The Covalent Bond and Carbon

1
Review Atomic Structure, Chemical Bonding and
Intro. To Molecular Polarity
2
I.  Atomic Structure Atoms are primarily
composed of 3 sub atomic particles.
Sub Atomic Particle Charge Mass(amu)
proton (p) 1 1
neutron(n) 0 1
electron(e-) -1 ?0
3
An atom is neutral if e-s ps.

• If a neutral atom gains extra electron(s) then it
  becomes a negatively charged species called an
  anion.
• If a neutral atom loses electron(s) then it
  becomes a positively charged species called a
  cation.

4
An atom is completely characterized by two
numbers the atomic (Z) and the atomic mass
(A).

• Atomic (Z) - the of protons in the nucleus -
  responsible for identity of the element.
• Mass (A)- the total of protons plus neutrons.

5
Representing Atoms of an Element

• An atom may be represented as its Symbol preceded
  by its subscripted atomic number, Z, and its
  superscripted atomic mass number, A.

A
Symbol
Z
In the case of the element Carbon
12
C
6
6
Fig. 5-6, p.125
7
Fig. 5-7, p.125
8
Arrangement of the subatomic particles within the
atom

• At the center is the nucleus which contains the
  protons and neutrons.
• electrons may be thought of as traveling in
  concentric shells or energy levels about the
  nucleus.
• the energy of the shells increase as one proceeds
  away from the nucleus.

9
There is a max. of e-s that can be
accommodated in each shell.
Shell Max. capacity e-s
1 2
2 8
3 18
4 32
10
Shell diagram for neutral atom of Phosphorus (P)
15 p
16 n
11
Further development of atomic model.

• Each shell is composed of 1 or more subshells.
• Each shell has as many subshells as its own
  number.
• 1st shell has 1 subshell.
• 2nd shell has 2 subshells.
• 3rd shell has 3 subshells.
• 4th shell has 4 subshells.

12
There are only four different kinds of
subshells.These subshells are labeled, in order
of increasing energy, by the letters s, p, d
f.Each subshell can accomodate a different of
e-s
Energy Increases subshell e- capacity
s 2
p 6
d 10
f 14
13
Thus the total capacity of shell is distributed
amongst its subshells.
8
14
Shell/subshell diagram for phosphorus
15 p
1s
2s
2p
3s
3p
16 n
The ground state electron configuration for
phosphorus
1s2
2s2
2p6
3s2
3p3
15
Atomic subshells in order of increasing energy,
filling order.

• NOTE
• Although the 4th shell is higher in energy than
  the 3rd shell, not all subshells of the 4th shell
  are higher in energy than all subshells of the
  3rd shell. In fact, the highest subshell of the
  3rd shell (3d) is higher in energy than the
  lowest subshell of the 4th shell (4s)

16
Further development of atomic model
Our latest model of the atom identifies electrons
as dots traveling about the nucleus in concentric
subshells. The truth is that we can never know
the exact position of an electron at any point in
time. In 1926, however, Erwin Schrödinger (of
University of Zurich) developed a theory known
as Quantum mechanics in which he worked out a
mathematical expression to describe the motion of
an electron in terms of its energy.
17
Further development of atomic model
These mathematical expressions are called wave
equations since they are based upon the concept
that e-s show properties not only of particles
but also of electromagnetic waves. These wave
equations have a series of solutions called wave
functions which allow us to predict the volume
of space around a nucleus in which there is a
high probability of finding a particular e-.
This volume of space in which an electron is most
likely to be found is called an orbital.
18
Now, to fully develop our theory of atomic
structure we must understand that the subshells
(s, p, d, f) of our earlier atomic model
consist of orbitals that are not all concentric
in shape. Furthermore, any one orbital can only
accommodate 2 e-s. Consequently, the number of
orbitals that comprise a subshell can easily be
calculated by simply dividing the subshell
capacity by 2.
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Number of Orbitals in each Subshell

• Any s subshell has a capacity of 2 e-s
• The number of orbitals that comprise any s
  subshell is 1.
• Any p subshell has a capacity of 6 e-s
• The number of orbitals that comprise any p
  subshell is 3.
• Any d subshell has a capacity of 10 e-s
• The number of orbitals that comprise any d
  subshell is 5.
• Any f subshell has a capacity of 14 e-s
• The number of orbitals that comprise any f
  subshell is 7.

