Chapter 8
Chemical Bonds
Models for chemical bonds
(pure) ionic bond --- most compounds --- (pure) covalent bond
Ions – held together by electrostatic forces, ionic lattices.
Atoms sharing pairs of electrons  covalent bonds, molecules.
NaCl(s) – Na+, Cl- in geometric pattern  crystal lattice.
Lewis symbols for atoms/ions:
Chemical symbol – for molecules + inner/core electrons
Dot – for each valence electron, paired or unpaired
e. g,
H
He
N
N
O
Cl
K
Mg
[He] 2s2 2px1 2py1 2pz1
Ionic Compounds:
 Remove valence electrons from cation (metal)
 Add valence electrons to anion (non-metal)
 Adjust numbers of cations/anion; include charges
Potassium Chloride
K
+
K
Cl
+
Cl
Lewis formula
Calcium chloride
Cl
Ca2+
Cl
Octet of electrons, ns2 np6 in anion/cation; except in 1st shell, n= 1; 1s2: H+ or H:-,
Li+, Be+ (1s2)
Anion  closed-shell configuration of following rare gas
1
Cation  closed-shell configuration of preceding rare gas or [ ] nd10 or [ ] (n – 1)
d10 ns2
Aluminum oxide:
Al +
O
2 [Al3+] + 3 O
2
Al2O3
Two Indium fluorides
F
[In:+][
In [Kr] 4d10 5s2 5p1
F ]
and
[In3+] 3 [ F ]
= In F3
(Another example, PbO and PbO2)
Driving force (reasons) for ionic compounds is the very favorable lattice
energy/enthalpy for ions in crystal arrangement. “more ionic” compounds have
most favorable lattice energies:
LiF (1046), LiCl (861), LiBr (818), LiI(751 kJ/mol)
Highest lattice energies between small, highly-charged ions:
KCl (717 kJ/mol)
MgO (3850 kJ/mol)
[Ions stack together in regular, 3-dimensional arrays to produce crystalline solids;
NaCl(s), no molecules, but each Na+ surrounded by Cl- ions (6 Cl- as “nearest
neighbors”) and each Cl- surrounded by 6 Na+ as nearest neighbors – “almost
infinite array of ions”. Crystalline solids – high melting, brittle, may dissolve to
produce ions (electrolyte)
NaCl(s)  Na+(aq) + Cl(aq)
Covalent Bonds
Covalent bond = pair of electrons shared between two atoms.
H2 molecule, H-H covalent bond
2
1s1
1s1
-
both electrons shared by/attracted to both nuclei
for 1s electrons, n = 1 , l = 0, ml = 0; must have paired spins.
H + H
H H or H-H , covalent bond in a Lewis structure.
Lewis Structures/Octet Rule
Atoms tend to share electrons in covalent bonds until a closed-shell configuration
is obtained  octet rule
ns2 np6 – true for 2nd period only; 1st period, n = 1, need two electrons;
3rd, 4th period of elements (P, S, Cl, As… etc) may have “expanded octets”
N --- N
, “needs” 3 electrons, makes 3 covalent bonds, “valence of 3”
Valence = number of covalent bonds formed
Cl-----
Cl
Ar ---
Ar
, “needs” one electron, makes one covalent bond, valence of
1
, not expected to form covalent bonds
H
H
F2
H H or H-H hydrogen expected to form 1 covalent bond
F
F
F
, can form a covalent bond --F
F
More simply,
Lone pairs or non-bonding electron pairs on F atom – valence electrons not used
in covalent bonding, normally spin-paired.
