Objectives

1
Objectives

• Use Lewis symbols to represent ionic bond
  formation.
• Review ionic radii trends and relate to lattice
  energy of ionic compounds.
• Use electronegativities to obtain relative bond
  polarity.
• Discuss covalent bond formation.
• Write Lewis structure given chemical formulas.
• Write necessary resonance structures.
• Use formal charges to determine the best Lewis
  structure.
• Relate bond order and bond length to lewis
  structures.
• Estimate DH from bond energies.

2
Chemical Bonding
3
We will now relate the concepts learned in
previous chapters to chemical bonding.
4
Chemical Bond
A strong attractive force that exists between
atoms in a molecule.
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Three Types of Bonds
Ionic Bond
Covalent Bond
Metallic Bond
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We will focus mainly on ionic and covalent bonds
in this chapter. But first, lets discuss Lewis
Symbols.
Lewis Symbol
A diagram in which the elemental symbol
represents the nucleus along with all inner
electrons, and the dots represent outer shell
electrons.
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Lewis symbols are very useful because they
portray the heart of a chemical bond valence
electrons!
Lets take a look at a few examples
Bohr Diagram
Lewis Diagram
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Dots represent valence electrons
Notice the negative sign on the chloride ion.
Dont forget to include the charge when writing
the Lewis symbols for ions.
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What would the Lewis symbols be for elemental
sodium and the sodium ion?
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Octet Rule
Atoms gain, share, or lose electrons with the
goal of achieving the electron configuration of
their neighboring noble gas. They do so because
as we know, the noble gases are stable. All
elements want to be like the NOBLE gases.
How many valence electrons do the noble gases
have?
8
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So, atoms gain, lose, or share electrons until
they are surrounded by 8 valence electrons.
Exceptions?
--Arent there always exceptions!-- These nasty
little buggers will show themselves later! For
now, stick to the rule!
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Its finally time to start talking about chemical
bonding. We will start with
Ionic Bond
Lets start by viewing a movie of a chemical
reaction that is the result of ionic bonding!
Click the icon below.
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Sodium chloride is composed of a crystal lattice
of sodium cations (Na) and chloride anions (Cl-).
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Take a closer look at the 3 dimensional array of
NaCl.
Display Model in Class
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Can you relate this reaction to the periodic
trends that you learned?
Sodium metal is on the far left of the periodic
table therefore it has a low ionization energy.
Remember, this is the energy required to remove
an electron. Since the required energy is so
low, sodium is willing to give up an electron
and thus, become positively charged.
Chlorine, on the other hand, is on the far right
of the periodic table, and therefore has a high
affinity for electrons. It is willing to take
the electron given up by the sodium atom, which
causes it to become negatively charged.
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Lets look at this from a Lewis symbol
perspective.
Notice how the electron is transferred from the
sodium atom to the chlorine atom resulting in
oppositely charged particles.
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The oppositely charged particles that result are
strongly attracted to one another just like the
opposite ends of a magnet.
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What is the chemical formula for Magnesium
Fluoride?
MgF2
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MgF2

• Magnesium has two electrons to give, whereas the
  fluorines have only one vacancy each.

• Consequently, magnesium can accommodate two
  fluoride ions.

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Ionic Bond
The electrostatic forces that exist between ions
of opposite charge as a result of the transfer of
one or more electrons from one atom to another.
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Do you think energy would be required to break
the ionic bonds that form the sodium chloride
crystal lattice, or do you think energy would be
released if the bonds of sodium chloride were
broken?
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Lattice Energy
Energy required to completely separate a mole of
a solid ionic compound into its gaseous ions.
NaCl (s) ? Na (g) Cl- (g) ?Hlattice 788
KJ/mol
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Energy Involved in Ionic Bonding

• The transfer of an electron from a sodium atom to
  a chlorine atom is not in itself energetically
  favorable it requires 147 kJ/mol of energy.
• 496 KJ/mol needed to remove an electron from Na.
• 349 KJ/mol released when Cl gains the electron.
• 496-349 147 KJ/mol for the transfer of the
  electron.

• However, 493 kJ of energy is released when these
  oppositely charged ions come together.
• An additional 293 kJ of energy is released when
  the ion pairs solidify.

