CONCEPTS OF
CHEMICAL BONDING
Chapter 8
Chem 2
8.1 Chemical Bonds
• Three basic types of bonds:
• Ionic
• •Electrostatic attraction
•
•
•
•
•
between ions
Covalent
•Sharing of electrons
Metallic
•Metal atoms bonded to
several other atoms.
•Electrons are free to move
around the structure.
Lewis Symbols
• Electrons involved in chemical bonding
are the valence electrons.
• G.N. Lewis (1875-1946) suggested a
simple way of showing the valence
electrons in an atom.
• Lewis electron-dot structures: consist of
the chemical symbol for the element plus a
dot for each valence electron.
Octet Rule
• Atoms tend to gain, lose, or share
electrons until they are surrounded by
eight valence electrons.
• An octet of electrons consists of full s
and p subshells in an atom.
• There are many exceptions to the
octet rule, but it provides a useful
framework for many important
concepts of bonding.
Ionic Bonding
• Na(s) + ½ Cl2(g) → NaCl(s)
∆Hºf = -409 kJ/mol
• Extremely exothermic
• NaCl is more stable than its constituent elements
• NaCl forms a regular structure in which each Na+ is
surrounded by six Cl- ions
• Closely packed ions
• Ionic lattice
Lewis Dot for Ionic Bonding
• SrCl2
Sr
+
Cl
Cl
[ Cl ]-1 [Sr]+2 [ Cl ]-1
Energetics of Ionic Bonding
• Removing an electron is always
endothermic.
• It takes +495 kJ/mol to remove electrons from
sodium (Na+).
• When a nonmetal gains an electron, it is
generally exothermic
• Adding an electron to chlorine releases
349 kJ/mol.
• 496 – 349 = +147 kJ/mol
-
Energetics of Ionic Bonding
• But these
numbers don’t
explain why the
reaction of
sodium metal
and chlorine gas
to form sodium
chloride is so
exothermic!
Energetics of Ionic Bonding
• What is as yet unaccounted for is the
electrostatic attraction between the newly
formed sodium cation and chloride anion.
• 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.
Lattice Energy
• The lattice energy is:
• NaCl(s) →Na+(g) + Cl-(g) ∆Hlattice= +788 kJ/mol
• The opposite process, the coming together of
Na+ and Cl- to form NaCl, is exothermic
(∆Hlattice = -788 kJ/mol)
Lattice Energy
• The energy associated with electrostatic
interactions is governed by Coulomb’s law:
Eel= ᴷ Q1Q2
d
• Q1 and Q2 are the charges, d is distance, and k is
constant, 8.99 x 109 J-m/C2
Lattice Energy
• Lattice energy, then, increases with the
charge on the ions.
• It also increases
with decreasing
size of ions
Energetics of Ionic Bonding
• These phenomena
also help 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
Example
• Arrange the following ionic compounds
in order of increasing lattice energy.
• NaF, CsI, and CaO
Solution
• CsI < NaF < CaO
• CaO has greatest ionic charge at +2 and -2
• Cs+ is larger than Na+
• I-is larger than F• The distance between the Na+ and F-ions in
NaF will be less than the distance between Cs+
and I- ions in CsI. As a result, the lattice energy
of NaF should be greater than that of CsI.
Example
• •Which substance would you expect to
have the greatest lattice energy?
• AgCl
• CuO
• CrN
Solution
• CrN
8.3 Covalent Bonding
• 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
• The attractions between nuclei and the
electrons cause electron density to
concentrate between the nuclei. As a
result, the overall electrostatic
interactions are attractive.
• A shared pair of electrons in any
covalent bond acts as a king of “glue”
to bind atoms together.
Lewis Structures
• •The formation of covalent bonds can
be represented using Lewis symbols
• Examples: Draw: H2, Cl2, NH3, CH4
• •Multiple Bonds: Doubles and Triples
• Examples: Draw CO2 and N2
• •As a rule, the distance between
bonded atoms decreases as the
number of shared electron pairs
increases.
8.4 Polar Covalent Bonds
• Although atoms often form compounds by
sharing electrons, the electrons are not
always shared equally.
• Flourine pulls harder on the electrons it
shares with hydrogen than hydrogen does.
• Therefore, the flourine end of the molecule
has more electron
density than the
hydrogen end.
Polarity
• Nonpolar covalent bond is one in which the
electrons are shared equally between two
atoms, as in the Cl2 and N2 bonds.
• Also compounds that are “balanced” like CCl4
• Polar covalent bond, one of the atoms
exerts a greater attraction for the bonding
electrons than the other. If the difference in
relative ability to attract electrons is large
enough, an ionic bond is formed.
Electronegativity
• The ability of atoms in
a molecule to attract
electrons to itself.
• 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.
Vocabulary
• Electronegativity – the ability of an
atom IN A MOLECULE to attract
electrons to itself.
• Ionization energy – how strongly an
atom holds on to its electrons.
• Electron affinity – the measure of how
strongly an atom attracts additional
electrons.
Polar Covalent Bonds and
Electronegativity
• The greater the
difference in
electronegativity
the more polar is
the bond.
