Chapter 8 Basic Concepts of
Chemical Bonding
 Lewis Symbols and the Octet rule
 Ionic bonding
 Sizes of ions
 Covalent bonding
 Bond Polarity and Electronegativity
 Drawing Lewis Structures
 Exceptions of the Octet rule
 Strengths of Covalent Bonds
 Oxidation Numbers
1
LEWIS SYMBOLS AND THE OCTET RULE
 It was observed that the electron configuration of many substances
after ion formation was that of an inert gas  octet rule.
Octet rule: Main-group elements gain, lose, or share in chemical
bonding so that they attain a valence octet (eight electrons in an
atoms valence shell).
 E.g. The electron configuration of each reactant in the formation of
KCl gives:
a. K+ is that of [Ar]
b. Cl is also that of [Ar].
 The other electrons in the atom are not as important in determining
the reactivity of that substance.
 The octet rule is particularly important in compounds involving
nonmetals.
2
IONIC BONDING
Ionic bonds are electrically neutral groups that are held together
by the attraction arising from the opposite charges of a cation and
an anion.
 Substances that have ionic bonds in a solid form a salt having high
melting point and high crystallinity.
 Bonding thought of as the result of the combination of neutral atoms
with transfer of an electron(s) from one atom to the other.
 For potassium and chlorine atoms approaching each other we have:
Ei = +418 kJ
K(g)  K+(g) + e
Eea = 349 kJ
Cl(g)+ e  Cl(g)
K(g) + Cl(g)  K+(g) + Cl(g) E = + 69 kJ
 Positive E = energy absorbed  energetically not allowed.
 Driving force must be the formation of the crystalline solid.
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BORN-HABER CYCLE AND LATTICE ENERGIES

Overall energetics for the formation of crystalline solids can be
determined from a Born-Haber cycle that accounts for all of the
steps towards the formation of solid salts from the elements. For
the formation of KCl from its elements we have:
1. Sub of t he metal
K (s)  K(g)
 89.2 kJ
2. Dissocia tion of chlorine
1/ 2 Cl2(g)  Cl(g)
122 kJ
3. Ionization of potassium
K(g)  K   e
 418 kJ
4. Formatio n of chloride ions (Eea )
Cl  e  Cl 
 348.6 kJ
5. Formatio n of solid KCl
K (g) Cl (g)  KCl(s)
 715 kJ
Sum reactions and energies
K(s)  1/ 2Cl2(g)  KCl(s)
 434.4 kJ
 Net energy change of 434 kJ/mol indicates energetically favored.
 Energy for the fifth step is the negative of the
lattice energy: energy required to break ionic bonds and sublime
(always positive).
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LATTICE ENERGIES AND PERIODICITY
 Lattice energy can also be determined from Coulomb’s law:
z1z2
U  k 
d
a. Directly proportional to charge on each ion.
b. Inversely proportional to size of compound (sum of ionic radii).
 Table 6.2 summarizes some of the lattice energies for alkali and
alkaline earth ionic compounds. The lattice energies
a. decrease for compounds of a particular cation with atomic
number of the anion.
b. decrease for compounds of a particular anion with atomic
number of the cation.
E.g. Determine the lattice energy of BaCl2 if the heat of sublimation of
Ba is 150.9 kJ/mol and the 1st and 2nd ionization energies are 502 and
966 kJ/mol, respectively. The heat for the synthesis of BaCl2(s) from
its elements is 806.06 kJ/mol.
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The Covalent Bond
 Repulsive forces of the electrons offset by the
attractive forces between the electrons and
the two nuclei.
 Most stable bond energy and bond distance
characterizes bonds between two atoms.
Strengths of Covalent Bonds:
 Bonds form because their formation produces
lower energy state than when atoms are
separated.
 Breaking bonds increases the overall energy
of the system. Energy for breaking bonds has
a positive sign (negative means that energy is
given off).
H - H (g)  2H(g) H = 436 kJ.
 Ionic vs. Covalent Bonds
1.Ionic compounds have high melting and
boiling points and tend to be crystalline;
2.covalently bound compounds tend to have
lower melting points since the attractive
forces between the molecules are relatively
weak.
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Bond Dissociation Enthalpies
Bond dissociation energy, D – the energy required to break one
mole of a type of bond in an isolated molecule in the gas phase.
 Useful for estimation of heat of unknown reactions.
 Average bond energies listed in tables (e.g. C – H bond); rest pf
structure not very important
1.
HO-H bond in H2O and CH3O-H bond are 492 and 435 kJ/mol.
 Hess’s law can be used with bond dissociation energies to estimate
the enthalpy change of a reaction. The breaking in a C – H bond
would be C – H(g)  C(g) + H(g) H = D = 410 kJ. Sign always
positive since energy must be supplied to break bond.
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Bond Dissociation Enthalpies2
E.g. Estimate the heat of formation of H2O(g) from bond dissociation
energies. Thus determine: H2(g) + ½ O2(g)  H2O(g) Hof = ?
Look in the book (p. 231 Table 7.1):
H – H (g)  2H(g)
H = D1 = 436 kJ
½ O=O  O(g)
H = D2 = 494/2 = 247 kJ
2H(g) + O(g)  H – O – H (g)
H = 2D3 = 2*459 kJ
H2(g) + ½ O2(g)  H2O(g)
Hof = 235 kJ
Actual Hof = 241.8 kJ
 Can be determined by suming all the energies for the bonds broken
and subtract from if the sum of the energies for the bonds formed.
E.g. 2 Estimate the energy change for the chlorination of ethylene:
CH2=CH2(g) + Cl2(g)  CH2ClCH2Cl.
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Bond Dissociation Enthalpies3
It may be necessary to include a phase change since many reactions
or reactants are not in the gas phase.
E.g. Determine the heat of formation of CCl4(l).
Solution: The reaction is: C(gr) + 2Cl2(g)  CCl4(l) Hof = ? Write
the reactions and sum energies:
C(gr)
2 Cl – Cl
C(g) + 4 Cl(g)
CCl4(g)
C(gr) + 2Cl2(g)
Actual is  139 kJ.
 C(g)
 4Cl(g)
 CCl4(g)
 CCl4(l)
 CCl4(l)
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H1 = 715 kJ
H2 = 486
H3 = 1320
H4 = 43
H = 162 kJ
Lewis Structures
Lewis structure valence electrons represented by dots and are
placed where they would be in any bonding that might exist.
Lewis structures of second row elements:
H2
BH3
CH4
NH3
H2O
HF
 Each has 8 electrons around the central atom; thus we can predict
the number of bonds that will form from the position in the periodic
.. ..
:Cl
:Cl
:
..
..
table. E.g. The structure of chlorine is
a) shared electrons are called "bonding" electrons
b) unshared electrons are called "non-bonding" or "lone pair"
electrons.
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Lewis Structures(cont’d)
 Octet can be filled by donation of electrons from each atom or one
atom can supply both electrons.
E.g. H+ + NH3  NH
4 . "co-ordinate covalent bond".


