Download Chemical Bonding

Chemical Bonding
Chapter 6
Chemical Bonding & Structure
Molecular bonding and structure play
the central role in determining the
course of chemical reactions.
Bonds
Forces that hold groups of
atoms together and make
them function as a unit.
Bond Energy
- It is the energy required to break a
bond.
- It gives us information about the
strength of a bonding interaction.
Bond Energies
Bond breaking requires energy
(endothermic).
Bond formation releases energy
(exothermic).
Chemical Bonds
Chemical Bond
Ionic
Cation
Anion
Covalent
Molecule
Ionic Bonds
-
Formed from electrostatic attractions of
closely packed, oppositely charged ions.
-
Formed when an atom that easily loses
electrons (metal) reacts with one that has a
high electron affinity(nonmetal).
-
2Na(s) + Cl2(g) ----> 2Na+(aq) + 2Cl-(aq)
Figure 11.8: The structure of lithium fluoride
Figure 11.1: The formation of a bond between two
hydrogen atoms
Covalent Bonding
Covalent bonds are formed by sharing
electrons between nuclei.
H. + .H ----> H-H
2 hydrogen atoms
hydrogen molecule
Types of Covalent Bonds
Polar covalent bond -- covalent bond in which
the electrons are not shared equally
because one atom attracts them more
strongly than the other. A dipole moment
exists. HOH, HCl, & CO
Nonpolar covalent bond -- covalent bond in
which the electrons are shared equally
between both atoms. No dipole moment
exists. CO2, CH4, & Cl2
Electronegativity
The ability of an atom in a molecule
to attract shared electrons to itself.
As electronegativity increases, the
attraction for electrons increases.
Fluorine has the highest value at 4.0
and cesium and francium are lowest
at 0.7.
Increasing electronegativity
08_132
H
Decreasing electronegativity
2.1
Li
Be
B
C
N
O
F
1.0
1.5
2.0
2.5
3.0
3.5
4.0
Na
Mg
Al
Si
P
S
Cl
0.9
1.2
1.5
1.8
2.1
2.5
3.0
K
Ca
Sc
Ti
V
Cr
Mn
Fe
Co
Ni
Cu
Zn
Ga
Ge
As
Se
Br
0.8
1.0
1.3
1.5
1.6
1.6
1.5
1.8
1.9
1.9
1.9
1.6
1.6
1.8
2.0
2.4
2.8
Rb
Sr
Y
Zr
Nb
Mo
Tc
Ru
Rh
Pd
Ag
Cd
In
Sn
Sb
Te
I
0.8
1.0
1.2
1.4
1.6
1.8
1.9
2.2
2.2
2.2
1.9
1.7
1.7
1.8
1.9
2.1
2.5
Cs
Ba
La-Lu
Hf
Ta
W
Re
Os
Ir
Pt
Au
Hg
Tl
Pb
Bi
Po
At
0.7
0.9
1.0-1.2
1.3
1.5
1.7
1.9
2.2
2.2
2.2
2.4
1.9
1.8
1.9
1.9
2.0
2.2
Fr
Ra
Ac
Th
Pa
U
Np-No
0.7
0.9
1.1
1.3
1.4
1.4
1.4-1.3
(a)
Increasing electronegativity
H
Decreasing electronegativity
2.1
Li
1.0
Na
0.9
K
B
Be
2.0
1.5
Al
Mg
1.2
Ca
Sc
Ti
V
Cr
Mn
Co
Ni
Cu
1.8
1.9
1.9
1.9
1.0
1.3
1.5
1.6
1.6
0.8
Y
Zr
Nb
Mo
Tc
Rh
Pd
Sr
Ru
Ag
Rb
1.6
1.8
1.9
2.2
2.2
2.2
1.9
W
Re
Os
Ir
Pt
1.7
1.9
2.2
2.2
2.2
0.8
Cs
1.0
Ba
1.2
1.4
La-Lu
Hf
Ta
1.5
1.5
Fe
0.7
0.9
1.0-1.2
1.3
Fr
Ra
Ac
Th
Pa
U
Np-No
1.1
1.3
1.4
1.4
1.4-1.3
0.7
0.9
Au
2.4
Zn
Si
P
1.5
1.8
2.1
Ga
Ge
As
4.0
3.5
3.0
2.5
F
O
N
C
S
2.5
Se
2.4
Cl
3.0
Br
2.8
1.6
1.8
2.0
Cd
In
Sn
Sb
1.7
1.7
1.8
1.9
2.1
Hg
Tl
Pb
Bi
Po
At
1.9
1.8
1.9
1.9
2.0
2.2
1.6
Te
(b)
Pauling Electronegativity Values
I
2.5
Electronegativity values for selected elements. See Figure 11.3
on page 334 in Zumdahl.
Percent Ionic Character
xA  xB
100% 
% Ionic Character (IC) 
xA
where xA is the larger electronegativity and xB
is the smaller value.
Watch significant figures!!!
Ionic Bond
Polar Covalent
Nonpolar Covalent
% IC > 50 %
% IC 5 - 50 %
% IC < 5 %
Percent Ionic Character
What type of bonding & % ionic character
does KCl have? Ionic
xA  xB
100% 
% Ionic Character (IC) 
xA
3.0  0.8
100% 
% Ionic Character (IC) 
3.0
2.2
100%
% Ionic Character (IC) 
3.0
% Ionic Character (IC)  73%
Percent Ionic Character
What type of bonding & % ionic character
does HOH have? Polar covalent
xA  xB
100% 
% Ionic Character (IC) 
xA
3.5  2.1
100%
% Ionic Character (IC) 
3.5
1.4
100%
% Ionic Character (IC) 
3.5
% Ionic Character (IC)  40.%
Percent Ionic Character
What type of bonding & % ionic character
does N2 have? Nonpolar covalent
xA  xB
100% 
% Ionic Character (IC) 
xA
3.0  3.0
100% 
% Ionic Character (IC) 
3.0
0
100%
% Ionic Character (IC) 
3.0
% Ionic Character (IC)  0%
Three Possible Types of Bonds
Nonpolar Covalent
(Electrons equally
shared.)
Polar Covalent
(Electrons shared
unequally.)
Ionic
(Electrons are
transferred.)
Figure 11.2: Probability representations of the
electron sharing in HF
Polarity
A molecule, such as HF, that has a center
of positive charge and a center of negative
charge is said to be polar, or to have a
dipole moment.
H F
+
partial positive charge

