Transcript Chemical Bonding Basic Concepts

Chemistry 100 Chapter 8
Chemical Bonding Basic Concepts
The Valance Electrons
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When atoms interact to form chemical bonds, only the
outer (valance) electrons take part.
Need a tool for keeping track of valence electrons, e.g.,
The Lewis dot symbol
Na .

+
..
:Cl :
.
NaCl
1 v.E.
7 v.E’s
When these two elements combine to form a compound
2 Na (s) + Cl2 (g)  2 NaCl (s)
What’s Happening?
..
Na . 1+ :Cl : 2 5 NaCl
.
[Ne]3s [Ne]3s
3p
(g)  Na+ (g) + e- (ionizes, loses e-)
.
Na

an electron configuration of [Ne]
.. + e-  Cl- (g)
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(g)
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: Cl :
.
an electron
configuration of [Ar]
In the crystal lattice,
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Na+ and Cl- ions; strong electrostatic attractions
The NaCl Crystal
Ionic Bonding
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Electrostatic attractions that hold ions
together in an ionic compound.
The strength of interaction depends on
charge magnitude and distance between
them.
Eionic
kq1q2

r
q1  magnitude of charge 1
q2  magnitude of charge 2
r  distance between the ionic
centres
Stability of Ionic Compounds
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The stability of ionic compounds depends on two main
factors
1.
2.
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Note
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The electron affinity of one of the elements
The ionization energy of the other
electron affinities and ionization potentials are gas-phase
reactions?
How are they related to the stability of solid materials?
The Lattice Energy
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A quantitative measure of just how strong the interaction
is between the ionic centres (i.e., a measure of the
strength of the ionic bond)
For the reaction
KCl (s)  K+ (g) + Cl- (g) H = 718 kJ/mol
Lattice energy (latH).