20
Orbitals (s p d f)

• All orbitals of the same kind have the same 3
  dimensional shape but different sizes. The size
  increases with the energy level. All s subshells
  consist of one s orbital that is spherically
  symmetrical about the nucleus. An s orbital can
  accommodate 2 e- This accounts for the 2e-
  capacity of the s subshell

21
s Orbitals
1s ORBITAL 2s ORBITAL
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Each p subshell actually consists of a set of
three p orbitals of equal energy px py pz.

• Each of the three p orbitals is dumbbell shaped
  and all are oriented in space perpendicular to
  one another.
• The max. capacity of each p orbital is 2e-.
• This accounts for the total capacity of the p
  subshell as being 6 e-s.

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Each p subshell consists of a set of three p
orbitals of equal energy, px py pz
25
Shown together the three p orbitals look like
this
26
The d subshell actually consists of a set of
five d orbitals of equal energy. Each d orbital
can hold a maximum of 2e-. This accounts for the
total capacity of the d subshell as being 10
e-s. We will not be focusing on the d orbitals
therefore their shapes and names need not be
memorized. However, FYI..
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Electron Spins

• Electrons spin on their axis
• Physics tells us that any charged species that
  spins, generates a magnetic moment. That is to
  say, it acts like a tiny bar magnet with a North
  and a South Pole.
• Furthermore, the Right Hand Rule tells us that
  if we wrap the fingers of our right hand around
  the spinning species, in the direction of the
  spin, then our thumb will be pointing to the
  magnetic north.

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S
N
N
S
30
Represeanting Electrons

• Therefore, because of their magnetic moments, we
  generally represent electrons using a single
  barbed arrow. The tip of the arrow points to the
  magnetic north of the electron.

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Atomic Orbitals in order of Increasing Energy
3d__ 3d__ 3d__ 3d__ 3d__
4s___
3px__ 3py__ 3pz__
3s ___
2px__ 2py__ 2pz__
2s___
1s___
ENERGY
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Ground - state electron configurations

• This refers to the lowest energy arrangement of
  e-s in orbitals about the nucleus.
• To obtain this ground - state electron
  configuration electrons are assigned to the
  orbitals of the previous slide according to the
  three rules.

34
Rules for Filling Orbitals

• Always fill the lowest energy orbitals first.
• The two electrons that occupy any orbital must
  have opposite spins.
• When filling orbitals of equal energy (those of
  the p,d,or f subshells) put one electron in each
  orbital with their spins parallel until all are
  half filled, then go back and pair them.

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Orbital Electron Configurations

• Write the orbital electron configuration for P
• Write the orbital electron configuration for O

1s2
2s2
2px2
3s2
3px1
3py1
3pz1
2py2
2pz2
1s2
2s2
2px2
2py1
2pz1
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Using the periodic table to write electron
configurations

• The P.T. is arranged such that each horizontal
  row (period) represents the filling of orbitals
  in their proper order.

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More information from the Periodic Table
Valence Electrons

• The term valence electron refers to the of e-s
  in the outermost energy level or shell of an
  atom.

40
For all main group elements the of the column
(family) of the Periodic Table in which the
symbol for the element occurs the of valence
electrons.
Element Number of Valence e- s
Na 1
B 3
Cl 7
41
Lewis Structures of Atoms

• These are shorthand techniques for emphasizing
  the outer shell or valence e-s of an atom by
  representing an atom as its symbol surrounded by
  its valence e-s, the e-s in the atoms outermost
  shell. Note that the symbol of the element
  represents the nucleus plus all inner shell e-s.

42
Write Lewis dot structures for carbon, hydrogen,
oxygen, nitrogen and chlorine.
carbon
C
nitrogen
N
hydrogen
H
chlorine
Cl
oxygen
O
43
Why do atoms react together to form compounds?

• Atoms react with one another to form compounds
  in an attempt to achieve the e- configuration of
  their nearest noble gas neighbor (family 8). The
  reason for this is that the e- configuration of
  the noble gases represents an extremely stable
  situation.

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Fig. 5-6, p.125
45
There are two ways in which atoms can bond
together so as to achieve the e- configuration of
their nearest noble gas neighbor.

1. They can loose or gain the necessary e-s and
   thereby become ions and ultimately form ionic
   bonds.
2. Two or more atoms can share e-s and form
   covalent bonds.

46
Ionic Bonds

• These are formed when ions anions/cations of
  opposite charge come together. Generally ionic
  compounds are formed between metals (left of
  step) and nonmetals (right of step).