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The magnitude of the lattice energy of a solid
depends on Charges of ions Size of
ions Type of arrangement
25
We wont be doing any calculations with this
equation, but it does allow us to compare lattice
energies of different ionic compounds.
Q1 and Q2 charges of particles d distance
between centers k 8.99 x 109 J-m/C2
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This equation tells us that as the charges
increase, and the radii decreases, the lattice
energy increases.
Compare the following in terms of increasing
lattice energy LiCl, LiF, MgCl2
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LiCl and LiF both have 1 and 1- charges.
Therefore, we must consider the sizes of the
ions. F- is smaller than Cl- So the distance
between Li and F- in the crystal structure will
be smaller than the distance between the Li and
Cl- ions. Therefore more energy is needed to
separate LiF LiCl lt LiF The fact that magnesium
has a 2 charge indicates that MgCl2 will have
the largest lattice of all three. LiCl lt LiF lt
MgCl2
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LiCl lt LiF lt MgCl2
Actual Lattice Energies MgCl2 2326 KJ/mol LiF
1030 KJ/mol LiCl 834 KJ/mol
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Review Sizes of Ions
Cations smaller than parent atom because the loss
of electron(s) increases the effective nuclear
charge and decreases electron-electron repulsion.
Anions larger than parent atom because the gain
of electron(s) increases electron-electron
repulsion causing electrons to spread out.
picture
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Anion
Parent
Cation
Parent
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Review Isoelectronic Series A group of ions that
have the same number of electrons. Example Increas
ing nuclear charge O2- F- Na Mg2 Al3 Decr
easing ionic radius
Electrons remain constant, while protons increase
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At room temperature, are these two compounds
solids, liquids, or gases?
Carbon monoxide CO
What can we conclude about the amount of heat
needed to boil CO and CO2?
Carbon dioxide CO2
Low Boiling Point
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Covalent Bonds
(Molecular Compounds)
A bond formed between two or more atoms as a
result of sharing of electrons.
Consists entirely of nonmetals.
Low Melting Points because the bonds between
molecules are not as strong as the bonds between
ions in crystal lattices.
Can be polar or nonpolar, which we will discuss
shortly.
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Covalent bonding explains why diatomic molecules
exist. H2, N2, O2, F2, Cl2, I2, Br2
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3 Types of Covalent Bonds
Single covalent bond one pair of electrons are
shared
Double covalent bond two pairs of electrons are
shared
Triple covalent bond three pairs of electrons
are shared
Examples
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Again, we start with Lewis Symbols!
Single , Double , or Triple?
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Orbital Diagram
H 1S _____
H2
H 1S _____
Which rule that you previously learned allows 1
electron from each atom to share the same orbital?
Pauli Exclusion Principle
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Pauli Exclusion Principle
The electrons that are shared spin in opposite
directions as explained in the Pauli Exclusion
Principle.
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Both hydrogen atoms now have a filled outer shell.
Both chlorine atoms now have a filled outer shell.
Notice that no ions are formed!
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Structural Formula Chemical formulas that show
the arrangement of atoms in molecules.
Each dash represents a shared pair of electrons.
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What is the structural formula for hydrogen?
A single dash shows that it is a single covalent
bond.
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O
O

O
O
Single , Double , or Triple?
What is the structural formula for Oxygen?
OO
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Triple Covalent Bond

Each needs 3 electrons to fill outer shell.
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What is the structural formula for Nitrogen?
N
N
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What do you think the structural formula of
methane would be? (CH4)
Hint Start by writing the Lewis symbol for
carbon.
H
C
H
H
H
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H
H
C
H
H
C
H
H

H
H
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Coordinate Covalent Bonds(Also called Dative
Bonds)

