Transcript Lecture 12

Chapter 8
Basic concepts in Chemical
Bonding
Lecture 26. Basic Concepts in Chemical
Bonding.
The properties of substances are determined largely
by the chemical bonds that hold them together. The
two extreme types of chemical bond are the ionic
bond, such as holds the Na+ and Cl- ions together in
NaCl, or the covalent bonds that hold molecules
together such as glucose.
red =
Cl- ion
cyan =
C atom
O atoms
H-atom
Na+ ion
Sodium chloride – ionic substance
no actual NaCl molecules
glucose molecule – molecule
persists even in solution
Lewis dot symbols:
Lewis Symbols.
Lewis suggested Lewis
dot diagrams as a way
of tracking the electrons
involved in bond formation.
These are only the valence
electrons, as the core
electrons do not participate
directly in chemical
bonding.
Gilbert Lewis (1875-1946)
Lewis dot structures for elements:
The Lewis symbol for each element consists of the symbol for
the element surrounded by dots for each valence electron:
.
.
.
.
H
Li
Be
B
.
.
.
C
N
O
F
Lewis Dot Structures
Lewis dot structures present a simple approach to bonding that
allows us to rationalize much molecular structure. The idea is that
atoms share electrons in the valence shell to form the chemical
bond, with one pair of electrons per bond. Note that each H-atom
now has two electrons, which is the structure of He, the next inert
gas.
Electron pair = single bond
Valence electrons
H-atom
H-atom
(Each H-atom has one valence electron)
H2 molecule
Lewis Dot Structures (contd.):
Two shared pairs of electrons
= double bond
O-atom
O-atom
O2 molecule
Periodic table
1
2
3
Oxygen has six valence electrons
4
5
6
7
8
The octet rule
Electrons are shared in forming bonds such that atoms have the
same number of electrons in their valence shells as the nearest
noble gas, including the electrons shared with the atom to which
they are bonded.
O-atom
O-atom
O2 molecule
Each oxygen atom in the O2 molecule now has eight
valence electrons, including those it shares with the
other oxygen atom = number of electrons (8 = octet)
in the nearest inert gas = neon.
Some more examples of Lewis dot structures:
The N2 molecule:
N-atom
N-atom
triple bond
N2 molecule
Periodic table
1
2
3
4
5
6
7
8
8.1 Chemical Bonds, Lewis Symbols, and
the Octet rule.
Chemical bonding involves mainly the attempt to
achieve the rare gas number of valence electrons,
i.e. an octet. This can be achieved in several
ways.
Ionic bond: Electrons are mainly the property of
one of the two atoms forming the bond.
Covalent bond: Electrons are shared so that each
atom has a noble gas electronic configuration.
Metallic bonds. Electrons are lost into the
conduction band.
8.2 Ionic Bonding.
This occurs between metallic elements from
the left-hand side of the periodic table and
non-metallic elements from the right hand
side of the periodic table.
Note that Na gives up its lone valence
electron to Cl, so that they both end up with
an octet of electrons.
Energetics of Ionic Bond Formation:
We have seen that Na has a low ionization energy, and Cl
has a high electron affinity. In fact the process:
Na(g) + Cl(g) = Na+(g) + Cl-(g)
is endothermic = 496 – 349 = + 147 kJ/mole.
What really stabilizes NaCl is the crystal lattice, where
there are strong attractive forces between the positively
charged Na+ ions and the negatively charged Cl- ions. This
is the lattice energy, which is the energy required to break
the NaCl lattice up into gaseous ions.
NaCl(s) = Na+(g) + Cl-(g) ΔHlattice = +788 kJ/mol
8.3 Covalent bonding.
Here the two atoms share the electrons
to achieve a covalent bond.
two pairs of electrons equally shared
between the two oxygen atoms
Multiple Bonds and bond order:
The sharing of a single pair of electrons
consititutes a single bond. Sharing of two
pairs of electrons constitutes a double bond,
and sharing three pairs of electrons
constitutes a triple bond.
H:H
Single bond
.. ..
:O::O:
double bond
:N:::N:
triple bond
Bond order: a single bond has bond order =
1, a double bond has bond order = 2, and a
triple bond has a bond order = 3.
Examples of Lewis dot diagrams:
Methane, CH4:
One shared
pair of electrons
= single bond
Carbon has four
valence electrons (red)
H
H
C
H
H
Hydrogens
achieve two
electrons like He
Carbon achieves
octet of electrons
single line
= single bond
Examples of Lewis dot diagrams:
Carbon dioxide: (CO2)
Carbon has four
valence electrons (red)
oxygens have six
valence electrons (black)
O=C=O
double line =
double bond
Carbon and both oxygens
achieve an octet of electrons
two shared
pairs of electrons
= double bond
Examples of Lewis dot diagrams:
Sulfur dioxide: (SO2)
double
bond?
single
bond?