47
Consider the formation of the ionic compound
magnesium bromide.

• Magnesium (Mg ) could achieve the e- config. of
  Neon by loosing 2e- .

? Mg 2 2e-
isoelectronic with Ne
48
Bromine could achieve the e- config. of krypton
by gaining one e-.
1e-
Note

• Consequently one magnesium combines with two
  bromine atoms to form MgBr2.

49
Note all atoms in MgBr2 are isoelectronic with
their nearest noble gas neighbor.

• Mg2 Br-1 MgBr2

50
Covalent Bond

• A covalent bond results from the sharing of an
  electron pair between two atoms.
• Whenever two atoms share a pair of e-s, it is as
  if each member of the bonded pair of atoms has
  gained an extra electron.
• As atoms bond together to become isoelectronic
  with their nearest noble gas neighbors, covalent
  bonds generally occur when two or more
  nonmetallic elements (right of step) bond
  together because the nearest noble gas neighbors
  for these elements lies ahead of them.
  Consequently, they all need to gain electrons to
  become isoelectronic with their nearest noble gas
  neighbors.

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How many hydrogen atoms bond to one carbon atom?
Kekulé or Lewis structure for covalently bonded
molecule
Can become isoelectronic with He by gaining 1e-
H
C
C
H
4

C
H
H
Can become isoelectronic with Ne by gaining 4e-
Lewis Structure for covalent molecule of CH4
H
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Rules for Creating Lewis Structures for more
Complicated Molecules

• Connect all atoms using single bonds
• Add the total of Valence Electons
• Subtract 2e-s from the total of Valence e-s
  for each single bond drawn in first step
• Sprinkle any remaining e-s so as to make all
  atoms isoelectronic with their nearest noble gas
  neighbors. This usually means 8 e-s. For H its
  2 e-s.
• If there are insufficient e-s to accomplish the
  previous step, make one or more nonbonded e-
  pairs perform double duty by forming multiple
  bonds.

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Now lets build the Ammonia Molecule
NH3
Connect all atoms using single bonds Total
Valence Electrons 8 Subtract 2e-s for each
single bond 8 (3 x 2) 2 e-s Sprinkle
remaining 2 e-s so that all atoms are
isoelectronic with their nearest noble gas
neighbor
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Lets Try CO2

• Connect all w single bonds
• Total Valence e-s 16
• Subtract 2e-s for each bond
• 16 (2 x 2) 12
• Sprinkle remaining e-s so that all atoms have
  8e-s. Peripheral atoms first.
• If the octet cannot be satisfied for all, force
  nonbonded pairs to perform double duty

55
Now Lets Try HCN

• Connect all w single bonds
• Total Valence e-s 10
• Subtract 2e-s for each bond
• 10 (2 x 2) 6 e-s
• Sprinkle remaining 6 e-s so that all atoms have
  8e-s. Peripheral atoms first.
• If the octet cannot be satisfied for all (except
  H), force nonbonded pairs to perform double duty

56
Lets Look at the Water Molecule
O

H
2
O
In the water molecule each oxygen is
isoelectronic with
Neon
In the water molecule each hydrogen is
isoelectronic with
Helium
57
Now Lets Try the Amino Acid Alanine -
NH2CH(CH3)COOH
O
H
H
C
O
H
N
C
H
H
C
H
H

• Connect all w single bonds, be careful not to
  exeeed the normal valences (combining capacities)
  for all atoms
• Total Valence e-s 36
• Subtract 2e-s for each bond
• 36 (2 x 12) 12 e-s
• Sprinkle remaining 12 e-s so that all atoms have
  8e-s.
• If the octet cannot be satisfied for all (except
  H), force nonbonded pairs to perform double duty

58
Amino Acids

• There are 20 different Amino Acids. Amino Acids
  differ from one another only in the nature of the
  R side chain.
• Different R Side groups gives different Amino
  Acids
• R side chain
• I
• H2H C COOH
• I
• H

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Examples of Amino Acids

• H
• I
• H2NC COOH
• I
• H glycine
• CH3
• I
• H2NC COOH
• I
• H alanine

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Different Types of R groups Different Amino
Acids

• Nonpolar R H, CH3, alkyl groups, aromatic
• O
• Polar ll
• R CH2OH, CH2SH, CH2CNH2,
• (polar groups with O-, -SH, -N-)
• Polar/Acidic
• R CH2COOH, or -COOH
• Polar/ Basic
• R CH2CH2NH2