• When bonds form between atoms that both donate an
  electron, you have

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Example
2 electrons (lone pair)
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Commonly occurs with transition metals to form
Complex ions. (Gives rise to colors of aqueous
solutions of transition metals)
Diamminesilver ion
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Coordinate Complex
Cu(NH3)42
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Before we discuss these Lewis STRUCTURES in more
detail, we need to discuss bond polarity and its
dependence upon electronegativity.
52
Nonpolar Covalent Bond electrons are shared
equally
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Polar Covalent Bond one atom in the molecule has
a stronger attraction for the bonding electrons.
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Electronegativity
the ability of an atom in a molecule to compete
for electrons.
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Electronegativities according to Pauling Scale
(Relative Scale from 0.7 - 4.0)
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It is the difference in electronegativity between
two atoms that allows us to determine the
polarity and ionic character of a bond.
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We dont consider any numerical fine line in
classifying the different bonds. Instead, we
consider whether a particular bond has ionic
character or covalent character.
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Most compounds with a difference gt 1.5 have more
ionic character than covalent character.
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Use the table below to determine which of the
following bonds are more polar.
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H O or C Cl
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H O or C Cl
3.5 2.1 1.4 3.0 2.5 .5
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H O or C Cl
3.5 2.1 1.4 3.0 2.5 .5
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The symbols used to indicate polarity are delta
plus and delta minus.
H O
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Polarity is also indicated using the following
symbol
H Cl
Indicates electron density being greater near Cl.
Electron density of 3 types of bonds
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Dipole A term used to describe the occurrence of
a charge separated from a - charge So, it is
basically another term for Polar molecule.
Dipole Moment µ Quantitative measure of a dipole.
(reported in debyes C-m) µ Qr Gets larger as
size of charge and distance between charge
increases.
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Writing Lewis Structures
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Writing Lewis Dot Structures

• The Lewis electron-dot structure of a covalent
  compound is a simple two-dimensional
  representation of the positions of electrons in a
  molecule.

• Bonding electron pairs are indicated by either
  two dots or a dash.
• In addition, these formulas show the positions of
  lone pairs of electrons.

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Writing Lewis Dot Structures

• The following rules allow you to write
  electron-dot structures even when the central
  atom does not follow the octet rule.

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Writing Lewis Dot Structures

• Step 1 Total all valence electrons in the
  molecular formula.

PCl3
26 e- total
5 e-
(7 e-) x 3

• For a polyatomic anion, add the number of
  negative charges to this total.
• For a polyatomic cation, subtract the number of
  positive charges from this total.

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Writing Lewis Dot Structures

• Step 2 Arrange the atoms radially, with the
  least electronegative atom in the center. Draw a
  short line between the central atom and each
  peripheral atom. (A line equals one pair of
  electrons).

Cl
Cl
P
Cl
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Writing Lewis Dot Structures

• Step 3 Distribute the remaining electrons to the
  peripheral atoms to satisfy the octet rule.

Cl
Cl
P
Cl
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Writing Lewis Dot Structures

• Step 4 Distribute any remaining electrons to the
  central atom. If there are fewer than eight
  electrons on the central atom, a multiple bond
  may be necessary.

Cl
Cl
P


Cl

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Writing Lewis Dot Structures

• Try drawing Lewis dot structure for the following
  covalent compound.

SCl2
20 e- total
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Writing Lewis Dot Structures

• Try drawing Lewis dot formulas for the following
  covalent compound.

COCl2
24 e- total
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Writing Lewis Dot Structures

• Note that the carbon has only 6 electrons.
• One of the oxygens must share a lone pair.

COCl2
24 e- total
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Writing Lewis Dot Structures

• Note that the carbon has only 6 electrons.
• One of the oxygens must share a lone pair.

COCl2
24 e- total
18 e- left


O
0 e- left
C
Note that the octet rule is now obeyed.
Cl
Cl
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Delocalized Bonding Resonance

• The structure of ozone, O3, can be represented by
  two different Lewis electron-dot formulas.

or

• Experiments show, however, that both bonds are
  identical.

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Delocalized Bonding Resonance

• According to theory, one pair of bonding
  electrons is spread over the region of all three
  atoms.

• This is called delocalized bonding, in which a
  bonding pair of electrons is spread over a number
  of atoms.

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Delocalized Bonding Resonance

• According to the resonance description, you
  describe the electron structure of molecules with
  delocalized bonding by drawing all of the
  possible electron-dot formulas.

and

• These are called the resonance formulas of the
  molecule.

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Double headed arrows are used to indicate
resonance structures.
If the resonance structure is an ion, brackets
are placed around both structures, and the charge
is included outside of the brackets.
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Draw the Lewis Structure for the nitrate ion.
I will draw it on the board when you are finished.
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Exceptions to the Octet Rule

• Although many molecules obey the octet rule,
  there are exceptions where the central atom has
  more than eight electrons.

• Generally, if a nonmetal is in the third period
  or greater it can accommodate as many as twelve
  electrons, if it is the central atom.
• These elements have unfilled d subshells that
  can be used for bonding.
• Nonmetals with atomic 15 or higher may consist
  of structures with exceptions.