O=S-O
(or O-S=O ?)
SO2 is an example where a
actual structure is average
molecule can be written in
of the two (bond order = 1½) :
two ways and actual structure
is the average of the two. This
is called RESONANCE (see later)
O
S
O
Note that as the bond order between
two atoms goes up, the bond length
gets shorter. Stronger bonds tend to be
shorter:
increasing bond order:
single
double
triple
N-N
N=N
N≡N
1.47 Å
1.24 Å
1.10 Å
bonds get shorter
bonds get stronger
Table. Average bond lengths (Å) for some
single and multiple bonds:
Bond Length
C-C
1.54
C=C
1.34
C≡C
1.20
C-N
1.43
C=N
1.38
C≡N
1.16
C-H
1.10
H-H
0.74
Bond Length
C-O 1.43
C=O 1.23
C≡O 1.13
N-N 1.47
N=N 1.24
N≡N 1.10
N-H 1.01
C-Br 1.94
Bond Length
N-O 1.36
N=O 1.22
O-O 1.48
O=O 1.21
O-H 0.96
C-F 1.38
C-Cl 1.78
C-I 2.14
Among the following examples, which
bond is shortest?
H-H
C-H
S
Cl-Cl
C C
Bond length depends on (a) radii of the bonded atoms
remember that the atomic radii
decrease along a period in the
P.T.
(b) the number of bonds between atoms
8.4 Bond Polarity and Electronegativity.
The concept of
electronegativity was
developed by Linus
Pauling. Electronegativity
is the ability of an element
to attract electrons to itself
in a molecule.
Electronegativity increases
across the periodic table
and is at a maximum in the
top right hand corner at
fluorine, and is at a
minimum at the bottom left
hand corner at Cesium.
Linus Carl Pauling
(1901-1994)
Electronegativities of the Elements
Electronegativities of some main group
elements:
H
2.1
Li
1.0
Na
0.9
K
0.8
Rb
0.8
Be
1.5
Mg
1.2
Ca
1.0
Sr
1.0
B
2.0
Al
1.5
Ga
1.6
In
1.7
C
2.5
Si
1.8
Ge
1.8
Sn
1.8
N
3.0
P
2.1
As
2.0
Sb
1.9
O
3.5
S
2.5
Se
2.4
Te
2.1
F
4.0
Cl
3.0
Br
2.8
I
2.5
Electronegativity (EN) and bond polarity:
The greater the difference in EN the more
ionic the bond. For EN differences less than
0.5 we can say the bond is covalent, and for
differences greater than 2.0 we can say the
bond is ionic, for 0.5-2.0 it is polar covalent:
Molecule:
F2
HF
LiF
______________________________________________________
EN diff:
Type:
4.0-4.0 = 0 4.0-2.1 = 1.9 4.0-1.0 = 3.0
non-polar
polar
ionic
covalent
covalent
Electronegativity and bonding:
Some typical ranges for EN differences are:
EN difference
range
bonding type Example EN difference
> 2.0
ionic
ionic
polar covalent
covalent
covalent
covalent
covalent
____________________________________________________________________________
0.5-2.0
<0.5
LiF
NaCl
HF
F-F
C-H
Li-Li
Au-C
4.0-1.0 = 3.0
3.0-0.9 = 2.1
4.0-2.1 = 1.9
4.0-4.0 = 0.0
2.5-2.1 = 0.4
1.0-1.0 = 0.0
2.5-2.4 = 0.1
______________________________________________________________________________
Bond polarity:
With greater EN difference, the electron density is pulled
onto the more electronegative of the two atoms forming
the molecule:
non-polar
covalent
electron
density
equally
shared
F2
polar
covalent
HF
ionic
LiF
electron
density
largely
on F
Among the following examples,
which bond is most polar?
all equally non-polar
C-F
S
(C-H is non-polar)
P-Cl
(furthest apart in
Periodic Table  largest DEN)
all equally non-polar
8.6 Resonance structures: Ozone (O3)
bond order = 1½
..
:
:
The ozone molecule can
be written with two
equivalent Lewis dot
structures. In such a
situation the actual
structure is the average
of these two structures,
with the two O-O bond
lengths equal.
O
=
O
O
O
:
O
O
:
:
O
::
: :
:
O
O
..
double arrow
= resonance
O-O bonds = 2.78 Å
O
O
O
The ozone molecule
Resonance structures – the nitrite anion: (NO2-)
In drawing up a Lewis dot diagram, if we are dealing with
an anion, we must put in an extra electron for each
negative charge on the anion:
negative charge
on anion
One extra electron
in Lewis dot
N
diagram because
O
Bond order
O
of single negative
= 1½
charge on anion
:
:
:
::
..
:
:
:
-
Two resonance structures
-
:
:
:
O
..
N
O
:
:
O
::
: :
:
O
..
