CHEMICAL BONDING
Chemical Bonds, Lewis Symbols and the Octet Rule
 Properties of materials can be understood in terms of their
microscopic properties. Theses properties include (1) the
connectivity between atoms and (2) the 3 dimensional shape
of the molecule.
 In chemical bonds, electrons are shared or transferred
 Types of bonds:
o Ionic ( electrostatic forces hold ions together, NaCl)
o Covalent(electron sharing, CH4
o Metallic( nuclei surrounded by electrons, Na)
Lewis Symbols
 Electrons involved in bonding are called valence electrons
 Valence electrons are found in incomplete, outermost shell
of the atom
 Valence electrons are represented by unpaired dots on the 4
sides of a square around the elements symbol
 Atoms tend to gain , lose, or share electrons until they are
surrounded by 8 valence electrons. This is the octet rule. An
octet consists of full s and p subshells
 S2P6 is a noble gas configuration. Very stable
Ionic Bonding
 Na(s) + ½ Cl2 NaCl(s) Hf= -410.9 kJ/mol
 The reaction is violently exothermic. We infer NaCl is more
stable than its constituent elements.
 Why? Na loses an electron and Cl gains an electron. Both
achieve a noble gas configuration.
Energetics of Ionic Bonds
 The heat of formation of NaCl is exothermic
 Separation of NaCl to sodium and chloride ions is
endothermic: NaCl Na+ + Cl- H =+788kJ/mol
 This is called the lattice energy
 The stability of an ionic compound comes from the
attraction between ions of opposite charges.
 Specifically: E = k Q1Q2/ d where Q1, Q2 are the particle
charges, d is the distance between their centers and k is a
constant.
Configuration of ions of representative elements
 Configuration can predict stable ion formation
o Mg : [Ne] 3s2
o Mg+: [Ne} 3s1 not stable
o Mg+2 : [Ne] stable
o Cl : [Ne] 3s23p5
o Cl- : [Ne] 3s23p6= [Ar] stable
Transition metal ions
 Lattice energies compensate for the loss of up to 3 electrons
 We typically encounter cations with 1+, 2+ and 3+ charges
 Most transition metals have more than 3 electrons beyond a
noble gas core. Silver has a [Kr] 4d105s1 configuration. In
forming Ag+ it loses the 5s electron, leaving a filled d shell.
Obviously, not a noble gas configuration. The octet rule
does have limitations.
Polyatomic ions
 Formed when there is an overall charge on a covalently
bonded compound
 SO4-, NO3Ion size
 Ionic size is important in determining lattice energy
 Generally, cations are smaller than the parent
o Electron removed from most spatially extended
orbital
o Effective nuclear charge increase
o Anions are larger than their parents
Isoelectronic series have the same number of electrons
 As nuclear charge increases the ion becomes smaller
 O2- F-Na+Mg+2Al+3
Covalent bonding
 A chemical bond formed by the sharing of electrons
 Both atoms acquire a noble gas configuration
Multiple bonds
 One shared pair is a single bond, 2 pair is a double bond,
3pair is a triple bond
 Bond lengths decrease from single to triple
Bond polarity and electronegativity
 Usually, shared electrons are shared unequally. This results
in partial charge son bonds
 In nonpolar covalent bonds the electrons are equally
shared(Cl2)
 In polar covalent bonds, one atom exerts a greater
attraction(HCl)
Electronegativity
 The ability of an atom to attract an electron
 Related to ionization energy and electron affinity
 Electronegativity increases from left to right across a period
and decreases down a group
 Differences in electronegativity result in bond polarity.
Dipole moment is the quantitative measure of the
magnitude of a dipole.
Drawing Lewis Structures
 Add up all valence electrons on all atoms
 For an anion, add electrons
 For a cation, subtract electrons
 Identify the central atom
 Place the central atom in the middle and the other atoms
around it. Place one bond(2 electrons) between each pair of
atoms
 Complete the octet for all atoms except the central atom
 Put remaining electrons on the central atom.
Formal charge
 It is possible to draw more than one Lewis structure with
the octet rule obeyed.
 To determine which structure is best we use formal charge
 The formal charge is the charge that an atom,in a molecule,
would have if all the atoms had the same electronegativity.
 To calculate formal charge
o Assign all unshared electrons to the atom on which
they are found
o Assign half of the bonding electrons to each atom in a
bond
o The formal charge= the number of valence electrons
minus the number of assigned electrons
 Consider CNo For carbon
 4 valence electrons
 In the Lewis structure, there are 2 nonbonding
electrons and 3 from the triple bond
 There are 5 electrons from the Lewis
 Formal charge= 4-5= -1
o For Nitrogen
 5 valence electrons
 Lewis gives 2 nonbonding and 3 from the triple
bond
 Formal charge= 5-5=0
 Using formal charge
o Preferred structure has the smallest formal charge
on each atom
o The most negative charge on the most electronegative
atoms
o Formal charges DO NOT REPRESENT REAL
CHARGES on atoms.
Resonance structures
 Resonance structures are attempts to represent a real
structure that is a mix of extremes. Ozone has 2 identical
bonds(experimentally) while the Lewis structure has a
single and a double bond
  is used to represent resonance
 O3, NO3-, SO3 , NO2 and benzene
Exceptions to the octet rule
 Molecules with an odd number of electrons: ClO2, NO , NO2
 Molecules where one atom has less than an octet:
compounds with boron or beryllium
 Molecules with more than an octet. The largest group.
Atoms from the third period onward can accommodate
more than 8 by using unfilled d subshells: PCl5
Covalent bond strength
 The energy required to break a bond is called bond
enthalpy.
 When more than one bond is broken: CH4(g) C + 4H
H= 1660 kJ/mol . The individual bond energy is ¼ H
Bond enthalpy and Enthalpy of reaction
 Hrxn =  bonds broken -  bonds formed
 CH4 + Cl2 CH3Cl + HCl : In this reaction, one C-H and
one Cl-Cl bond is broken and one H-Cl bond and one C-Cl
bond is formed. Therefore, Hrxn = [ D(C-H) + D(Cl-Cl)] –
[D(C-Cl) = D(H-Cl)]
Bond enthalpy and bond length
 Bond length is the distance between atomic nuclei.
 As the number of bonds increases, the bond grows shorter
and stronger. Note that the enthalpy of a double bond is not
2x single bond