-
lone pair repulsion on F atoms weakens F-F covalent bond, making F2
very reactive – more reactive than Cl2, Br2, I2
3
no lone pair
HBr,
H
Br
3 lone pairs
(6 electrons)
H- Br
Polyatomic Species –
Methane, CH4
C
+
4H
4 + 4(1) = 8 valence electrons
(need to “unpair” C valence electrons to make 4 covalent bonds)
H
H
H C H
H
=
H C H
H
Lewis structure (C is tetrahedral in most
compounds)
Single bond = 1 shared pair of electrons
F
H H or H-H
F
Double Bond = 2 shared electron pairs between same atoms
C
O
C
O
; C
C
C
C
4 electrons shared
Triple Bond = 3 shared electrons pairs between two atoms
C
C
C
C
, six electrons shared; multiple bonds
Selecting central atom in polyatomic molecule/ion:
a. Element with lowest ionization energy (lowest electronegativity) – so P,
As, S, even Cl, Br, I; more apt than O, F; never H
b. Bonded atoms usually symmetric around central atom; SO2  O-S-O
O
O
SO3
SO42-
S
O
O
O
S
O
2
O
OF2  F-O-F
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c. Central atom usually written first: PO43, SO42, NO3, NO2 etc
d. Oxoacids – acidic H bonded to O, not central atom
H2SO4  (HO)2SO2;
HClO2  (HO)ClO;
H3PO4  (HO)3PO
HClO4  (HO)ClO3
Ions: Adjust valence electrons (dots) for electrons lost (cations) or gained
(anions)
+
H
H
5 + 4(1) – 1 = 8
NH4+
N
H
H
2
O
O
SO426 + 4(6) + 2 = 32 electrons/dots
16 electron pairs/pairs of dots
S
O
O
Treat ions separately; cations/anions – are not linked by electron pairs
Example: (NH4)2 SO4
ionic/electrostatic attraction
Charge belongs to whole ion, not individual atoms.
Examples of Lewis Structures for Polyatomic Species
HCN
1. Count valence electrons/electron pairs:
1 + 4 + 5 = 10 electrons ( 5 pairs)
2. Write down skeleton/arrangement of atoms, putting single bonds between
bonded atoms: H-C-N
3. Complete octet using all remaining valence electrons; use multiple bonds if
apparently electron deficient.
H C
N = H C N
CS2
5
1. 4 + 2(6) = 16 electrons, 8 pairs
2. S-C-S
S
C S
3.
HNO3
1. 1 + 5 + 3(6) = 24 electrons / 12 pairs
O
H
O
N
O
2.
O
H
O
N
O
3.
CO32-
1. 4 + 3(6) + 2 = 24 electrons / 12 pairs
O
O
C
O
2.
2
O
O
C
O
3.
OF2
1. 6 + 2(7) = 20 electrons / 10 pairs
2. F-O-F
F
3.
O
F
CO2 vs CO
O
C
O
C
O
Resonance / Resonance Structures
6
Sometimes two or more valid Lewis structures possible for molecules / ions:
NO3
5 + 3(6) +1 = 24 electrons /12 pairs
O
O
N
O
O
O
O
N
O
O
N
O
Individual structures: 2 single N-O bond (140 pm), 1 double N=O bond (120 pm);
experimental: 3 equivalent N-O bonds (124 pm).
Actual structure = blend/averaging of structures, not resonating
Delocalized multiple bonds, between single and double
O3, ozone: 3(6) = 18 electrons / 9 pairs
Equivalent bond lengths
O
CH3CO2:
O
O
O O O
2(4) + 3(1) + 2(6) + 1 = 24 electrons / 12 pairs
H
O
H C
C
O
H
H
O
H C
C
O
O
C
H
O
“equivalent resonance structures” ; non-equivalent structures
SCN, thiocyanate: 6 + 4 + 5 + 1 = 16 e / 8 pair
S C N
S
C
N
S C N
Equivalent resonance structures have same energy, contribute equally to
resonance hybrid; non-equivalent structures have different energies and
contribute unequally to hybrid – lower energy structure contribute more.
CH3CH2OH, CH3OCH3 are isomers, not resonance structures, because atom
arrangement is changed.
Formal Charge
Assignment of electrons to individual atoms in Lewis structure leads to “apparent
changes” on atoms = formal charge
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Shared (bonding) electrons divided equally between atoms bonded; lone pair
electrons assigned to that atom.
Formal Charge = valence electrons – non-bonding electrons – ½ bonding
electrons
FC = V – (L + ½ S)
Example:
NO3, choose on equivalent resonance structure:
-1
-1
O
0
O
N
O
+1
N
FC = 5 – [0 + ½ (8)] = +1
-O
FC = 6 – [ 6 + ½ (2)] = -1 x 2
=O
FC = 6 – [4 + ½ (4) ] = 0
Overall charge = -1
0
+1
-1
O
O
O
O1 – FC = 6 – 4 – ½ (4) = 0
O2 – FC = 6-2 – ½ (6) = +1
O3 – FC = 6 – 6 – ½ (2) = -1
Use FC to predict/anticipate “most reasonable resonance structure”
S C N
S
C
N
FC = 6 – 6 – ½ (2) = -1
FC = 4 – 0 ½ (6) = 0
FC = 5 – 2 – ½ (6) = 0
S
C
N
6 – 2 -3 = +1
4–0–0=0
5 – 6 – 1 = -2
S C N
6–4–2=0
4–0–0=0
5 – 4 – 2 = -1
Which formal charges are “most reasonable”?
a. Best Lewis structure(s) have formal charges closes to zero
b. “Good Lewis structures” must have formal charges that are
reasonable with regard to electronegativity values.
