Chemical Bonding
Why atoms bond and bond types
Lewis structures
●Resonance structures
●Shapes of molecules and ions
●Some example problems
●
Chemical Bonding
●
Why do atoms bond?
Chemical compounds are formed by the joining of
two or more atoms. A stable compound occurs when
the total energy of the combination has lower
energy than the separated atoms.
The bound state implies a net attractive force
between the atoms ... a chemical bond.
Chemical Bonding
●
Why do atoms bond?
The two extreme cases of chemical bonds are:
Covalent bond: bond in which one or more pairs
of electrons are shared by two atoms.
Ionic bond: bond in which one or more electrons
from one atom are removed and attached to
another atom, resulting in positive and negative
ions which are held together by Coulomb force.
Chemical Bonding
●
Why do atoms bond?
Other types of bonds include metallic bonds and
hydrogen bonds.
Chemical Bonding
●
Why do atoms bond?
processes are spontaneous if they result in a
system with lower potential energy
●
chemical bonds form because they lower the
potential energy between the charged particles
that compose atoms
●
the potential energy between charged particles
is directly proportional to the product of the
charges
●
Chemical Bonding
●
Why do atoms bond?
the potential energy between charged particles
is inversely proportional to the distance
between the charges
●
1 q1⋅q2
E potential=
4 π ϵ0 r
for charges with the same sign, Epotential is
positive and the magnitude gets less positive as
the particles get farther apart
●
Chemical Bonding
●
Why do atoms bond?
for charges with the opposite signs, Epotential is
negative and the magnitude gets more negative as
the particles get closer together
●
remember: the more negative the potential
energy, the more stable the system becomes
●
1 q1⋅q2
E potential=
4 π ϵ0 r
Chemical Bonding
Potential energy
●
Why do atoms bond?
Two charges of
same sign
Distance
© 2012 K. Brown
Chemical Bonding
Potential energy
●
Why do atoms bond?
Two charges of
opposite sign
Distance
© 2012 K. Brown
Chemical Bonding
●
Why do atoms bond?
a chemical bond forms when the potential energy
of the bonded atoms is less than the potential
energy of the separate atoms
●
have to consider following interactions:
● nucleus-to-nucleus repulsion
● electron-to-electron repulsion
● nucleus-to-electron attraction
●
Chemical Bonding
●
Why do atoms bond?
Chemical Bonding
Lewis Structures
G.N. Lewis 1875 - 1946
Chemical Bonding
Lewis Structures
G. N. Lewis (in 1916 .. before modern quantum
mechanics!) suggested that covalent bonds formed
to enable elements to attain a "noble gas
configuration".
We can extend this idea to ionic compounds, in a
compound such as sodium chloride, one element
loses electron(s) to gain this stable electronic
configuration whilst the other gains electron(s)
to achieve the same result.
Chemical Bonding
Lewis Structures
The Lewis theory of chemical bonding has as it's
basic ideas, several points:
*The electrons most involved in chemical bonding
to form compounds are the outer shell or valence
electrons.
*Electrons are sometimes completely transferred
between atoms in compound formation to form
ionic bonds and sometimes electrons are shared
between atoms in the formation of covalent
bonds.
Chemical Bonding
Lewis Structures
*The transfer or sharing of electrons happens in
such a way that each atom attains an especially
stable electron configuration ... usually that
of a noble gas with eight outer shell electrons
(the “octet rule”).
Chemical Bonding
Lewis Structures of Atoms
electron dot symbols
●
use symbol of element to represent nucleus and
inner electrons
●
use dots around the symbol to represent valence
electrons
● pair first two electrons for the s orbital
● put one electron on each open side for p
electrons
● then pair rest of the p electrons
●
Chemical Bonding
Lewis Structures of Atoms
Li
Be
B
C
N
O
F
Ne
Chemical Bonding
Lewis Structures of Compounds
The essence of the Lewis structure of a compound
is that, if it is ionic, a transfer of electrons
is shown and if covalent, a sharing of electrons
is shown.
