CHEMISTRY
The Central Science
9th Edition
Chapter 8
Basic Concepts of Chemical
Bonding
David P. White
Prentice Hall © 2003
Chapter 8
Chemical Bonds, Lewis
Symbols, and the Octet
Rule
• Chemical bond: attractive force holding two or more
atoms together.
• Covalent bond results from sharing electrons between
the atoms. Usually found between nonmetals.
• Ionic bond results from the transfer of electrons from a
metal to a nonmetal.
• Metallic bond: attractive force holding pure metals
together.
Prentice Hall © 2003
Chapter 8
Chemical Bonds, Lewis
Symbols, and the Octet
Rule
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Lewis Symbols
As a pictorial understanding of where the electrons are in
an atom, we represent the electrons as dots around the
symbol for the element.
The number of electrons available for bonding are
indicated by unpaired dots.
These symbols are called Lewis symbols.
We generally place the electrons one four sides of a
square around the element symbol.
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Chapter 8
Chemical Bonds, Lewis
Symbols, and the Octet
Rule
Lewis Symbols
Prentice Hall © 2003
Chapter 8
Chemical Bonds, Lewis
Symbols, and the Octet
Rule
The Octet Rule
• All noble gases except He has an s2p6 configuration.
• Octet rule: atoms tend to gain, lose, or share electrons
until they are surrounded by 8 valence electrons (4
electron pairs).
• Caution: there are many exceptions to the octet rule.
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Chapter 8
Ionic Bonding
Consider the reaction between sodium and chlorine:
Na(s) + ½Cl2(g)  NaCl(s)
DHºf = -410.9 kJ
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Chapter 8
Ionic Bonding
• The reaction is violently exothermic.
• We infer that the NaCl is more stable than its constituent
elements. Why?
• Na has lost an electron to become Na+ and chlorine has
gained the electron to become Cl-. Note: Na+ has an Ne
electron configuration and Cl- has an Ar configuration.
• That is, both Na+ and Cl- have an octet of electrons
surrounding the central ion.
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Chapter 8
Ionic Bonding
• NaCl forms a very regular structure in which each Na+
ion is surrounded by 6 Cl- ions.
• Similarly, each Cl- ion is surrounded by six Na+ ions.
• There is a regular arrangement of Na+ and Cl- in 3D.
• Note that the ions are packed as closely as possible.
• Note that it is not easy to find a molecular formula to
describe the ionic lattice.
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Chapter 8
Ionic Bonding
Prentice Hall © 2003
Chapter 8
Ionic Bonding
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Energetics of Ionic Bond Formation
The formation of Na+(g) and Cl-(g) from Na(g) and Cl(g)
is endothermic.
Why is the formation of Na(s) exothermic?
The reaction NaCl(s)  Na+(g) + Cl-(g) is endothermic
(DH = +788 kJ/mol).
The formation of a crystal lattice from the ions in the gas
phase is exothermic:
Na+(g) + Cl-(g)  NaCl(s) DH = -788 kJ/mol
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Chapter 8
Ionic Bonding
Energetics of Ionic Bond Formation
• Lattice energy: the energy required to completely
separate an ionic solid into its gaseous ions.
• Lattice energy depends on the charges on the ions and the
sizes of the ions:
Q1Q2
El  k
d
k is a constant (8.99 x 10 9 J·m/C2), Q1 and Q2 are the
charges on the ions, and d is the distance between ions.
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Chapter 8
Ionic Bonding
Energetics of Ionic Bond Formation
• Lattice energy increases as
• The charges on the ions increase
• The distance between the ions decreases.
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Chapter 8
Ionic Bonding
Electron Configurations of Ions of the
Representative Elements
• These are derived from the electron configuration of
elements with the required number of electrons added or
removed from the most accessible orbital.
• Electron configurations can predict stable ion formation:
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Mg: [Ne]3s2
Mg+: [Ne]3s1
not stable
Mg2+: [Ne]
stable
Cl: [Ne]3s23p5
Cl-: [Ne]3s23p6 = [Ar] stable
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Chapter 8
Ionic Bonding
Transition Metal Ions
• Lattice energies compensate for the loss of up to three
electrons.
• In general, electrons are removed from orbitals in order
of decreasing n (i.e. electrons are removed from 4s before
the 3d).
Polyatomic Ions
• Polyatomic ions are formed when there is an overall
charge on a compound containing covalent bonds. E.g.
SO42-, NO3-.
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Chapter 8
Covalent Bonding
• When two similar atoms bond, none of them wants to
lose or gain an electron to form an octet.
• When similar atoms bond, they share pairs of electrons to
each obtain an octet.
• Each pair of shared electrons constitutes one chemical
bond.
• Example: H + H  H2 has electrons on a line connecting
the two H nuclei.
