Chapter 9:
 Basic Concepts of
Chemical Bonding
 NaCl versus C12H22O11.
Types of Bonds
 We can classify bonds based on the kinds of atoms that are
bonded together.
Types of Atoms
metals to
nonmetals
nonmetals to
nonmetals
metals to
metals
Tro: Chemistry: A Molecular Approach, 2/e
Type of Bond
Ionic
Covalent
Metallic
Bond
Characteristic
electrons
transferred
electrons
shared
electrons
pooled
2
Types of Bonding
3
Tro: Chemistry: A Molecular Approach,
2/e
Lewis Dot Symbols
 Combines the element symbol plus the valence
electrons as dots placed around symbol.
 Dots are first placed on each of the four sides (N-S-EW), then paired up after that.
 Symbols for period 2.
Octet Rule
 Atoms tend to gain, lose, or share electrons until they
are surrounded by eight electrons.
 Ionic = gain or lose
 Covalent = share
 While there are exceptions to this rule, it is the
important driving force for the formation of
compounds.
Ionic Bonding
 When Na(s) and Cl2(g)
are combined, a very
violent, exothermic
reaction results.
 LEP #1
Ionic Bonding
Ionic Bonding
 The energetics of ionic bond formation can be
explained by a series of steps.
 Step 1: Loss of electron by Na
Na(s)  Na(g) ;
Na(g)  Na+(g) + 1e- ;
DH = +108 kJ
DH = +496 kJ
 Step 2: Gain of electron by Cl
½ Cl2(g)  Cl(g) ;
Cl(g) + 1e-  Cl-(g) ;
DH = +122 kJ
DH = -349 kJ
Ionic Bonding
 Lattice energy is the energy required to completely
separate a mole of a solid ionic compound into gaseous
ions.
NaCl(s)  Na+(g) + Cl-(g) ; DH = 788 kJ
 Reverse this AND add it to the four previous
reactions yields:
Na(s) + ½ Cl2(g)  NaCl(s) ; DH = -411 kJ
Lattice Energy
 The electrostatic attraction of
two charged particles is ruled
by the equation:
E
Q1Q 2
d
 Q1, Q2 are the magnitudes of
the charges and d is the
distance between the two
nuclei.
Lattice Energy
 As the magnitude of the charge increases, the lattice
energy will increase.
 Na+1 – Cl-1 U = +788 kJ
 Sr+2 – Cl-1 U = +2127 kJ
 Sr+2 – O-2 U = +3217 kJ
Lattice Energy
 As the distance between the ions
increases, the lattice energy
decreases.
Summary of Lattice Energies
 The force of attraction between
oppositely charged particles is directly
proportional to the product of the
charges
 Larger charge means the ions are more
strongly attracted
 larger charge = stronger attraction
Lattice Energy =
−910 kJ/mol
 stronger attraction = larger lattice energy
 Of the two factors, ion charge is generally
more important
 LEP #2
13
Tro: Chemistry: A Molecular Approach,
2/e
Lattice Energy =
−3414 kJ/mol
Covalent Bonding
 When two electrons are shared by two atoms, this
is known as a covalent bond.
 Formation of H2
Covalent Bonding
 Formation of Cl2 – each Cl atom has seven valence
electrons. Each has one unpaired electron that can
pair up to make a bond.
Lewis Structures

1.
2.
For molecules or ions containing three or more
atoms, we can follow a set of rules to guide in the
process.
Add up the total valence electrons from all the
atoms in the compound. Ex) CF4
Make a skeleton structure – the first element in
the formula is usually the central atom – all
others are then placed around this atom and
connected with a bond.
Lewis Structures
Fill the external atoms until they have an octet.
4. Compare total number of electrons used to step
1. If all are used, then go to step 5. If some are
left, place on central atom as lone pair(s).
5. Check the central atom for an octet. If no octet,
may need multiple bond(s) by moving nonbonding pair(s).
3.
Lewis Structures
 General Guidelines
 Group 7A as an external atom will NOT do multiple
bonds.
 Carbon will almost always have four bonds.
 Group 2A, 3A as a central atom may be deficient of octet.
 Group 5A, 6A, 7A, and Xe as a central atom may exceed
the octet.
 Oxygen will do up to two bonds.
Electronegativity
 Non-polar covalent bond – the electrons are shared
equally.
 Ex) F2 , Br2 , I2 , O2 , etc.
 Polar covalent bond – electrons are not shared equally
– one atom has a greater desire for the electron pair.
 Electronegativity Scale
Electronegativity
Bond Polarity
 If the electronegativity difference is zero, then the
bond is non-polar covalent.
 If there is a difference AND the two elements are
non-metals, then the bond is polar covalent.
 If there is a difference AND one element is a metal
and the other is a non-metal, then the bond is
ionic.
Bond Polarity
 Simple molecules like HCl have a polar covalent bond.
 The more electronegative element will have a partial
negative charge and the less electronegative element
will have a partial positive charge.
Dipole
Moment
 This is the quantitative measurement of the polar bond.
 m=Qxd
 m is measured in a unit called the Debye or Coulomb x
meter.
 LEP #4
Resonance Structures
 In some cases, a molecule or ion may be described by
more than one Lewis Structure.
 Ex) O3
 LEP #5
Bond Strength
 The energy required to break a covalent bond is its





strength.
Some are fairly simple.
Cl2(g)  2 Cl(g) ; DH = 242 kJ
Others are more complicated.
CH4(g)  C(g) + 4 H(g) ; DH = 1660 kJ
Note: Energy is ALWAYS required to break a bond.
Using Bond Enthalpies
 One method for estimating the enthalpy of a reaction
is:
DH = S(Bonds Broken) – S(Bonds Made)
 LEP #6
Chemistry of Explosives
 Many explosives are solids or liquids that contain the
nitro (NO2) or nitrate (NO3) group.
 The nitrogen atoms are typically weak (singly) bonded
to carbon (293 kJ).
 They then become N-N strong triple bonds (941 kJ).
Chemistry of Explosives
Bond Order
 Bond order is the number of bonds connecting two
atoms.
 Can be 1, 2, or 3.
C2H2
Bond Length
 As bond order increases, the bond lengths decrease.
 Ex) N-N N=N
1.47Å 1.24Å
 Ex) C-C C=C
1.54Å 1.34Å
NN
1.10Å
CC
1.20Å
Resonance Structures
 The bond orders and lengths for resonance structures
must be averaged.
 Ex) SO2 , NO3-1 , C6H6
Formal Charges
 Formal Charge is a fictitious charge assigned to each
atom in a Lewis Structure.
 It helps to evaluate the validity of competing
structures.
 Formal Charge = #Valence Electrons – #Nonbonding
electrons – ½ #Bonding Electrons
Formal Charges
 General Rules:
The sum of all formal charges in a neutral molecule
will equal zero.
2. The sum of all formal charges in an ion must equal
the charge of the ion.
3. Small (+1 or -1 or 0) formal charges are preferred over
larger ones.
4. When formal charges cannot be avoided, a negative
formal charge will preferentially go on the more
electronegative atom.
1.
Formal Charges
 HCN molecule = 10 electrons.
 Is skeleton structure H – C – N or H – N – C?
 Which Lewis Structure for the cyanate (OCN-) ion is
best?