The energy required to completely separate one mole of the solid
ionic compound into its gas-phase ions.
Lattice Energies of Various Ionic
Compounds
Determined using a thermochemical cycle the Born-Haber cycle (a Hess’s Law application)
Covalent Bonding
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In a wide variety of molecules, the bonding atoms fulfill
their valance shell requirements by sharing electrons
between them.
Covalent bonds - a bond in which the electrons are
shared by two atoms.
H2  H-H, F2  F-F, Cl2  Cl-Cl
For many electron atoms (like F and Cl), we again to
worry only about the outermost (valence) electrons.
Covalent Bonding
Examples of Covalent Bonding
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Let’s look at the Cl2 example.
Each Cl atom has 7 valence shell electrons 3
Lone pairs and one unpaired electron
Lone pairs
..
: Cl :
.
Unshared electron
The Cl2 Molecule
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lone pairs (non bonding)
bonding
electrons
The structure we have just drawn are
called
..
..
Lewis structures. :Cl
Cl
.. :
..
The dash in between the atomic centres
represents the bonding electrons
Redraw F2
..
:F
..
..
.F. :
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Note both Cl2 and F2 satisfy their valence shell
requirements by the formation of a single bond.
What about O2? How can we satisfy the octet rule for 2
O atoms?
..
O
..
..
O
..
Valence shell requirements are satisfied by the
formation of a double bond.
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check out N2  :NN: (triple bond)
Note that the octet rule works mainly for the
second row elements.
Filled valence shells can have more than 8
electrons after Z=14 (Si). This is generally
termed octet expansion.
Covalent Compounds
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Compounds that contain only covalent bonds
are called covalent compounds.
There are two main of covalent compounds,
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Molecular covalent compounds (CO2, C2H4)
Network covalent compounds (SiO2, BeCl2).
The network covalent compound are
characterized by an extensive “3-D” network
bonding
Comparison between Ionic and Covalent
Compounds
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Ionic Compounds
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usually solids with very high melting points
 Covalent Compounds
conduct electricity when molten (melted)  usually low melting solids,
gases or
usually quite water soluble and they are electrolytes
in liquids
aqueous solution
 don’t conduct electricity when
NaCl
molten
 aren’t very soluble in water and
are non electrolytes
 CCl4
The Filled Valence Shell rule
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Filled Valence Shell rule
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Atoms participate in the formation of bonds (either ionic or
covalent) in order to satisfy their valence shell requirements.
Atoms other than H tend to form bonds until they end up
being surrounded by 8 valence electrons (the noble gas
configuration).
This is known as the octet rule.
Electronegativity
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Electronegativity is defined as the ability of an atom to
attract electrons towards itself in a molecule (
(pronounced ‘chi’))
Examine the H-F covalent bond
+H-F
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 denotes a partial “+” charge on the H atom
- denotes a partial “-“ charge on F atom
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Electronegativity is related to the electron
affinity and the ionization energy.
Compare the following elements.
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Na  low I1, small negative E.A.  low 
F  high I1, large, negative E.A.,  high 
Trends in the  Values
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Across a row
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Down a group
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The  values generally increase as we proceed
from left to right in the periodic table.
The  values generally decrease as we descend
the group.
Transition metals
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Essentially constant  values
Plot of  Values
Electronegativity and Bond Type
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Can we use the electronegativity values to help us
deduce the type of bonding in compounds?
 values
bond type
0.0 <  < 0.5
non-polar covalent
0.5    1.9
polar covalent
2.0    3.3
Ionic bond
An Outline for Drawing Lewis Structures
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Predict arrangement of atoms (i.e., predict
the skeletal arrangement of the molecule or
ion).
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The H is always a terminal atom, bonded to ONE
OTHER ATOM ONLY. A halogen atom is usually
a terminal atom.
Note that the central atom usually has the least
negative electron affinity.
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Count total number of valence shell electrons (include
ionic charges).
Place 1 pair electrons (sigma bond, ) between each
pair of bonded atoms (i.e., the central atom and each
one of the terminal atoms).
Place remaining electrons around the terminal atoms to
satisfy the filled valence shell rule. (lone pairs).
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All remaining electrons are assigned to the
central atom. Atoms in the 3rd or higher row
can have more than eight electrons around
them.
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If a central atom does not have a filled valence
shell, use a lone pair of electrons from a terminal
atom to make a pi () bond.
Formal Charges
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Definition: formal charge on atom = number of valence
electrons – number of non-bonding - ½ the number of
bonding electrons.
Formal charge in a Lewis Structure is a bookkeeping
“device”
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keeps track of the electrons “associated” with certain atoms in
the molecule vs. the valence e-‘s in the isolated atom!
How does it work?
Rules for Formal Charges
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Neutral molecules  S formal charges = 0
Ions S formal charges = charge of ion
For molecules where the possibility of multiple Lewis
Structures with different formal charges exist
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Neutral molecule - choose the structure with the fewest formal
charges.
Structures with large formal charges are less likely than ones
with small formal charges
Two Lewis Structures with similar formal charge distribution 
negative formal charges on more electronegative atom
Resonance Structures
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Examine the NO3- anion.
O
O
O N O
O N O
• The structures differ in the location of the N=O
double bond.
• They are said to be resonance structures.
• The actual structure of the molecule is a combination
of three resonance structures (the resonance hybrid).
Experimental Evidence for Resonance.
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The resonance structures for benzene C6H6
• We would expect to find two different bond
lengths in benzene (C=C and C-C bonds).
• C= C  bond length = 133 pm = 0.133 nm
• C- C  bond length = 0.154 nm
• Experimentally, all benzene carbon-carbon
bond lengths are equivalent at 0.140 nm
Exceptions to the Filled Valence Shell
Rule
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Be compounds  BeH2, BeCl2,
Boron and Al compounds  BF3, AlCl3, BCl3
BF3 is stable  The B central atom has a tendency to pick up an
unshared e- pair from another compound
BF3 + NH3  BF3NH3
the B-N bond is an example of a coordinate covalent bond, or a
“dative” bond  i.e. a bond in which one of the atoms donates both
bonding electrons.
Odd e- molecules
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These molecules have uneven numbers of electrons \
no way that they can form octets.
Examples
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NO and NO2. These species have an odd number of
electrons.
. ..
:N O
..
.. . ..
:O
.. N O
..
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Look at the dimerization reaction of NO2.
2 NO2 (g) ⇄ N2O4 (g) Keq = 210
:O :
N
:O
.. :
..
:O :
N
:O :
Valence Shells having more than 8 Electrons
(Expanded Octets)
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A central atom having more than 8 valance shell
electrons is possible with atomic number 14 and above.
Cl
Cl
Cl P Cl
Cl
Reason - elements in this category can use
the energetically low-lying d orbitals to
accommodate extra electrons
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Look at HClO3
..
:O :
..
..
:O
.. Cl
.. H
.. O
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High formal charge on the electronegative Cl
atom (f.c.(Cl) = 7-2-1/2 (6) = +2)
This resonance structure would make a very
small contribution to the overall resonance
hybrid.
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With the possibility of using the low lying d-orbitals on
the Cl atom to accommodate extra electron pairs, we
may write other Lewis structures
..
:O :
..
..
:O.. Cl
.. O.. H
..
:O :
:O :
..
..
O.. Cl
.. O.. H
..
..
: O Cl O H
.. .. ..
:O :
..
..
O.. Cl
.. O.. H

Note: the final three structures reduce the
formal charges
Bond Energies and Thermochemistry
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Look at the energy required to break 1 mole of gaseous diatomic
molecules into their constituent gaseous atoms.
H2 (g)  H (g) + H (g)
H° = 436.4 kJ
Cl2 (g)  Cl (g) + Cl (g)
H° = 242 kJ
These enthalpy changes are called bond dissociation energies. In
the above examples, the enthalpy changes are designated D (H-H)
and D (Cl-Cl).
For Polyatomic Molecules.
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CO2 (g)  C (g) + 2 O (g)
H = 745 kJ
Denote the H of this reaction D(C=O)
What about dissociating methane into C + 4 H’s?
CH4 (g)  C(g) + 4 H (g) H° = 1650 kJ
Note 4 C-H bonds in CH4 \ D (C-H) = 412 kJ/mol
H2O (g)  2 H (g) + O (g)
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H° = 929 kJ/mol H2O
It takes more energy to break the first O-H bond.
H2O (g)  H (g) + OH (g)
H° = 502 kJ/mol H2O
HO (g)  H (g) + O (g)H = 427 kJ/mol H2O
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Note: we realize that all chemical reactions involve the
breaking and reforming of chemical bonds.
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Break bonds  add energy.
Make bonds  energy is released.
rxnH°  S D(bonds broken) - S D(bonds formed)
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These are close but not quite exact. Why?
The bond energies we use are averaged bond energies,
i.e.,
This is a good approximate for equations involving
diatomic species.
We can only use the above procedure for GAS PHASE
REACTIONS ONLY.