Chemical Bonding
Lewis Structures of Compounds
Na
+
Cl
Na+ +
Cl-
Mg
+
O
Mg2++
O2-
Chemical Bonding
Lewis Structures of Compounds
H
+ H
H H
Shared electron
pair
Cl
+
Cl
Cl Cl
Unshared
electron pair
Chemical Bonding
Lewis Structures of Compounds
H
+ Cl
H
+
O
H
H Cl
H O
H
or
or
H Cl
H O
H
Chemical Bonding
Lewis Structures of Compounds
O
+ C
+
O
O
C
O
or
O
C
O
Chemical Bonding
Lewis Structures of Compounds
N
+
N
N
N
or
N
N
Chemical Bonding
Lewis Structures of Compounds
H
H
C
C
H
H
H
H
C
H
C
H
or
H
H
C
H
C
H
Chemical Bonding
Coordinate Covalent Bonds
Occasionally, the 'shared' pair of electrons in
a bond actually formally originate on one atom.
The bond is called in this case a coordinate
covalent bond
H
H N H + H+
H
H N H
H
+
Chemical Bonding
Coordinate Covalent Bonds
Other examples of compounds with coordinate
covalent bonds are BF4- , H3O+ and AlCl4-.
Chemical Bonding
Polar Covalent Bonds
Binary compounds with identical atoms have pure
covalent bonds since there can be no asymmetry
in their electron distribution.
In all other compounds in which the bonded atoms
are different, there will be an unequal sharing
of electrons such that the more electronegative
element will have a greater share of the
electrons.
Chemical Bonding
Polar Covalent Bonds
In the extreme, this unequal sharing becomes an
ionic bond where one atom has a
disproportionately large share of the electrons.
Chemical Bonding
Electronegativity
If difference in electronegativity between
bonded atoms is 0, the bond is pure covalent
● equal sharing
●
If difference in electronegativity between
bonded atoms is 0.1 to 0.4, the bond is nonpolar
covalent
●
Chemical Bonding
Electronegativity
If difference in electronegativity between
bonded atoms 0.5 to 1.9, the bond is polar
covalent
●
If difference in electronegativity between
bonded atoms larger than or equal to 2.0, the
bond is ionic
●
Chemical Bonding
Electronegativity
Chemical Bonding
Electronegativity
ENCl = 3.0
3.0 - 3.0 = 0
Pure Covalent
ENCl = 3.0
ENH = 2.1
3.0 – 2.1 = 0.9
Polar Covalent
ENCl = 3.0
ENNa = 1.0
3.0 – 0.9 = 2.1
Ionic
Chemical Bonding
Dipole Moment
A dipole is a pair of electric charges or magnetic
poles of equal magnitude but opposite polarity,
separated by some (usually small) distance.
Dipoles can be characterized by their dipole moment,
a vector quantity with a magnitude equal to the
product of the charge or magnetic strength of one of
the poles and the distance separating the two poles.
⃗ = q ⃗d
μ
Chemical Bonding
Dipole Moment
Even though the total charge on a molecule is zero,
the nature of chemical bonds is such that the
positive and negative charges do not completely
overlap in most molecules.
Such molecules are said to be polar because they
possess a permanent dipole moment. A good example is
the dipole moment of the water molecule. Molecules
with mirror symmetry like oxygen, nitrogen, carbon
dioxide, and carbon tetrachloride have no permanent
dipole moments.
Chemical Bonding
Percent Ionic Character
Percent ionic character=
actual dipole moment (μ )
dipole moment if the bond were ionic
Example: LiH m = 6.00
ionic dipole moment = 7.66
percent ionic character = (6.00/7.66) x 100
= 78.3%
Chemical Bonding
Percent Ionic Character
Another way to characterize the types of bonds in
compounds is to say that if the percent ionic
character is 50% or greater it is ionic … otherwise
it is covalent.
Chemical Bonding
Drawing Lewis Structures
Lewis structures are to this day, the most used
way to write the structural formulas of chemical
compounds where all the atoms and valence
electrons are shown.
Sometimes, people have a lot of trouble learning
how to do this. However, using the following
method, you should have very little trouble.
Chemical Bonding
Drawing Lewis Structures
1. Consult the molecular formula and sum up all
the valence electrons from the separate atoms.
Remember: the group number in the periodic table
= the number of valence electrons for an atom.
Add one for each (-) charge (extra electron)
Subtract one for each (+) charge (missing
electron)
Chemical Bonding
Drawing Lewis Structures
2. Choose the central atom(s). Almost always the
least electronegative atom is the central atom.
