the chemical bond:
Chapter 9
Ionic and Covalent
Bonding
◆
the force that holds atoms or ions
together as an aggregate unit
bond energy (or bond dissociation enthalpy, ∆HBDE):
◆
◆
energy required to break a chemical bond
Cl–Cl (g) ! 2 Cl (g); ∆H = bond
energy
bond energy is always endothermic
generally 3 types of bonds - ionic, covalent, metallic
Lewis Symbols
developed by G.N. Lewis to represent an element and
its number of valence electrons
◆
◆
◆
each side of element’s symbol may have
0, 1, or 2 dots
each dot represents a valence electron
for main group elements:
# valence e–’s = group #
The Octet Rule
atoms tend to lose, gain, or share e-’s in such a way
that they attain a noble gas configuration (ns2 np6)
◆
ionic compounds:
typically metals with low ionization energy
lose e-‘s
nonmetals with favorable electron affinity
gain e-’s
◆
The Octet Rule
molecular compounds:
typically nonmetals will share e-’s to form
covalent bonds
Energy Considerations of Ionic Compounds
Lattice Energy, U: the energy required to
separate an ionic solid into its gas phase
ions
MX (s) ! M+ (g) + X– (g);
Lattice Energy and Coulomb’s Law
Coulomb’s Law:
◆ describes the energy of electrostatic
interaction between 2 charged particles
separated by some distance
Q1●Q2
Ecoul ∝ ––––––
r
Q1 and Q2 = charges on ions
r = distance between ions
Electron Configurations of Ions
cations
◆
◆
this relationship is also true for lattice energy
U ∝ charges on ions
U ∝ 1/distance separating ions
formed by the loss of electrons
e–’s are lost from the highest energy,
populated atomic orbital
anions
◆
◆
◆
endothermic
formed by the gain of electrons
e–’s are added to the lowest energy available
atomic orbital
How does the e– configuration of an atom change as an ion is formed?
Electron Configurations of Ions
Ionic Radii
some examples of cations that do not have a noble
gas configuration:
◆
◆
◆
size of cation or anion relative to the parent atom
cations are smaller than
their parent, neutral atom
heavy main group metals:
ex. Sn2+
rNa+ < rNa
anions are larger than their
parent, neutral atom
transition metals:
ex. Fe3+
rCl– > rCl
Periodic Trend in Ionic Radii
◆
◆
Similar to atomic radii, ionic radii increase from
the top to the bottom of the periodic table.
To understand the trend moving left to right across
the periodic table, consider isoelectric species.
isoelectric – same number of electrons
O2–
F–
Na+
Mg2+
Al3+
# e–’s
10
10
10
10
10
# p’s
8
9
11
12
13
ionic
radius
140 pm
136 pm
95 pm
65 pm
50 pm
Molecular Compounds:
Covalent Bonds
In molecular compounds, atoms share electrons to
achieve a noble gas configuration
◆
◆
◆
covalent bonds result from shared pairs of
electrons between atoms
electrons are shared in the region between
atoms where atomic orbitals overlap
orbital overlap results in a concentration of
electron density between 2 nuclei
Nonpolar vs Polar Covalent Bonds:
How is the electron density distributed between
atoms in a bond?
Electronegativity, !:
the ability of an atom in a polyatomic species to
attract electrons to itself
Some Introductory Thoughts About Structure and
Bonding in Molecular Compounds
◆
How many covalent bonds is a central atom in a
molecule likely to form?
consider the element’s group # and # valence e–’s
use the octet rule as a guideline
◆
◆
bonding vs. nonbonding pairs of electrons
The greater the difference in electronegativity (∆ !)
between 2 atoms in a bond, the more polar the bond.
Some Introductory Thoughts About Structure and
Bonding in Molecular Compounds
◆
Drawing Lewis Structures
single vs. multiple bonds
1 pair of e–’s shared between atoms (i.e. 2 shared e–’s)
! single bond
2 pairs of e–’s shared between atoms (i.e. 4 shared e–’s)
! double bond
3 pairs of e–’s shared between atoms (i.e. 6 shared e–’s)
! triple bond
a 1st step to understanding:
◆
◆
◆
◆
◆
◆
◆
◆
molecular structure
atom connectivity
arrangement & distribution of valence e–’s
numbers and types of bonds
numbers of nonbonding pairs of e–’s
3-D shape
bond angles
molecular polarity
Drawing Lewis Structures
start by thinking about structures of molecules
using the localized electron model:
◆
◆
localized electron model assumes molecules
are collections of atoms bonded together by
covalent bonds
pairs of electrons are either localized on
atoms (lone or nonbonding e– pairs), or
localized in the space between 2 atoms
(bonding e– pairs)
Drawing Lewis Structures
We will use the localized electron model to:
◆
describe the atom arrangement and
distribution of valence e–’s in a molecule
Lewis Dot Structures (now)
◆
predict molecular geometry, bond angles,
and polarity
VSEPR Theory (Ch. 10)
◆
describe the types of atomic orbitals used
by atoms in bonding or to house lone pairs
Valence Bond Theory (Ch. 10)
Drawing Lewis Structures
1. Determine the total number of valence electrons in
the molecule:
total # valence e–’s = ∑valence e–’s of atoms in
molecule
◆
◆
◆
main group elements: # valence e–’s = group #
add 1e– for each unit negative charge on anion
subtract 1e– for each positive charge on cation
! this is the total number of electrons you will
need to have in your final structure
2. Write symbols for atoms in order of connectivity;
connect appropriate atoms with single bonds
Drawing Lewis Structures
3. Complete the octets of atoms bonded to the
central atom(s).
