Chapter 8 Basic Concepts of Chemical Bonding
Chemical Bonds, Lewis Symbols, and the Octet Rule
ionic bonds - electrostatic forces
covalent bonds - sharing of electrons between two atoms
metallic bonds - each atom is bonded to several neighboring atoms
Lewis Symbols
Chemical symbol + valence electrons
.
S
2
[Ne] 3s 3p4 :S:
∙
Number of valence electrons = Group number (s and p elements)
Octet Rule
Atoms tend to gain, lose, or share electrons until they are surrounded by eight
valence electrons.
Chapter 8 Basics of Chemical Bonding 1
Electron Configurations () of Ions of Representative Elements
∙Na ≈ 1s22s22p63s1 ≡ [Ne] 3s1
∙Na+ ≈ 1s22s22p6 ≡ [Ne]
∙Cl ≈ 1s22s22p63s23p5 = [Ne] 3s23p5
∙Cl– ≈ 1s22s22p63s23p6 = [Ne] 3s23p6
= [Ar]
Transition metals
In forming ions, transition metals lose the valence shell s electrons first (first in,
first out) then as many d electrons as are required to reach the charge on the ion.
∙Fe ≈ [Ar]3d64s2
∙Fe2+ ≈ [Ar]3d6
∙Fe3+ ≈ [Ar]3d5
Sizes of Ions
Cations are smaller than their parent atoms
removal of outermost electrons &fewer electron-electron repulsions
Anions are larger than their parent atoms
increased electron-electron repulsions
Isoelectronic series “same number of electrons”, same el. Config 
increasing nuclear charge
8
O2–
1.40
 ≈ 1s22s22p6
9
11
12
13
F–
1.33
Na+
0.97
M g2+
0.66
Al3+
0.50 Å
decreasing radius
Chapter 8 Basics of Chemical Bonding 2
Cations from atoms
Order of electron loss: p then s then d (p_s_d)
ns2 occupation explains why main group metal cations vary by 2+
e.g. In+,In3+; Pb2+, Pb4+; Sn2+, Sn4+; Sb3+, Sb5+ ; Bi3+,Bi5+
Chapter 8 Basics of Chemical Bonding 3
Ionic Bonding
Nas  12 Cl2 g  NaCl(s) H f  410.9 kJ mol
Na
+
Cl  Na+ + [ Cl ]–
2s22p6
3s23p6
Lattice Energy is the 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
Magnitude depends on size of charge and size of ions (Coulomb’s Law)
Ek
Q1Q2
d
Qi = charge on ion, d = internuclear distance or sum of ionic radii
Estimating relative lattice energies of solids
1) First compare Q,
2) then d.
i) KCl and CaS
QiQj = +1 x -1 = -1 for KCl and QiQj = +2 x -2 = -4 for CaS
EKCl = k ∙ (-1) / dKCl
ECaS = k ∙ (-4) / dCaS
But dKCl = 138pm ∙ 181pm = 2.5x104 pm2 and dCaS = 100pm ∙ 184pm = 1.8x104
pm2
dKCl ≈ 2x104 pm2 and dCaS ≈ 2x104 pm2
Chapter 8 Basics of Chemical Bonding 4
So, QiQj differences predominate over d differences.
ii) CaO and MgO,
QiQj both = -4, but r(Mg2+) = 72pm and r(Ca2+) = 100pm and dCaO > dMgO and ECaO
< EMgO
Self-test 2.3 A and B
Ionic solids typically have high melting points andare brittle. The interaction
between ios is larger when the charges are greater and sizes are smaller
Covalent Bondingsharing pairs of electrons
Chapter 8 Basics of Chemical Bonding 5
Lewis Structures
H
+
Cl
+
H
H H
Cl
Cl Cl
each shared pair of electrons is drawn as a line
H H
H F
Cl
H O
Cl
H N H
H
C
H
The number of bonds an atom normally forms is called the VALENCY.
Multiple Bonds
If two pairs of electrons are shared, it is a double bond
If three pairs of electrons are shared, it is a triple bond
N +
N
N N
The aim is to give an octet of electrons on each atom
Bond distances decrease as the number of shared electron pairs increase because
there are more pairs of e- between the positive nuclei.