For example, in ClO2, the Cl is the central atom;
in SF5 the S is the central atom. Occasionally,
you will need to choose the unique atom, even
when it is the most electronegative: e.g., the O
in Cl2O. A wrong choice usually will be signaled
by your being unable to write a valid structure.
Chemical Bonding
Drawing Lewis Structures
Arrange the other atoms around the central atom,
in accord with the normal valences of the atoms.
That is, do not place more atoms around a
central one than it normally can bond to.
For first and second row elements, the maximum
valence = the group number through Group IV;
after that, it is 8 - (the group number).
Hydrogen is never the central atom. It usually
forms only one bond, so it must generally be in
the outer layer of atoms.
Chemical Bonding
Drawing Lewis Structures
3. Insert pairs of electrons between all pairs
of atoms that are to be bonded together. These
are the bonding electron pairs. If this uses up
all available electrons, go to Rule 6.
4. Place any remaining electrons on peripheral
atoms as unshared pairs, starting with the most
electronegative such atom. Fill this atom up to
an octet. Then proceed to the next most
electronegative, and so on. Remember that
hydrogens can only have two electrons, and so
cannot have any unshared pairs.
Chemical Bonding
Drawing Lewis Structures
5. If electrons still remain unused, place them
on the central atoms as unshared pairs, again
beginning with the most electronegative atom.
Fill that atom to an octet. Then proceed to the
next most electronegative, and so on.
Chemical Bonding
Drawing Lewis Structures
6. Examine the resulting structure. You will
observe one of four situations:
i. All of the atoms in the structure will have
octets of electrons, except the hydrogens,
which will have two electrons each. Go to Rule
7.
ii. The molecule has an odd number of electrons,
which results in one of the central atoms
having only 7 electrons. Go to Rule 7.
Chemical Bonding
Drawing Lewis Structures
iii. The central atom has Z > 11, and has other
than an octet. Go to Rule 7.
iv. The central atom is a second or third row
element and
a. the number of electrons in the molecule is
even and the central atom lacks an octet; or
b. the number of electrons in the molecule is
odd and the central atom has fewer than 7
electrons.
Chemical Bonding
Drawing Lewis Structures
In either case, move an unshared pair from a
peripheral atom to make a double bond to the
central atom. If the central atom still has too
few electrons, move another pair from the same
atom to make a triple bond, or a pair from
another atom to make a second double bond.
Chemical Bonding
Drawing Lewis Structures
7. Examine every atom in the structure and
assign it a formal charge as follows:
formal charge = (number of valence electrons on
the neutral, uncombined atom) - (number of
covalent bonds to the atom in the current
structure) - (the number of unshared electrons
[not pairs!] on the atom in the current
structure)
= group number – number of bonds – number of
unshared electrons
Chemical Bonding
Drawing Lewis Structures
O
Cl P Cl
Cl
O : 6−6−1=−1
Cl : 7−6−1=0
P : 5−0− 4=+1
Chemical Bonding
Drawing Lewis Structures
7. Examine every atom in the structure and
assign it a formal charge as follows:
If there is a formal charge difference between
two atoms in the structure, decrease the
difference by using an unshared pair on the
peripheral atom to make an extra shared pair
between the atoms. Do NOT violate the octet rule
for C, N, O and F.
Chemical Bonding
Drawing Lewis Structures
O
Cl P Cl
Cl
O : 6−4−2=0
Cl : 7−6−1=0
P : 5−0−5=0
Chemical Bonding
Drawing Lewis Structures
8. Rules 1-7 may allow some collections of atoms
to form several valid structures with differing
arrangements of atoms. This is OK; the
alternative arrangements of atoms are called
structural isomers. Isomers are particularly
common among compounds of C, N, and O. For
example, two isomeric Lewis structures can be
written for C2H6O and four can be written for
HCNO.
Chemical Bonding
Drawing Lewis Structures
9. Application of Rule 1-7 also may lead to
multiple structures having the same arrangement
of atoms but different placement of electrons.
This is OK too. These are resonance structures:
the molecule actually resembles an average of
all of the structures rather than any single
one.