H, He, Li, and Be will only have a duplet of e–’s
4. Place any left over electrons on the central
atom, even if it results in greater than an octet.
5. If there are not enough electrons to give the
central atom a full octet, try multiple bonds.
Drawing Lewis Structures
examples:
PCl3
What if you can draw more than one Lewis
Structure that obeys the octet rule?
CH2Cl2
Which one is the “right” one?
2 concepts to help interpret Lewis structures:
HCN
ClO2–
NH4+
N2H4
◆
◆
COBr2
◆
Considering formal charges
to determine formal charges of elements in structure:
1. all unshared e–’s (lone pairs) are assigned to the atom
on which they are found
e–’s
2. bonding
are assumed to be shared evenly between
the atoms participating in the bond;
◆
homolytic clevage of the bond
◆
! of the bonding e–’s are assigned to each atom in the
bond
3. formal charge = # valence e–’s of isolated atom –
# e–‘s assigned by Lewis structure
Formal Charge
same atom arrangement; different
e– arrangement
Resonance Structures
same atom arrangement; same net e–
arrangement; same formal charge
distribution
In general, the more stable Lewis structure is
considered to be the one in which:
◆
◆
atoms bear formal charges closest to 0
any negative formal charge resides on more
electronegative element
ex: Consider 2 possible Lewis structures of CO2.
ex: Consider 3 possible Lewis structure of NCS–.
◆
Considering resonance structures
sometimes one Lewis structure does not
adequately describe e– arrangement
The 2 resonance structures A and B are equivalent
contributors to the overall resonance hybrid structure.
supported by experimental evidence
resonance hybrid
structure
example: Consider 2 possible Lewis structures for
ozone, O3:
example: carbonate ion, CO32–
experimental data:
◆ O3 is a bent molecule
◆ both O–O bond lengths equivalent
Exceptions to the Octet Rule
1. odd number of electrons
ex. NO
Exceptions to the Octet Rule
3. central atom has more than an octet of e–’s
◆
◆
2. central atom has less than an octet
ex. PCl5
ex. BF3
XeF4
expanded valence
possible for larger central atoms in the
3rd period and below
A Way to Think About Expanded Valence
for PCl5:
◆ central atom: P
◆ valence e– configuration: 3s2 3p3
◆ empty 3d orbitals sit at slightly higher energy
A Way to Think About Expanded Valence
for XeF4:
◆ central atom: Xe
◆ valence e– configuration: 5s2 5p6
◆ empty 5d orbitals sit at slightly higher energy
◆
◆
a 3s e– absorbs E and is promoted to a
higher E, empty 3d orbital
↑↓
↑
3s
↑
↑
↑
↑↓
3d
3p
↑
3s
↑
↑
3p
↑
↑↓
3d
5s
Using Covalent Radii to Approximate Bond Length
A
A
A–A bond length = rA + rA
◆
A
↑↓
5s
B
A–B bond length = rA + rB
knowing periodic trend in atomic radii can help
you make predictions about relative lengths of
bonds
example: Would you predict that a N–Cl or
a P–Br bond would be longer?
2 5p e–’s absorb E and are promoted
to higher E, empty 5d orbitals
↑↓
↑↓
5p
↑↓
↑
5p
5d
↑
↑
↑
5d
Bond Multiplicity (or Bond Order) and
Relationship to Bond Length and Bond Energy
bond
length
decreases
A Word About Tabulated Bond Energy Data
bond
order
# e– pairs
shared
# e–’s
shared
single
1
1
2
double
2
2
4
triple
3
3
6
bond
energy
increases
Using Bond Energies to Approximate ∆Hrxn:
∆Hrxn = ∑ E of bonds broken " ∑ E of bonds formed
recall:
◆
bond
type
bond energy (or bond enthalpy) is the energy
required to break a bond
◆
endothermic
◆
units kJ/mol
example:
What is the C–H bond energy?
if: CH4 (g) ! C (g) + 4 H (g); ∆H = 1660 kJ
then: we can approximate the average C–H
bond energy as 1660 ÷ 4 = 415 kJ
Using Bond Energies to Approximate ∆Hrxn:
∆Hrxn = ∑ E of bonds broken " ∑ E of bonds formed
example:
Calculate ∆H for the following reaction using
bond energies:
C2H4 (g) + H2O (l) ! C2H5OH (l)