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Exceptions to the Octet Rule

• For example, the bonding in phosphorus
  pentafluoride, PF5, shows ten electrons
  surrounding the phosphorus.

P
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Exceptions to the Octet Rule

• In xenon tetrafluoride, XeF4, the xenon atom must
  accommodate two extra lone pairs.

Xe

85
Draw the Lewis Structure of the following SF4
F
F
F
S
F
34 e-
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Exceptions to the Octet Rule

• There are two other elements where an exception
  may occur.
• B and Be may consist of less than an octet.

Cl
Cl
B
Cl
BCl3 24 e-
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Draw the Lewis Structure of the following BeF2
F
F
Be
16 e-
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Formal Charge and Lewis Structures

• In certain instances, more than one feasible
  Lewis structure can be illustrated for a
  molecule. For example,

H
C
N
C
N
H
or

• The concept of formal charge helps us determine
  the structure that is most stable and most likely
  to form.

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• The formal charge of an atom is determined by the
  following formula
• number of valence electrons
• minus
• ½ number of bonding electrons
• minus
• number of unshared valence electrons

Formal Charge VE - 1/2B - U
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Formal Charge and Lewis Structures

• The most likely structure is the one with the
  least number of atoms carrying formal charge. If
  they have the same number of atoms carrying
  formal charge, choose the structure with the
  negative formal charge on the more
  electronegative atom.

• In this case, the structure on the left is most
  likely correct.

91
The following are possible Lewis Structures for
carbon dioxide. Based on formal charge, which
structure is most likely the correct structure?
0
0
0
O C O
or
O C O
-1
0
1
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Bond Length and Bond Order

• Bond length (or bond distance) is the distance
  between the nuclei in a bond.

Knowing the bond length in a molecule can
sometimes give clues as to the strength of the
bond in a molecule. Covalent radii are values
assigned to atoms such that the sum of the radii
of atoms A and B approximate the A-B bond
length.
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Bond Length and Bond Order

• Table 8.5 on page 305 lists bond lengths for some
  covalent bonds.

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Bond Length and Bond Order

• The bond order, determined by the Lewis
  structure, is the number of pairs of electrons in
  a bond.

• Bond length depends on bond order.
• As the bond order increases, the bond gets
  shorter and stronger.

Bond length Bond energy
C
C
154 pm
346 kJ/mol
C
C
134 pm
602 kJ/mol
C
C
120 pm
835 kJ/mol
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Bond length Bond energy
C
C
154 pm
346 kJ/mol
C
C
134 pm
602 kJ/mol
C
C
120 pm
835 kJ/mol
Since this bond is stronger, it requires more
energy to break this bond. Bond Enthalpy Energy
needed to break a bond in 1 mole of a gaseous
substance.
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The strength of a bond is also related to its
stability. Consider N2 versus Cl2.
N
N
Cl
Cl
Bond Length 1.10 A Bond Enthalpy 941 KJ /
mol More Stable
Bond Length 1.96 A Bond Enthalpy 242 KJ /
mol Less Stable
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Bond enthalpies are determined experimentally
using calorimetry.
Determining bond enthalpies for diatomic
molecules is a straight forward experimental
process. When dealing with large molecules,
individual bond enthalpies can be calculated by
considering the average of each bond in the
molecule Example
98
?H of breaking methane into its original atoms
(atomization) is experimentally determined to be
1660 KJ / mol
H
C
H
H
C (g) 4H (g)
H
Average bond enthalpy of C-H bond is therefore
1660 / 4 415 KJ / mol Experiments of other C-H
containing molecules yield similar results.
99
A Return to Hesss Law Bond enthalpies can be
used to determine heats of reactions
?Hrxn. Allows us to see what causes a reaction
to be endo. or exo.
Basic Concept Reactant Reactant ? Product
Product
Energy required to Break Bonds
Energy released when new bonds form
The difference between the two determines
endothermic versus exothermic.
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Bond Energy
Table 8.4 lists values of some bond
energies. (Handout)

• To illustrate, lets estimate the heat of
  reaction for reaction between methane and
  chlorine gas.

CH4 (g) Cl2 (g) ? CH3Cl (g) HCl (g)
101
CH4 (g) Cl2 (g) ? CH3Cl (g) HCl (g)
Calculation

• In this reaction, one C-H bond and one Cl-Cl bond
  must be broken.
• In turn, one C-Cl bond and one H-Cl bond are
  formed.