Chemical Bonding
Lewis Structure of HNO3
1. Count valence electrons:
N
5eH
1e3xO
18etotal: 24e-
Chemical Bonding
Lewis Structure of HNO3
2. Choose central atom and arrange peripheral
atoms and connect to central atom with a single
bond:
O
O
N
O
H
Chemical Bonding
Lewis Structure of HNO3
3. Use remaining electrons to complete octets,
starting with the most electronegative elements:
O
O
N
O
H
Chemical Bonding
Lewis Structure of HNO3
4. Is the octet rule obeyed on all atoms? If
not, 'borrow' non-bonding pairs peripheral atoms
to complete the octet:
O
O
N
1
O
O
H
O
N
2
O
O
H
O
N
3
O
H
Chemical Bonding
Lewis Structure of HNO3
Structure 3 is not acceptable .. too many
separated formal charges
O
0
−1
−1
O
O
N
+1
1
O
0
H
0
−1
O
N
O
+1 +1
3
H 0
Chemical Bonding
Lewis Structure of HNO3
Structures 1 and 2 are acceptable and are called
resonance structures
O
O
N
1
O
O
H
O
N
2
O
H
Chemical Bonding
Lewis Structure Examples
H2O, SiH4, NF3, HNO3, NO2-, H2SO4, POCl3, CO2 , SO2
CH4, C2H4, C2H6, C2H6O, ClOF3, ClO4-, BF3
Chemical Bonding
Resonance Structures
The Lewis structure of ozone (O3):
O
O
O
Ozone would appear to have one single bond, and
one double bond.
Chemical Bonding
Resonance Structures
However... known facts about the structure of
ozone
The bond lengths between the central oxygen and
the other two oxygens are identical:
Public Domain
Chemical Bonding
Resonance Structures
We would expect that if one bond was a double
bond that it should be shorter than the other
(single) bond
Since all the atoms are identical (oxygens)
which atom is chosen for the double bond?
O
O
O
or
O
O
O
Chemical Bonding
Resonance Structures
These Lewis structures are equivalent except for
the placement of the electrons (i.e. the
location of the double bond)
Equivalent Lewis structures are called resonance
structures, or resonance forms
O
O
O
O
O
O
Chemical Bonding
Resonance Structures
This indicates that the ozone molecule is
described by an average of the two Lewis
structures
O
O
O
O
O
O
Chemical Bonding
Resonance Structures
The important points to remember about resonance
forms are:
* The molecule is not rapidly oscillating
between different discrete forms
* There is only one form of the ozone
molecule, and the bond lengths between the
oxygens are intermediate between characteristic
single and double bond lengths between a pair of
oxygens
Chemical Bonding
Resonance Structures
* We draw two Lewis structures (in this
case) because a single structure is insufficient
to describe the real structure using the Lewis
method
Chemical Bonding
Exceptions to the Octet Rule
There are three general ways in which the octet
rule breaks down:
1. Molecules with an odd number of electrons
2. Molecules in which an atom has less than an
octet of electrons
3. Molecules in which an atom has more than an
octet of electrons
Chemical Bonding
Odd Number of Electrons
Nitrogen Monoxide
1. Total electrons: 6+5=11
2. Bonding structure:
N
O
3. Octet on "outer" element
N
O
Chemical Bonding
Odd Number of Electrons
Nitrogen Monoxide
4. Remainder of electrons (11-8 = 3) on
"central" atom:
N
O
Chemical Bonding
Odd Number of Electrons
Nitrogen Monoxide
5. There are currently 5 valence electrons
around the nitrogen. A double bond would place 7
around the nitrogen, and a triple bond would
place 9 around the nitrogen.
We appear unable to get an octet around each
atom
N
O
Chemical Bonding
Less than an Octet of Electrons
Beryllium hydride
1. Add electrons (2*1) + 2 = 4
2. Draw connectivities:
H
Be H
3. At this point we are finished. We cannot
obtain an octet of electrons around beryllium.
Chemical Bonding
More than an Octet of Electrons
Phosphorus pentachloride
1. Count electrons: 5 + 5x7 = 40
2. Bonding structure
Cl
Cl P
Cl
Cl
Cl
Chemical Bonding
More than an Octet of Electrons
Phosphorus pentachloride
3. Octets for peripheral elements
Cl
Cl P
Cl
Cl
Cl
Chemical Bonding
Shapes of Molecules
Thalidomide
Public Domain
Chemical Bonding
Shapes of Molecules
Public Domain
Chemical Bonding
Shapes of Molecules
Valence Shell Electron Pair Repulsion theory
(VSEPR) is a set of rules whereby the chemist
may predict the shape of an isolated molecule.