partial negative charge
08_131
H
F

H




 H
F
F



 H
F

F

H
F
 H
H


F
H



 H
F
F

 H

(a)
F
(b)
The Effect of an electric field on hydrogen fluoride molecules.
08_133


+
H
O



H

(a)
(b)
Dipole Moment for the water molecule.
Polar Water Molecule
The polarity of water allows it to dissolve
ionic materials which are essential for life.
The polarity of the water molecule allows
water molecules to attract each other
strongly (hydrogen bonds). Because of
this fact water remains as a liquid at room
temperatures and allows the existence of
life as we know it.
08_134
3
+

N
H
H


H




(a)
(b)
Dipole moment for the ammonia molecule.
08_151
Nonpolar molecule--zero dipole moment.
Achieving Noble Gas Electron
Configurations (NGEC)
Two nonmetals react: They share
electrons to achieve NGEC.
A nonmetal and a representative group
metal react (ionic compound): The
valence orbitals of the metal are emptied
to achieve NGEC. The valence electron
configuration of the nonmetal achieves
NGEC.
Noble Gas Configuration
When a Group I, II, or III metal reacts with a
nonmetal to form a binary ionic
compound, the nonmetal gains electrons to
obtain the configuration of the next noble
gas. The metal loses electrons to gain the
configuration of the previous noble gas.
Na ----> Na+ + e- configuration of Ne
Cl + e- ----> Cl-
configuration of Ar
Noble Gas Configuration
Continued
When two nonmetals react to form a covalent
bond, they share electrons to form noble
gas configurations for both.
..
..
H. + : Cl : ----> H-- Cl:
..
.
Hydrogen gains the noble gas configuration of
helium, while Chlorine gains the
configuration of Argon.
Anion Size
Anions are always larger than the parent
atom because they have added
electrons which repel each other. As
well, the number of protons is less
than the number of electrons so they
are not held as tightly.
Cation Size
Cations are always smaller than the
parent atom because they have lost an
entire electron shell. As well, the
number of protons is greater than the
number of electrons so the electrons
are held tighter.
Relative sizes of some ions and their parent atoms.
Lewis Structure
-
Shows how valence electrons are arranged
among atoms in a molecule.
-
Reflects central idea that stability of a
compound relates to noble gas electron
configuration.
-
Developed by G.N. Lewis in 1902.
Lewis Structures
Na. sodium atom
..
.S: sulfur atom
.
[Na]+ sodium ion
.. 2 
[ : S:] sulfide ion
..
Lewis Structures
Ionic Compounds
1
K
..
1
: Br:
..
Covalent
Compounds
..
..
: F: F:
..
..
In ionic compounds, electrons are transferred
and ions are formed. In covalent compounds,
electrons are shared to form a molecule.
Potassium gains the stability of argon,
bromine of krypton, and fluorine of neon.
Lone Pairs & Bonding Pairs
..
..
: F F:
..
..
Electrons shared between atoms are
bonding pairs. Electrons that are not
involved in bonding are called lone pairs.
Each fluorine has three lone pairs and
one bonding pair shared between them.
Octet
Neon does not form bonds because it has a
full outer shell of electrons--an octet. An
octet is four pairs of electrons and
represents extra stability for atoms and
ions.
Rules for Writing Lewis
Structures
•
Sum the valence electrons from all the
atoms.
•
Use a pair of electrons to form a bond
between each pair of bound atoms.
•
Arrange remaining electrons to satisfy the
duet rule for hydrogen and the octet rule
for the second-row elements.
Lewis Structures
NO+
•
5 e- + 6 e- - 1 e- = 10 e-
•
[:NO:]+
•
Each atom has an octet and is satisfied.
Single, Double, & Triple Bonds
Single bonds -- one shared pair of
electrons.
Double bonds -- two shared pairs of
electrons.
Triple bonds -- three shared pairs of
electrons.
•Bond Strength = Triple > Double > Single
–For bonds between same atoms, CN >
C=N > C—N
–Though Double not 2x the strength of
Single and Triple not 3x the strength of
Single
•Bond Length = Single > Double > Triple
–For bonds between same atoms, C—N >