• F2
• 4.0 – 4.0 = 0
• Nonpolar
• HF
• 4.0 – 2.1 = 1.9
• Polar Covalent
• LiF
• 4.0 – 1.0 = 3.0
• Ionic
Example
• Which bond is more polar? Indicate in
each case which atom has the partial
negative charge.
• B-Cl or C-Cl
Solution
• Use Figure 8.6 (pg 285)
• The difference in the electronegativities of
chlorine and boron is 3.0 –2.0 = 1.0
• The difference between chlorine and carbon is
3.0 –2.5 = 0.5
• Therefore B-Cl is more polar. The chlorine atom
carries the partial negative charge because it
has a higher electronegativity.
Polar Covalent Bonds
• When two atoms share electrons unequally, a
bond dipole results.
• The dipole moment, μ, produced by two equal but
opposite charges separated by a distance, r, is
calculated:
μ = Qr
• •It is measured in debyes (D).
• •A unit that equals 3.34 x 10-34 Coulomb-meters
What this means…
μ = Qr
• For molecules, we usually measure charge in
units of electronic charge, e,1.6 x 10-19C, and
distance in units of angstroms.
• Suppose we have a charge of 1+ and 1-and
are separated by a distance of 1 Angstrom. A
dipole moment produced is
=(1.60 x 10-19C)(1A)(10-10m/1A)( 1D/3.34x10-30C-m)
= 4.79 D
Example
• The bond length in the HCl molecule is
1.27 Angstroms.
• Calculate the dipole moment, in debyes,
that would result if the charges on the H
and Cl atoms were 1+ and 1-respectively.
Solution
μ = Qr
• = (1.60 x 10-19C)(1.27A) (10-10m/1A)
(1D/3.34 x 10-30C-m)
• = 6.08 D
8.5 Lewis Structures
• Lewis structures are representations of
molecules showing all electrons, bonding
and nonbonding.
Writing Lewis Structures
1. Find the sum of
valence electrons
of all atoms in the
polyatomic ion or
molecule.
PCl3
If it is an anion, add
one electron for each
negative charge.
2. If it is a cation, subtract
one electron for each
positive charge.
1.
5 + 3(7) = 26
Writing Lewis Structures
Keep track of the
electrons:
26 – 6 = 20
2.The central atom is
the least
electronegative
element that isn’t
hydrogen. Connect
the outer atoms to it
by single bonds.
Writing Lewis Structures
the
3. Fill the octets of
outer atoms.
Keep track of the electrons
26 – 6 = 20 – 18 = 2
3. Fill the octet of the
central atom.
Keep track of the electrons
26 – 6 = 20 – 18 = 2
Writing Lewis Structures
5. If you run out of
electrons before
the central atom
has an octet…
…form multiple
bonds until it
does.
Writing Lewis Structures
• Then assign formal charges.
• For each atom, count the electrons in lone pairs
and half the electrons it shares with other atoms.
• Subtract that from the number of valence
electrons for that atom: The difference is its
formal charge.
Writing Lewis Structures
• The best Lewis structure…
• …is the one with the fewest charges.
• …puts a negative charge on the most
electronegative atom.
Resonance
• This is the Lewis structure we would draw
for ozone, O3.
Resonance
• But this is at odds
with the true,
observed structure of
ozone, in which…
• …both O—O bonds
are the same length.
• …both outer oxygens
have a charge of -1/2.
Resonance
• One Lewis structure
cannot accurately
depict a molecule
such as ozone.
• We use multiple
structures, resonance
structures, to describe
the molecule.
Resonance
Just as green is the
synthesis of blue and
yellow…
…ozone is a synthesis
of these two
resonance structures
Resonance
• In truth, the electrons that form the second C—
O 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.
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.
8.7 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).
Odd Number of Electrons
Though relatively rare and usually quite
unstable and reactive, there are ions
and molecules with an odd number of
electrons.
Examples: ClO2, NO, and NO2
Complete pairing of electrons is
impossible.
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.
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.
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.
• Presumably d orbitals
in these atoms
participate in bonding.
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.
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.
8.8 Covalent Bond Strength
• Most simply, the strength of a bond is measured
by determining how much energy is required to
break the bond.
• Cl2(g) → 2Cl(g)
• This is the bond enthalpy. Always a positive
quantity.
• The bond enthalpy for a Cl—Cl bond,
D(Cl—Cl), is measured to be 242 kJ/mol.
Average Bond Enthalpies (pg 301)
• 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.
Average Bond Enthalpies
NOTE: These are
average bond
enthalpies, not
absolute bond
enthalpies; the C—
H bonds in
methane, CH4, will
be a bit different
than the
C—H bond in
chloroform, CHCl3.
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= Σ(bond enthalpies of bonds broken) Σ(bond enthalpies of bonds formed)
Enthalpies of Reaction
• CH4(g)+ Cl2(g) →
CH3Cl(g)+ HCl(g)
• In this example, one
C—H bond and one
Cl—Cl bond are
broken; one C—Cl
and one H—Cl bond
are formed.
Enthalpies of Reaction
• So,
∆Hrxn= [D(C—H) + D(Cl—Cl) - [D(C—Cl) +
D(H—Cl)
= [4(413 kJ) + 1(242 kJ)] - [4(413 kJ) +1(328 kJ) +
1(431 kJ)]
= (1894 kJ) - (1998 kJ)
= -104 kJ
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.