It is based on the premise that groups of
electrons surrounding a central atom repel each
other, and that to minimize the overall energy
of the molecule, these groups of electrons try
to get as far apart as possible.
Chemical Bonding
Shapes of Molecules
VSEPR is useful for predicting the shape of a
molecule when there are between 2 and 6
substituents around the central atom.
That means that there are only five unique
electronic geometries to remember. For each
electronic geometry, there may be a number of
different molecular geometries
Chemical Bonding
Shapes of Molecules
Electron Geometry: the geometry of all electron
groups around a central atom. Includes bonding
and non-bonding electron groups.
Molecular Geometry: the geometry of bonding
electron groups only.
Chemical Bonding
Shapes of Molecules
Since the molecular geometry is determined by
how many bonding and non-bonding electron groups
surround the central atom, the first thing one
needs to do is count how many of each there are.
Note that bonding "electron groups" does not
necessarily imply single bonds; it can mean
double or triple bonds as well.
Chemical Bonding
Shapes of Molecules
For the purpose of determining the molecular
shape, double and triple bonds are counted as a
single “electron pair”. So the CO2 molecule is
considered to have a single pair of bonding
electrons between the carbon and each oxygen,
for shape determination purposes.
O
C
O
Chemical Bonding
Determining the Molecular Shape
1. Draw the correct Lewis structure
2. Count the number of electron groups around
the molecule, bonding and nonbonding. Remember
that double and triple bonds count as one
bonding pair only.
3. Assign an electron groups geometry based on
step 2
Chemical Bonding
Determining the Molecular Shape
4. The molecular shape will be a 'sub shape' of
the electron group geometry of step 3. In other
words, when describing the shape of the molecule
the non-bonding electron pairs are not
considered.
Chemical Bonding
Chemical Bonding
Chemical Bonding
axial
equatorial
Chemical Bonding
axial
equatorial
Chemical Bonding
Determining the Molecular Shape
Recall that there are two shapes or geometries
to consider:
Electron group geometry which will be one of the
five basic geometries
Molecular geometry which will include only
bonding electron groups.
Chemical Bonding
Chemical Bonding
Determining the Molecular Shape
Examples:
CO2, SO2, NH3, PCl3,
CH4, C2H6, ClOF3
Chemical Bonding
Polarity of Molecules
●
in order for a molecule to be polar it must
● have polar bonds
● electronegativity difference - theory
● bond dipole moments - measured
● have an unsymmetrical shape
● vector addition
Chemical Bonding
Polarity of Molecules
polarity affects the intermolecular forces
of attraction
● therefore boiling points and solubilities
● like dissolves like
●
nonbonding pairs affect molecular polarity,
strong pull in its direction
●
Chemical Bonding
Polarity of Molecules
The H-Cl bond is polar. The bonding
electrons are pulled toward the Cl end of
the molecule. The net result is a polar
molecule.
●
The H-Cl bond dipole moment is represented
as a vector pointing towards the Cl (the
more electronegative element).
●
Chemical Bonding
Polarity of Molecules
© 2012 K. Brown
Chemical Bonding
Polarity of Molecules
In molecules with multiple bonds each bond
is polar and the overall polarity of the
molecule is arrived at by vector addition of
the bond moments.