61
Amino Acids and Proteins

• Amino Acids are the building blocks of proteins
• In fact proteins are simply combinations of amino
  acids linked in a head to tail fashon

62
Types of Proteins

• Type Examples
• Structural tendons, cartilage, hair, nails
• Contractile muscles
• Transport hemoglobin
• Storage milk
• Hormonal insulin, growth hormone
• Enzyme catalyzes reactions in cells
• Protection immune response

63
Kekulé or Lewis structure for water molecule.
Bonded electron pairs
Nonbonded Electron Pairs
O
H
H
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The Covalent Bond and Electronegativity

• The sharing of an e- pair between two atoms may
  be equal .
• If this is the case then the resulting covalent
  bond is a nonpolar covalent bond.
• If, on the other hand the sharing is unequal then
  a polar covalent bond results.

65
The reason for this variance in bond polarity is
due to the fact that different elements have
different tendencies to attract to themselves
extra electrons. In other words, each element
has a different electronegativity
Electronegativity
66
Electronegativity

• The tendency of an atom, when in combination
  with other atoms, to attract to itself the bonded
  (extra) e-s.

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• Electronegativity values increase from left to
  right across any horizontal row (period) of the
  P.T. and they decrease going down any vertical
  column (family) of the P.T.
• Consequently the most electronegative elements
  are
• N, O, F, Cl, Br

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Electronegativity values for selected elements.
70

• If two atoms are covalently bonded and one has a
  high electronegativity and the other has a low
  electronegative then the electron pair comprising
  that bond is not shared equally but spends more
  of its time closer to the more electronegative
  atom. The immediate result of this unequal
  sharing is that the more electronegative atom
  gains a partial negative charge (?-) while the
  less electronegative element gains a partial
  positive charge (? ). This type of bond is
  called a polar covalent bond.

71
The degree to which a covalent bond is polarized
is indicated by the electronegativity difference
between the two bonded atoms. Refer to next
slide for electronegativity values of elements.

• If the electronegativity difference is greater
  than .5 but less than 2.0 then the covalent bond
  is polar.
• If the electronegativity difference is less than
  .5 then the covalent bond is nonpolar.

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Polar Covalent Bonds in H2O
Electronegativity Difference Between Oxygen and
Hydrogen is
? -
1.4
O
H
H
?
?
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A molecule typical of those found in petroleum.
The bonds are not polar.
Electronegativity Difference Between Carbon and
Hydrogen is
0.4
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Ionic Bond and Electronegativity

• Consideration of electronegativity can
  demonstrate that ionic bonds are nothing more
  than an extreme case of a polar covalent bond. In
  fact
• if the electronegativity difference between two
  atoms is greater than 2.0, then any bond between
  these two atoms would be ionic.

75
Molecular Polarity

• If a molecule contains polar bonds, and if those
  polar bonds are located such that the ? charges
  are at one end of the molecule and the ? -
  charges are at the other end, then the molecule
  is a polar molecule. The measure of molecular
  polarity is a quantity called the dipole moment
  (D).

76
Like Dissolves Like

• Polar molecules dissolve in Polar Solvents
• Nonpolar molecules dissolve in nonpolar solvents
• Polar molecules do not dissolve in nonpolar
  solvent
• Nonpolar molecules do not dissolve in polar
  solvents

77
An oil layer floating on water. The oil is
nonpolar and the water is polar
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Polar water molecules interact with the positive
and negative ions of a salt. Ionicly bonded
materials are the extreme case of polar substances
79
The Covalent Bond In Organic Chemistry
CHAPTER 01 (FUNDAMENTALS Org. Chem) CONTINUED

• The covalent bond is of chief importance in
  organic chemistry

80
The Covalent Bond and Valence Bond Theory

• Valence Bond theory offers a description of the
  covalent bond in terms of the orbital model of
  the atom. Valence Bond theory maintains that
  covalent bonds are formed by an overlapping of
  two half-filled (1e-) atomic orbitals.
• Consider the formation of the H2 molecule from
  two isolated hydrogen atoms

81
Formation of the H2 molecule


The reaction is accompanied by the evolution of
104 kcal/mole H2 formed. This means that the
product (H2 molecule) is more stable than the
reactants (isolated H atoms) by 104 K cal/mole.
The bond strength of the H2 molecule is 104
kcal/mole. This means that it would take 104
kcal to rupture the bonds in 1 mole of H2
molecules.
82
The Valence Bond representation of covalent bond
formation
1
83
Valence Bond Theory identifies two types of
Covalent Bonds

• Sigma (d) Bonds the bond in the H2 molecule is
  a sigma bond. Sigma bonds result from the head
  to head overlap of two half filled atomic
  orbitals. Sigma bonds are cylindrically
  symmetrical about a line joining the two nuclei
• Pi (p) Bonds These result from the side-to-side
  overlap of two half filled atomic orbitals.