N
=
-
N
O
O
average structure
The nitrate anion:
O
:
Number of canonical structures
:
O
: O:
:
.. N O
O..
:
:
O
O
N
B.O. = 1
..
:
O
N
..
:
:
:
:
average bond
order (B.O.)=
2 + 1 + 1 = 1⅓
3
O
..
O
B.O. = 1
..
:
O
N
:
B.O. = 2
-
..
..
O
:
..
to work out bond order,
pick the same bond in
each structure and
average the bond order
for that bond
Resonance in benzene.
H
C
H
C
C
H
C
C
C
H
H
H
H
H
C
C
C
C
H
H
or
C
C
H
H
There are two
canonical structures
for benzene, which
means that the C to C
bonds have a bond
order of (2+1)/2 = 1.5.
The benzene ring has
a very high stability
due to this resonance,
which is called
aromaticity.
Short-hand versions for the benzene ring
8.7. Exceptions to the octet rule.
BF3. This can be written as F2B=F with three
resonance structures. To complete its octet, BF3
readily reacts with e.g. H2O to form BF3.H2O. The
actual structure of BF3 appears not to involve a
double bond and does not obey the octet rule:
Possible resonance
structure for BF3,
but is not important
as this would
involve the very
electronegative
F donating e’s to B
Best representation of
BF3 with B
having only
6 electrons
in its valence
shell
Exceptions to the octet rule: free radicals
There are some molecules that do not obey the octet
rule because they have an odd number of electrons.
Such molecules are very reactive, because they do
not achieve an inert gas structure, and are known as
free radicals. Examples of free radicals are chlorine
dioxide, nitric oxide, nitrogen dioxide, and the
superoxide radical:
odd electrons
nitric oxide
chlorine dioxide
Exceptions to the Octet rule: Heavier atoms (P, As,
S, Se, Cl, Br, I) may attain more than an octet of
electrons:
Example: PF5.
In PF5, the P atom has ten electrons in its valence
shell, which occurs commonly for heavier non-metal
atoms:
leave off F
F
electrons not
shared with P
P
F
F
F
P has
10 valence
electrons
F
PF5
Many phosphorus compounds do obey the
octet rule:
PF3 and [PO4]3- :
three blue electrons are
from charge on anion
Some compounds greatly exceed an
octet of electrons:
IF7
XeF6
(both I and Xe have 14 valence e’s)
(Think about [XeF8]2-)
8.8 Strengths of Covalent Bonds.
The strength of a covalent bond is measured
by the energy required to break that bond
into atoms:
Cl-Cl(g) → 2 Cl(g) ΔH = 242 kJ/mol
The strength of the C-H bond comes from:
CH4(g) → C(g) + 4 H(g) ΔH = 1660 kJ/mol
Divide by 4 = 415 kJ/mol ( a ‘best’ value is
413 kJ/mol)
Table of bond enthalpies (kJ/mol):
Single bonds:
C-H 413
C-C 348
C-N 293
C-O 358
H-H 436
Multiple bonds:
C=C 614
C≡C 839
C=N 615
C-F
C-Cl
C-Br
C-I
H-F
485
328
276
240
567
C≡N 891
C=O 799
C≡O 1072
N-H
N-N
O-H
O-O
HCl
391
163
463
146
431
N=N 418
N≡N 941
N=O 607
F-F
Cl-Cl
Br-Br
I-I
H-I
155
242
193
151
299
O=O 495
S=O 523
S=S 418
Which of the following bonds do you think would be most difficult
(require the largest energy) to break?
H-H
C-F
S
Cl-Cl
C C
The shortest bonds tend to be the hardest to break in
a series of similar bonds. Bonds involving F also tend
to be very strong, because greater electronegativity
difference leads to stronger bonds.
Bond enthalpies and the enthalpy of
reactions.
The enthalpy of a reaction can be
calculated by application of Hess’s Law
(ΔH for overall reaction is sum of H for
the individual steps).
ΔHrxn = Σ(bond enthalpies of bonds
broken) – Σ(bond enthalpies of bonds
formed)
Example:
Consider the reaction between CH4 and Cl2 to
give methyl chloride:
CH3-H
+ Cl-Cl →
Bonds broken:
Bonds formed:
ΔHrxn
=
CH3-Cl + H-Cl
C-H =
Cl-Cl =
C-Cl =
H-Cl =
413 kJ/mol
242 kJ/mol
328 kJ/mol
431 kJ/mol
413 + 242 - 328 - 431 = -104 kJ/mol
Estimate DHo for the following reaction:
CH4 (g)
+
2 O2 (g)
→
CO2 (g)
+
2 H2O (g)
DHorxn = Σ n x Dbroken – Σ m x Dformed
DHorxn = [4 x 414 + 2 x 498] – [2 x 803 + 4 x 463] = - 806 kJ