Chemical Bonding
Polarity of Molecules
 A 
C=
B

A
+

B

C

A

B
Chemical Bonding
Polarity of Molecules
 A 
C=
B=0

A

B
+

A

B
Chemical Bonding
Polarity of Molecules
No net dipole moment
© 2012 K. Brown
Chemical Bonding
Polarity of Molecules
Net dipole moment
© 2012 K. Brown
Chemical Bonding
Polarity of Molecules
Net dipole moment
© 2012 K. Brown
Chemical Bonding
Polarity of Molecules
Non polar
Polar
Non polar
Chemical Bonding
Polarity of Molecules
Non polar
Polar
Chemical Bonding
Polarity of Molecules
SO3
NOCl
3.5O
3.0
3.0
Cl
N
O
3.5
1) polar bonds, N-O
2) asymmetrical shape
polar
3.5 O
S
O 3.5
1) polar bonds, all S-O
2) symmetrical shape
nonpolar
© 2012 K. Brown
Chemical Bonding
Bond Energies
chemical reactions generally involve
breaking bonds in reactant molecules and
making new bond to create the products
●
the DH°reaction can be calculated by comparing
the cost of breaking old bonds to the profit
from making new bonds
●
Chemical Bonding
Bond Energies
the amount of energy it takes to break one
mole of a bond in a compound is called the
bond energy
● in the gas state
● homolytically – each atom gets ½ bonding
electrons
●
Chemical Bonding
Bond Energies
the more electrons two atoms share, the
stronger the covalent bond
● C≡C (837 kJ) > C=C (611 kJ) > C−C (347 kJ)
● C≡N (891 kJ) > C=N (615 kJ) > C−N (305 kJ)
●
the shorter the covalent bond, the stronger
the bond
● Br−F (237 kJ) > Br−Cl (218 kJ) > Br−Br
(193 kJ)
● bonds get weaker down the column
●
Chemical Bonding
Bond Energies
the actual bond energy depends on the
surrounding atoms and other factors
●
we often use average bond energies to
estimate the DHrxn
● works best when all reactants and products
in gas state
●
Chemical Bonding
Bond Energies
bond breaking is endothermic, DH(breaking) =
+
●
●
bond making is exothermic, DH(making) = −
DHrxn = S (DH(bonds broken)) + S (DH(bonds
formed))
Chemical Bonding
Chemical Bonding
Chemical Bonding
Bond Energies
Estimate the enthalpy of the following reaction:
H2(g) + O2(g)
H2O2(g)
Reaction involves breaking 1mol H-H and 1 mol
O=O and making 2 mol H-O and 1 mol O-O
Chemical Bonding
Bond Energies
H2(g) + O2(g)
H2O2(g)
bonds broken (energy cost)
(+436 kJ) + (+498 kJ) = +934 kJ
bonds made (energy release)
2(464 kJ) + (142 kJ) = -1070 kJ
DHrxn = (+934 kJ) + (-1070. kJ) = -136 kJ
(Appendix DH°f = -136.3 kJ/mol)
Chemical Bonding
Bond Lengths
Bond lengths tend to follow bond energies
inversely. The higher the bond energy the
shorter the bond length.
Chemical Bonding
Bond Lengths
Chemical Bonding
Lattice Energies
1
Li ( s)+ Cl 2 ( g) → LiCl ( s)
2
Δ H f =−408.8 kJ / mol
+
Δ H I =+520 kJ / mol
-
Δ H eaf =−349 kJ / mol
Li ( g) → Li ( g)
Cl ( g) → Cl ( g)
Chemical Bonding
Lattice Energies
Enthalpy
Li + ( g)+Cl ( g)
1
+
Li ( g)+ Cl 2 ( g)
2
Δ H3
1
Li ( g)+ Cl 2 ( g)
2
Δ H2
1
Li ( s)+ Cl 2 ( g)
2
Δ Hf
Δ H1
LiCl ( s)
Δ H4
Li + ( g)+Cl - ( g)
Δ H5
Chemical Bonding
Lattice Energies
Δ H f = Δ H 1+ Δ H 2 + Δ H 3 + Δ H 4 + Δ H 5
Δ H 5 = Δ H f −( Δ H 1+ Δ H 2 + Δ H 3 + Δ H 4 )
Δ H 5 =−409−(+132+520+122−349)
=−834 kJ / mol
Chemical Bonding
Lattice Energies
LiCl −834 kJ / mol
NaCl −788 kJ / mol
KCl −701 kJ / mol
CsCl −657 kJ / mol
© 2012 K. Brown
Chemical Bonding
Lattice Energies
NaF −910 kJ / mol
CaO −3414 kJ / mol
© 2012 K. Brown
Chemical Bonding
Potential energy
Lattice Energies
1 q1⋅q2
E potential=
4 π ϵ0 r
Two charges of
opposite sign
Distance
© 2012 K. Brown
Examples
1. Draw the Lewis structures for OF2, N2F2, IOF3
and NO2.
2.Determine the shapes for OF2, N2F2, IOF3 and NO2
and whether or not they are polar molecules.