84
Orbital overlap to form Sigma (d) Bonds
85
The Sigma Bond in the H2, HCl and the Cl2
molecules
86
Formation of the Pi (p) Bonds
Pi bonds are always accompanied by a Sigma Bond.
A Pi Bond cannot form without first forming a
Sigma Bond.
87
Pi Bond
88
Consider the Formation of the O2 molecule
The Orbital electron configuration for oxygen is
1s2 2s2 2px2 2py1 2pz1

O

O
89
Hybridization of Atomic Orbitals
mix or hybridize

• Certain atoms their
  atomic orbitals before bonding to other atoms and
  forming molecules.
• The reason for this is that Hybridized Atomic
  Orbitals are more directional and offer more
  effective overlap than do unhybridized atomic
  orbitals. As the strength of a covalent bond is
  directly related to the extent of overlap of the
  two ½ filled atomic orbitals, hybridized atomic
  orbitals form stronger bonds than do unhybridized
  atomic orbitals.

90
Evidence for the Hybridization of Atomic Orbitals

• Consider the H2O molecule.
• We know that the electron pair geomery is
• We therefore know that its HOH bond angle is

Tetrahedral
109.5 degrees
91
However

• If a covalent bond results from the overlap of
  two ½ filled atomic orbitals and if the orbital
  electron configuration for Oxygen is
• Then the HOH bond angle should be 90 degrees and
  look like this

1s2 2s2 2px2 2py1 2pz1
or
92
and CH4

• The same proof of Hybridization can be obtained
  by comparing the actual shapes and bond angles in
  ammonia and methane to what they would be if N
  and C used their unhybridized atomic orbitals to
  bond the Hs.
• The orbital electron configurations for N and C
  are N 1s2 2s2 2px1 2py12pz1 and C 1s2 2s2
  2px12py1
• In each case the expected bond angle would be 90
  degrees as compared to the actual bond angle of
  109.5 degrees

93
Hybridizations States of Carbon

• Carbon can adopt any one of three hybridization
  states, sp3( tetrahedral molecular geometry),
  sp2(trigonal planar), sp (linear) depending upon
  the number of electron pairs about the carbon.
• If C has 4 electron pairs as in CH4 then it sp3
  hybridizes. Bond angles 109.5
• If C has three electron pairs as in ethene,
  it sp2 hybridizes.
• Bond
  angles 120
• If C has two electron pairs as in ethyne,
  it sp hybridizes. 180

94
Practice Problems

• Identify the hybridization states and the bond
  angles for each carbon atom in the following
  molecule.

sp3
sp3
sp2
sp
sp2
sp3
sp
sp
All sp3 Cs 109.5 degrees
All sp2 Cs 120 degrees
All sp Cs 180 degrees
95
sp3, sp2 and sp Hybridized Orbitals

• How are they formed from the atomic orbitals?
• What do they look like?

96
The orbital electron configurations for C is 1s2
2s2 2px1 2py1 2pz0
Formation of four sp3 hybrid orbitals
or
97
sp3 Hybrid Orbitals

• sp3 Hybrid orbitals are formed from one s and
  three p orbitals. Therefore, there are four
  large lobes.
• Each lobe points towards the vertex of a
  tetrahedron.
• The angle between the large lobs is 109.5?.
• All molecules with tetrahedral electron pair
  geometries are sp3 hybridized.

98
sp3 Hybrid Orbitals

• Tetrahedral e- pair geometry
• 109.5 bond angle

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sp2 Hybrid Orbitals

• Important when we mix n atomic orbitals we must
  get n hybrid orbitals.
• sp2 hybrid orbitals are formed with one s and two
  p orbitals. (Therefore, there is one
  unhybridized p orbital remaining.)
• The large lobes of sp2 hybrids lie in a trigonal
  plane.
• All molecules with trigonal planar electron pair
  geometries have sp2 orbitals on the central atom.