BF

F

BF
3
4
E.g.2
 Multiple bonds may form as a result when the two atoms forming
the bond do not have enough electrons.
a) O=O
b) NN
 Multiple bonds are shorter and stronger than single bonds because
of the extra electrons holding the two atoms together.
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Lewis Structures of Polyatomic Molecules
Procedure for more complicated molecules:
1 Determine the total number of valence electrons from each atom.
2 Distributed atoms around the central atom (least electronegative.
Hydrogen atoms are usually attached to any oxygen.
3 Satisfy the octet of the atoms bonded to the central atom.
4 Satisfy the octet of the central atom by distributing the remaining
electrons as electron pairs around it. (multiple bonds may be
necessary)
E.g. Determine the Lewis structure of H2SO4.
E.g. Draw the Lewis dot structures of XeF4, ICl3, and XeOF4.
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Lewis Structures and Resonance
 Quantum theory indicates that any position possible for an electron.
 Equivalent electron positions often possible:
E.g. SO2: O=S-O: and :O-S=O. Each structure equally likely.
a) the true form of the molecule is a hybrid of these and is called
resonance and the hybrid form is called a resonance hybrid.
13
Polar Bonds: Electronegativity
Electronegativity is a measure of the atom’s ability to gain or lose
electrons. It is directly related to its ionization tendency and its ability to
Ei  Eea
form the inert gas configuration. Obtained by:  
where Ei =
2
ionization energy and Eea = the electron affiinity.
E.g. Li has a very low ionization energy and electron affinity, while Cl has a
both a high ionization energy and high electron affinity. Electronegativity will
be high for Cl and low for Li.
 Fluorine has the highest electronegativity of 4.0.
 Electronegativities increase from bottom to top of periodic table and
increase to a maximum towards the top right.
 Combination of elements with intermediate electronegativities forms
bonds that are intermediate between covalent and ionic.
 can provide an insight as to the type of bond that would be expected.
1. Ionic bonds formed when   2
2. covalent bonds forms when   1.
3. Polar covalent forms when 1    2, the bonding is "intermediate"
between the two.
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Polar Bonds: Electronegativity2
E.g.1 Determine the polarity of the N – H in NH3.
E.g. 2 Predict the type of bond formed in CCl4.
 The magnitude of  indicates if electrons are polarized around one
element in preference to the other.
 bond polar. With intermediate , a small charge on the atom due to that
bond develops. + and  designates which is the positive and negative
side respectively.
E.g.3 Determine the relative polarities of HF, HCl, HBr and HI.
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FORMAL CHARGES
Formal Charge (of an atom in a Lewis formula) the hypothetical
charge obtained by assuming that bonding electrons are equally
shared between the two atoms involved in the bond. Lone pair
electrons belong only to the atom to which they are bound.
1
FC  # Val e   # Bonding e   # nonbonding e 
2
E.g. determine the formal charge on all elements: PCl3, PCl5, and
HNO3.
 formal charge (FC)allows the prediction of the more likely
resonance structure.
To determine the more likely resonance structure:
1.FC should be as close to zero as possible.
2.Negative charge should reside on the most electronegative and
positive charge on the least electronegative element.
E.g. draw the resonance structures of H2SO4; determine the formal
charge on each element and decide which is the most likely
structure.
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