

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

53
Average Bond Enthalpies

• NOTE These are average bond enthalpies, not
  absolute bond enthalpies the CH bonds in
  methane, CH4, will be a bit different than the
• CH bond in chloroform, CHCl3.

54
Enthalpies of Reaction

• 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.

• In other words,
• ?Hrxn ?(bond enthalpies of bonds broken) ?
• ?(bond enthalpies of 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.

56
Enthalpies of Reaction

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

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Bond Enthalpy and Bond Length

• We can also measure an average bond length for
  different bond types.
• As the number of bonds between two atoms
  increases, the bond length decreases.

58
DH0 SBE(reactants) SBE(products)
DH0 436.4 156.9 2 x 568.2 -543.1 kJ
9.10