(Å=10-10m = 100pm = 0.1nm)
N–N
N=N
NN
1.47 Å
1.24 Å
1.10 Å
Chapter 8 Basics of Chemical Bonding 6
Drawing Lewis Structures (Tool Box 2.1)
1. Sum the valence electrons from all atoms and charge on anion (-1 means + 1 e-)
2. Write the symbol for the atoms to show which atoms are attached to which and
connect them with single bonds
3. Complete the octets of the atoms bonded to the central atom
4. Place any left over electrons on the central atom.
5. If there are not enough electrons to give the central atom an octet, try multiple
bonds.
Guide to valency in periodic groups
Chapter 8 Basics of Chemical Bonding 7
Formal Charge is a way of keeping count of electrons, based on the number of
lone pairs in the free atom and its unpaired electrons that form bonds. The sum of
formal charges equals zero in a neutral molecule or the charge of an ion.
Formal Charge = (group number) – (non bonding e’s) – 1/2(shared e’s)
Chapter 8 Basics of Chemical Bonding 8
Chapter 8 Basics of Chemical Bonding 9
Resonance Structures
+1
+1
-1
1.278 Å
O
O
O
O
O
O
O
-1
1.278 Å
O
O
O
O
O
O
O
O
Resonance Structure Rules:
1) All resonance structures must be valid Lewis structures (octet CNOF)
2) Only the placement of electrons can be changed (atoms cannot be moved,
especially H’s)
3) The number of unpaired (not lone pairs) electrons must stay the same
4) The major resonance contributor is the one of lowest energy
5) Resonance stabilization is best when delocalising a charge over 2 or more
atoms.
General points:
Normally lone pairs or multiple bond electrons are the most common to move for
resonance structures.
Good contributors will have all octets satisfied, as many bonds as possible, and as
little charge separation as possible.
Negative charges are most stable on electronegative elements.
Chapter 8 Basics of Chemical Bonding 10
Exceptions to the Octet Rule
1) Molecules with an odd number of electrons (broken bonds) radicals;
2) Molecules in which an atom has less than an octet (B and Al);
3) Molecules in which an atom has more than an octet (P,S,Cl).
Cl
Cl
P Cl
Cl
Cl
3s
3p
3d
Chapter 8 Basics of Chemical Bonding 11
3s23p3 in atom,
3s13p33d1 in molecule PCl5
S atom in sulfate ion, SO42-
Must Know!!!
Sulfate, sulfite
Nitrate, nitrite
Phosphate
Perchlorate, chlorate, chlorite, hypochlorite
Chapter 8 Basics of Chemical Bonding 12
Strengths of Covalent Bonds
Bond enthalpy is the enthalpy change, ∆H , for breaking a particular bond in a
mole of gaseous substance.
The bond enthalpy is always a positive quantity.
Bonding Polarity and Electronegativity
nonpolar covalent bond - electrons are shared equally between atoms.
polar covalent bond – one atom attracts electrons more strongly than the other.
electronegativity – the ability of an atom in a molecule to attract electrons to
itself.
Pauling scale,
F 4.0(highest)
Cs 0.7(lowest
H
2.1
Li
Be
B
C
N
O
F
1.0
1.6
1.8
2.5
3.2
3.4
4.0
Na
Mg
Al
Si
P
S
Cl
0.9
1.3
1.6
1.9
2.2
2.5
3.2
FONClBrICSH
Notice: EN increases left to right, and down to up.
Chapter 8 Basics of Chemical Bonding 13
Bond polarity
Compound
F2
HF
LiF
EN difference
4.0 – 4.0 = 0
4.0 – 2.1 = 1.9
4.0 – 1.0 = 3.0
Type of Bond
Nonpolar
Polar Covalent
Ionic
covalent
In general:
EN difference < 0.5
Nonpolar
0.5 < EN difference < 2.0
Polar
EN difference > 2.0
Ionic
Dipole moments
H F
 
H F
H F
Na+ F H
H F
H F
F H
Cl–
F
H
H
F
H F
Bond Enthalpy and Bond Length
C–C
C=C
CC
1.54 Å
1.34 Å
1.20 Å
348 kJ/mol
614 kJ/mol
839.kJ/mol
Chapter 8 Basics of Chemical Bonding 14
As the number of bonds between two atoms increases, the bond grows shorter and
stronger.
N–N
N=N
NN
1.47 Å
1.24 Å
1.10 Å
163 kJ/mol
418 kJ/mol
941 kJ/mol
Chapter 8 Basics of Chemical Bonding 15