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Chapter 8
Covalent Bonding
Prentice Hall © 2003
Chapter 8
Covalent Bonding
Lewis Structures
• Covalent bonds can be represented by the Lewis symbols
of the elements:
Cl + Cl
Cl Cl
• In Lewis structures, each pair of electrons in a bond is
represented by a single line:
H
H O
H N H
Cl Cl
H F
H C H
H
H
H
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Chapter 8
Covalent Bonding
Multiple Bonds
• It is possible for more than one pair of electrons to be
shared between two atoms (multiple bonds):
• One shared pair of electrons = single bond (e.g. H2);
• Two shared pairs of electrons = double bond (e.g. O2);
• Three shared pairs of electrons = triple bond (e.g. N2).
H H
O O
N N
• Generally, bond distances decrease as we move from
single through double to triple bonds.
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Chapter 8
Bond Polarity and
Electronegativity
• In a covalent bond, electrons are shared.
• Sharing of electrons to form a covalent bond does not
imply equal sharing of those electrons.
• There are some covalent bonds in which the electrons are
located closer to one atom than the other.
• Unequal sharing of electrons results in polar bonds.
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Chapter 8
Bond Polarity and
Electronegativity
Electronegativity
• Electronegativity: The ability of one atoms in a
molecule to attract electrons to itself.
• Pauling set electronegativities on a scale from 0.7 (Cs) to
4.0 (F).
• Electronegativity increases
• across a period and
• down a group.
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Chapter 8
Bond Polarity and Electronegativity
Electronegativity
Bond Polarity and
Electronegativity
Electronegativity and Bond Polarity
• Difference in electronegativity is a gauge of bond
polarity:
• electronegativity differences around 0 result in non-polar
covalent bonds (equal or almost equal sharing of electrons);
• electronegativity differences around 2 result in polar covalent
bonds (unequal sharing of electrons);
• electronegativity differences around 3 result in ionic bonds
(transfer of electrons).
Prentice Hall © 2003
Chapter 8
Bond Polarity and
Electronegativity
Electronegativity and Bond Polarity
• There is no sharp distinction between bonding types.
• The positive end (or pole) in a polar bond is represented
+ and the negative pole -.
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Chapter 8
Bond Polarity and
Electronegativity
Dipole Moments
• Consider HF:
• The difference in electronegativity leads to a polar bond.
• There is more electron density on F than on H.
• Since there are two different “ends” of the molecule, we call HF
a dipole.
• Dipole moment, m, is the magnitude of the dipole:
m  Qr
where Q is the magnitude of the charges.
• Dipole moments are measured in debyes, D.
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Chapter 8
Bond Polarity and
Electronegativity
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Bond Types and Nomenclature
In general, the least electronegative element is named
first.
The name of the more electronegative element ends in
–ide.
Ionic compounds are named according to their ions,
including the charge on the cation if it is variable.
Molecular compounds are named with prefixes.
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Chapter 8
Prentice Hall © 2003
Chapter 8
Bond Polarity and
Electronegativity
Bond Types and Nomenclature
Ionic
Molecular
MgH2
Magnesium hydride H2S
Hydrogen sulfide
FeF2
Iron(II) fluoride
Oxygen difluoride
Mn2O3 Manganese(III)
oxide
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OF2
Cl2O3 Dichlorine trioxide
Chapter 8
Drawing Lewis Structures
1. Add the valence electrons.
2. Write symbols for the atoms and show which atoms are
connected to which.
3. Complete the octet for the central atom the complete the
octets of the other atoms.
4. Place leftover electrons on the central atom.
5. If there are not enough electrons to give the central atom
an octet, try multiple bonds.
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Chapter 8
Drawing Lewis Structures
Formal Charge
• It is possible to draw more than one Lewis structure with
the octet rule obeyed for all the atoms.
• To determine which structure is most reasonable, we use
formal charge.
• Formal charge is the charge on an atom that it would
have if all the atoms had the same electronegativity.
Prentice Hall © 2003
Chapter 8
Drawing Lewis Structures
Formal Charge
• To calculate formal charge:
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All nonbonding electrons are assigned to the atom on which
they are found.
Half the bonding electrons are assigned to each atom in a
bond.
• Formal charge is:
valence electrons - number of bonds - lone pair electrons
Prentice Hall © 2003
Chapter 8
Drawing Lewis Structures
Formal Charge
• Consider:
C N
• For C:
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There are 4 valence electrons (from periodic table).
In the Lewis structure there are 2 nonbonding electrons and 3
from the triple bond. There are 5 electrons from the Lewis
structure.
Formal charge: 4 - 5 = -1.
Prentice Hall © 2003
Chapter 8
Drawing Lewis Structures
Formal Charge
• Consider:
C N
• For N:
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There are 5 valence electrons.
In the Lewis structure there are 2 nonbonding electrons and 3
from the triple bond. There are 5 electrons from the Lewis
structure.
Formal charge = 5 - 5 = 0.
• We write:
C N
Prentice Hall © 2003
Chapter 8
Drawing Lewis Structures
Formal Charge
• The most stable structure has:
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the lowest formal charge on each atom,
the most negative formal charge on the most electronegative
atoms.