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Electronegativity, ability of an atom to attract (pull) electrons to itself in a covalent
bond. Several different “scales”, Pauling electronegativities, Mulliken Scale ½ (I1
+ EA)
Electronegativity Trends
a. decrease down groups F > Cl > Br > I
b. Increases from left to right (except noble gases): Li < Be < B < C < N < O
<F
c. Small diagonal difference, upper left > lowr right
C > P, N > S, O > Cl
d. metals have low electronegativity values: They are normally referred to as
“electropositive” rather than electronegative.
cH, 2.6 vs. 2.2; C-H bonds are non-polar

Most electronegative elements: F > O > N > Br
S C N
-1 0 0
Reasonable
S
C
N
S C N
+1 0 -2
“unreasonable”
0 0 -1
best
Expanded Valence Shells / the Expanded Octet
Common formulas: PF5, SF6, PF6, IF5, IF7
F
F
F
F
PF5; 5 + 5(7) = 40 electrons or 20 pair
F
P
F
F
F
S
F
F
SF6; 6 + 6(7) = 48 electrons or 24 pairs
F
Central atom in 3rd, 4th, 5th period: P, S, As, Se, Sb, Te, Cl, Br, I have availability
of empty nd orbitals; also larger atomic size.
PCl5(s) = [PCl4]+[PCl6]
5 + 4 (7) -1 = 32 electrons/16 pairs
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+
Cl
Cl
Cl
P
Cl
Cl
Cl
P
Cl
Cl
Cl
Cl
Octet
expanded octet
Variable covalence = ability to form different numbers of covalent bonds. P forms
3, 4, 5, 6 covalent bonds.
P
Cl
Cl
Cl
PCl3 =
More Subtle Examples:
F
F
S
F
F
SF4: 6 + 4 (7) = 34 electrons / 17 pairs
F
F
Xe
F
F
XeF4: 8 + 4(7) = 36 electrons / 18 pairs
Sometimes expanded octet structures are better resonance structures than
others:
SO42
2
-1
O
-1
O
S
+2
O
0
O
S
O
6 - o - 4 = +2
6 - 6 - 1 = -1
2
-1
O
O
-1
+1
-1
FC: S
-O
=O
2
-1
-1
O
0
O
S
6 - 0 - 5 = +1
6 - 6 - 1 = -1
6 - 4 -2 = 0
O
-1
-1
O
0
0
O
6-0-6=0
6 - 6 - 1 = -1
6 - 4 -2 = 0
10
2nd best
Poorest, but possible
best
2
0
O
0
O
S
0
-2
O
0
S
6 - 0 - 8 = -2
unreasonable
=O 6 - 4 - 2 = 0
O
SO2:
O
S
O
O
S
O
equivalent structures
FC: S 6 - 2 - 3 = +1
-O 6 - 6 - 1 = -1
=O 6 - 4 -2 = 0
O
S
O
no-equivalent structure
6-2-4=0
6-4-2=0
Ionic vs Covalent Bonds
A - B = | A = B|
 = 0, purely covalent, equally shared electron pair = non – polar bond
0 <  < 2, increasingly polar covalent bond = unequal sharing
≥ 2 ionic bond; electron transfer
“% ionic” or “% covalent” nature of bonds.
NaCl, = | 0.93 – 3.2 | = 2.3, ionic
CsF,  = | 0.79 – 4.0 | = 3.2 “more ionic”
BeCl,  = | 1.6 – 3.2 | = 1.6 highly polar  ionic bond
Al-Cl,  = | 1.6 – 3.2 | = 1.6 highly polar Al2Cl6(s)
Al-F  = | 1.6 – 4.0 | = 2.4 ionic AlF3(s)
C-H, = | 2.6 – 2.2 | 0.4 essentially non-polar
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