C=N > CN
Comments About the Octet Rule
-
2nd row elements C, N, O, F observe the
octet rule.
-
2nd row elements B and Be often have fewer
than 8 electrons around themselves - they are
very reactive.
-
3rd row and heavier elements CAN exceed
the octet rule using empty valence d orbitals.
-
When writing Lewis structures, satisfy octets
first, then place electrons around elements
having available d orbitals.
Electron Deficient Molecules
Beryllium chloride -- BeCl2 -- is electron deficient
with four electrons. It forms a linear molecule.
Boron trifluoride -- BF3 -- is electron deficient
with six electrons. It forms a trigonal planar
molecule.
See page 351 for the reaction between boron
trifluoride and ammonia.
Four Failures of Lewis
Structures
Lewis Structures cannot adequately explain:
1. electron-deficient molecules.
2. the paramagnetism of oxygen and other
similar substances.
3. odd-electron molecules.
4. resonance.
Odd-Electron Molecules
NO2
• contains 17 electrons.
• cannot satisfy the octet rule.
• a more sophisticated model is neededthe molecular orbital model.
Resonance
Occurs when more than one valid Lewis
structure can be written for a particular
molecule.
These are resonance structures. The actual
structure is an average of the resonance
structures called a resonance hybrid.
See the resonance structures for the nitrate
ion on page 362 in Zumdahl.
Resonance
Resonance structures have Lone Pairs and
Multiple Bonds in different positions.
The actual molecule is a combination of all the
resonance forms – it does not resonate between
the two forms, though we often draw it that
way!
••
••
••
•• O ••
••
•• S •• O
••
••
•• O
••
••
••
••S ••
•• O ••
Stereochemistry
The study of the threedimensional arrangement
(molecular structure) of atoms or
groups of atoms within molecules
and the properties which follow
such arrangement.
VSEPR Model
Valence Shell Electron Pair
Repulsion -- The structure
around a given atom is
determined principally by
minimizing electron pair
repulsions.
Predicting a VSEPR Structure
1. Draw Lewis structure.
2. Put pairs as far apart as possible.
3. Determine positions of atoms from the
way electron pairs are shared.(Parent
Geometry)
4. Determine the name of molecular structure
from positions of the atoms.(Actual
Geometry)
Molecular Geometry
Parent Geometry is
Actual Geometry is the
electron pair
arrangement about
the central atom.
arrangement of atoms
about the central
atom.
•linear
•linear
•trigonal planar
•bent
•tetrahedral
•trigonal pyramid
08_142
Lone
pair
N
N
H
H
H
(a)
(b)
Lone pair of electrons on the ammonia molecule.
08_143
Lone pair
Bonding
pair
O
Bonding
pair
O
H
(a)
H
Lone pair
(b)
O
H
(c)
H
Lone pairs on the water molecule.
VSEPR
Two pairs of electrons are placed 180o apart -linear arrangement.
Three pairs of electrons are placed 120o apart
-- trigonal planar arrangement.
Four pairs of electrons are placed 109.5o apart
-- tetrahedral arrangement.
Double bonds and triple bonds count as one
effective pair of electrons.
Electron pair arrangement is the parent geometry. Molecular
structure is the actual geometry.
Parent & Actual Geometry
When every pair of electrons on the central
atom is shared with another atom, the
parent and actual geometry are the same.
When one or more pair of electron pairs
around a central atom are unshared(lone
pairs), the parent and actual geometry are
different.
VSEPR Model Summary
•
Determine the Lewis structure(s) for the
molecule.
•
For molecules with resonance structures, use any
of the structures to predict the molecular
structure.
•
Sum the electron pairs around the central atom to
determine the parent geometry.
•
The arrangement of the pairs is determined by
minimizing electron-pair repulsions.(Actual
Geometry)
VSEPR Model Summary
(Continued)
Lone pairs require more space
than bonding pairs since
they are tightly attracted
to only one nucleus. Lone
pairs produce slight
distortions of bond angles
less than 120o.