Resonance Structures
• Some molecules are not well described by Lewis
Structures.
• Typically, structures with multiple bonds can have
similar structures with the multiple bonds between
different pairs of atoms
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Chapter 8
Drawing Lewis Structures
Resonance Structures
• Example: experimentally, ozone has two identical bonds
whereas the Lewis Structure requires one single (longer)
and one double bond (shorter).
O
O
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Chapter 8
O
Drawing Lewis Structures
Resonance Structures
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Chapter 8
Drawing Lewis Structures
Resonance Structures
• Resonance structures are attempts to represent a real
structure that is a mix between several extreme
possibilities.
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Chapter 8
Drawing Lewis Structures
Resonance Structures
• Example: in ozone the extreme possibilities have one
double and one single bond. The resonance structure has
two identical bonds of intermediate character.
O
O
O
O
O
O
• Common examples: O3, NO3-, SO42-, NO2, and benzene.
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Chapter 8
Drawing Lewis Structures
Resonance in Benzene
• Benzene consists of 6 carbon atoms in a hexagon. Each
C atom is attached to two other C atoms and one
hydrogen atom.
• There are alternating double and single bonds between
the C atoms.
• Experimentally, the C-C bonds in benzene are all the
same length.
• Experimentally, benzene is planar.
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Chapter 8
Drawing Lewis Structures
Resonance in Benzene
• We write resonance structures for benzene in which
there are single bonds between each pair of C atoms and
the 6 additional electrons are delocalized over the entire
ring:
or
• Benzene belongs to a category of organic molecules
called aromatic compounds (due to their odor).
Prentice Hall © 2003
Chapter 8
Exceptions to the Octet
Rule
• There are three classes of exceptions to the octet rule:
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Molecules with an odd number of electrons;
Molecules in which one atom has less than an octet;
Molecules in which one atom has more than an octet.
Odd Number of Electrons
• Few examples. Generally molecules such as ClO2, NO,
and NO2 have an odd number of electrons.
N O
N O
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Chapter 8
Exceptions to the Octet
Rule
Less than an Octet
• Relatively rare.
• Molecules with less than an octet are typical for
compounds of Groups 1A, 2A, and 3A.
• Most typical example is BF3.
• Formal charges indicate that the Lewis structure with an
incomplete octet is more important than the ones with
double bonds.
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Chapter 8
Exceptions to the Octet
Rule
More than an Octet
• This is the largest class of exceptions.
• Atoms from the 3rd period onwards can accommodate
more than an octet.
• Beyond the third period, the d-orbitals are low enough in
energy to participate in bonding and accept the extra
electron density.
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Chapter 8
Strengths of Covalent
Bonds
• The energy required to break a covalent bond is called
the bond dissociation enthalpy, D. That is, for the Cl2
molecule, D(Cl-Cl) is given by DH for the reaction:
Cl2(g)  2Cl(g).
• When more than one bond is broken:
CH4(g)  C(g) + 4H(g) DH = 1660 kJ
•
the bond enthalpy is a fraction of DH for the
atomization reaction:
D(C-H) = ¼DH = ¼(1660 kJ) = 415 kJ.
• Bond enthalpies can either be positive or negative.
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Chapter 8
Strengths of Covalent
Bonds
Bond Enthalpies and the Enthalpies of
Reactions
• We can use bond enthalpies to calculate the enthalpy for
a chemical reaction.
• We recognize that in any chemical reaction bonds need
to be broken and then new bonds get formed.
• The enthalpy of the reaction is given by the sum of bond
enthalpies for bonds broken less the sum of bond
enthalpies for bonds formed.
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Chapter 8
Strengths of Covalent
Bonds
Bond Enthalpies and the Enthalpies of
Reactions
• Mathematically, if DHrxn is the enthalpy for a reaction,
then
DH rxn   Dbonds broken  -  Dbonds formed 
• We illustrate the concept with the reaction between
methane, CH4, and chlorine:
CH4(g) + Cl2(g)  CH3Cl(g) + HCl(g)
DHrxn = ?
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Chapter 8
Strengths of Covalent Bonds
Strengths of Covalent
Bonds
Bond Enthalpies and the Enthalpies of
Reactions
• In this reaction one C-H bond and one Cl-Cl bond gets
broken while one C-Cl bond and one H-Cl bond gets
formed.
DH rxn  DC - H   DCl - Cl  - DC - Cl   DH - Cl 
 -104 kJ
• The overall reaction is exothermic which means than the
bonds formed are stronger than the bonds broken.
• The above result is consistent with Hess’s law.
Prentice Hall © 2003
Chapter 8
Strengths of Covalent
Bonds
Bond Enthalpy and Bond Length
• We know that multiple bonds are shorter than single
bonds.
• We can show that multiple bonds are stronger than
single bonds.
• As the number of bonds between atoms increases, the
atoms are held closer and more tightly together.
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Chapter 8
Strengths of Covalent Bonds
End of Chapter 8:
Basic Concepts of Chemical
Bonding
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Chapter 8