Review Chapter 9: The Basics of Chemical Bonding Chemistry: The
Molecular Nature of Matter, 6 th edition By Jesperson, Brady, &
Hyslop
Slide 2
Chapter 9 Concepts 2 Know the difference between an ionic and
covalent bond Electronegativity: Identify polar bonds and estimate
relative dipole moments Ionic Bond: Energetics of ionic bonding:
lattice energy Predict possible ionic compounds For transitions
metals predict ions that will have half, full, or empty d-orbitals
Covalent Bond: Draw lewis dot structures Predict most reasonable
lewis dot structure Formal charges Electronegativity Identify
possible reasonable resonance structures Draw hybrid resonance
structures Understand electron delocalization & stability from
it.
Slide 3
Covalent vs Ionic Bonds Ionic Bonds result from electrostatic
attraction between a cation and anion: metal-nonmetal (with the
exception of NH 4 + and H 3 O + cations). Covalent bonds result
from the sharing of electrons between two atoms: nonmetal-nonmetal.
Li F Ionic Bonds Covalent Bonds
Slide 4
Ionic Bonds 4 Coulombs law determines the potential energy of
two ions (q 1 and q 2 ) separated by a distance (r), where k is a
proportionality constant. Lattice Energy For a stable ionic
compound to form the potential energy must be lowered. Lattice
Energies are usually very large negative numbers. Lattice Energies
always exothermic Transition Metals Hard to predict ions electron
configurations, but transition metal ions with exactly filled or
half-filled d subshells are extra stable and therefore tend to
form. This is transition metals form multiple oxidation states
Slide 5
Electronegativity & Polar Bonds + HCl + on H = +0.17 on Cl
= 0.17 EN = |EN 1 EN 2 | = q r Atoms participating in covalent
bonds will not share electron density equally if they have a
difference in electronegativities of between 0.5 and 1.7.
Slide 6
Drawing Lewis Dot Structures Step [1] Determine the valence
electrons for each element and draw the lewis symbol. Step [2]
Count the valence electrons. The sum gives the total number of e
that must be used in the Lewis structure. For each atom the number
of bonds = 8 valence electrons. Step [3] Arrange the atoms next to
each other that you think are bonded together. Place H and halogens
on the periphery, since they can only form one bond. Step [4]
Arrange the electrons around the atoms. Place one bond (two e )
between every two atoms. Use all remaining electrons to fill octets
with lone pairs, beginning with atoms on the periphery.
Slide 7
Drawing Lewis Dot Structures 1.Decide how atoms are bonded
Skeletal structure = arrangement of atoms. Central atom Usually
given first Usually least electronegative 2.Count all valence
electrons (all atoms) 3.Place two electrons between each pair of
atoms Draw in single bonds 4.Complete octets of terminal atoms
(atoms attached to central atom) by adding electrons in pairs
5.Place any remaining electrons on central atom in pairs 6.If
central atom does not have octet Form double bonds If necessary,
form triple bonds
Slide 8
The Octet Rule Exceptions Holds rigorously for second row
elements like C, N, O, and F B and Be sometimes have less than
octet BeCl 2, BCl 3 2 nd row can never have more than eight
electrons 3 rd row and below, atoms often exceed octet When atoms
form covalent bonds, they tend to share sufficient electrons so as
to achieve outer shell having eight electrons
Slide 9
Formal Charges & Reasonable Structures FC = # valence e [#
bonds to atom + # unshared e ] Most Stable Lewis Structure 1.Lowest
possible formal charges are best 2.All FC 1 3.Any negative FC on
most electronegative element +2 0 0 0 0 0 0 0
Slide 10
Resonance Stabilization = Resonance Structures Resonance Hybrid
Drawing Good Resonance Structures: 1.All must be valid Lewis
structures 2.Only electrons are shifted - Usually double or triple
bond and lone pair - Nuclei can't be moved - Bond angles must
remain the same Resonance Structures are equivalent lewis dot
structures: - Atom connectivity does not change (ie, nuclei stay
the same) - Only electrons change, ie, # of bonds between atoms can
rearrange Hybrid Resonance Structures show how electrons are
delocalized over the molecule, the more atoms that are
participating in the delocalization the more stable the molecule
is. 4.Number of unpaired electrons, if any, must remain the same
5.Major contributors are the ones with lowest potential energy (see
above) 6.Resonance stabilization is most important when
delocalizing charge onto two or more atoms
Slide 11
Coordinate Covalent Bonds
Slide 12
Problem Set A 1.Identify the covalent and ionic bonds: a.CaF 2
b.CCl 4 c.NaOH d.NH 4 NO 3 2.Which ionic solid is likely to have
the smallest exothermic lattice energy? LiCl; CsCl; NaCl, KCl
3.What ion will form for each element? Draw the elements lewis
symbol and write the ions electron configuration. a.I b.Rb c.P
4.Predict the ionic compound that will form between the following:
a. Aluminum (Al) and Chlorine (Cl) b.Strontium (Sr) and Bromine
(Br)
Slide 13
Problem Set B/C 5.Draw lewis dot structures for the following
molecules: a.C 6 H 6 b.C 5 H 10 c.C 2 HCl 3 O 2 d.C 2 H 7 N e.CH 2
O 3 f.NH 4 CN 6.Which has the least polar bond? HCl, HF, HI, HBr.
7.What are 2 possible structures for XeO4? Please indicate the
formal charge on Xe in both structures and decide which is the most
reasonable. 8.Draw 3 resonance structures of NCO -. Evaluate the
formal charges on each and decide which is the best structure.
9.Draw the best lewis structure for the following: a.HClO 4 b.XeF 4
c.I 3 d.BrF 5
Slide 14
Solutions 1.Identify the covalent and ionic bonds: a.CaF 2
Ionic b.CCl 4 Covalent c.NaOH Ionic, OH - a polyatomic ion with
covalent bonds d.NH 4 NO 3 Ionic, NH 4 + and NO 3 - are polyatomic
ions with covalent bonds 2.CsCl, because Cs has the largest radius,
therefore r large and E less negative. 3.What ion will form for
each element? Write the ions electron configuration and the lewis
symbol. a.[Kr] 5s 2 4d 10 5p 6 b.[Kr] or [Ar] 4s 2 3d 10 4p 6
c.[Ne] 3s 2 3p 6 4.Predict the ionic compound that will form:
a.AlCl 3 b.SrBr 2 I Rb