104
The orbital electron configurations for C is 1s2
2s2 2px1 2py1 2pz0
Formation of three sp2 hybrid orbitals
These three atomic orbitals mix to form three sp2
hybrid orbitals
All together the 3 sp2s and the unhybridized p
look like this
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• Show Brown and LeMay Clip here
• The Shape of Molecule Chapter

112
sp Hybrid Orbitals

• sp hybrid orbitals are formed with one s and one
  p orbitals. (Therefore, there are two
  unhybridized p orbital remaining.)

The lobes of sp hybrid orbitals are 180º apart.
113
The orbital electron configurations for C is 1s2
2s2 2px1 2py1 2pz0
Formation of two sp hybrid orbitals
These two atomic orbitals mix to form two sp
hybrid orbitals
All together the 2 sps and the 2 unhybridized
ps look like this
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Multiple Bonds

• A double bond (2 pairs of shared electrons)
  consists of a sigma bond and a pi bond.
• A triple bond (3 pairs of shared electrons)
  consists of a sigma bond and two pi bonds.

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Sample Problems

• Predict the hybridization, geometry, and bond
  angle for each atom in the following molecules
• Caution! You must start with a good Lewis
  structure!
• NH2NH2
• CH3-C?C-CHO

118
Acids and Bases The BrønstedLowry Definition

• The terms acid and base can have different
  meanings in different contexts
• The Arrhenius definition of an acid is any
  substance that increases the H conc, of water.
  An Arrhenius base increases the OH- conc. of
  water.
• The idea that acids are aqueous solutions
  containing a lot of H and bases are solutions
  containing a lot of OH- is not very useful in
  organic chemistry
• Instead, BrønstedLowry theory defines acids and
  bases by their role in reactions that transfer
  protons (H) between donors and acceptors

119
Brønsted Acids and Bases

• A Brønsted acid is a substance that donates a
  hydrogen ion (H) also called a proton.
• A Brønsted base is a substance that accepts the
  H
• proton is a synonym for H - loss of an
  electron from H leaving the bare nucleusa proton

120
The Reaction of HCl with H2O

• When HCl gas dissolves in water, a Brønsted
  acidbase reaction occurs
• HCl donates a proton to water molecule, yielding
  hydronium ion (H3O) and Cl?
• The reverse rxn. is also a Brønsted acidbase
  reaction between the conjugate acid and conjugate
  base

Acids are shown in red, bases in blue. Curved
arrows go from bases to acids
121
Conjugate Acids and Conjugate Bases

• Every acid has association with it a conjugate
  base formed from the acid by loss of a proton.
• The conjugate base of HCl is Cl-.
• Every base has associated with it a conjugate
  acid formed from the base by addition of a
  proton.
• The conjugate acid of H2O is H3O.

122
Acid Strengths

• Acid strengths are indicated by the extent to
  which they donate protons to water. This extent
  of proton donation is indicated by the
  equilibrium constant for the reaction of the acid
  with water.
• HA H2O A- H3O
• Keg A-H3O products
• HAH2O reactants
• The higher the keg, the greater the tendency for
  the acid to donate protons to water. Therefore,
  the higher the keg, the stronger the acid, the
  lower the keg, the weaker the acid

123
Ka Values

• As the H2O remains constant for most keq
  measurements, we may therefore rewrite the
  equilibrium expression using a new term called
  the acidity constant Ka.
• Ka keq H2O H3O A-
• HA
• Therefore, the stronger the acid, the greater the
  Ka, the weaker the acid, the lower the Ka.

124
pKa Values

• Acid strengths are usually defined in terms of
  pKa values
• pKa -log Ka
• The stronger the acid, the lower the pKa and the
  weaker the acid, the higher the pKa value.

125
Predicting AcidBase Reactions from pKa Values

• An Acid will react with a Base if and only if the
  conjugate acid that is formed is weaker (has a
  higher pKa value) than the original acid.
• This simple concept will allow one to predict
  whether or not an acid/base reaction will go by
  simply comparing the pKa values of the original
  acid and its conjugate acid. Will the following
  reaction go?

Yes!
126
Predicting Acid Base Reactions

• Will the following reaction go?

Yes!
127
Organic Acids and Organic Bases

• The reaction patterns of organic compounds often
  are acid-base combinations
• The transfer of a proton from a strong Brønsted
  acid to a Brønsted base, for example, is a very
  fast process and will always occur along with
  other reactions

128
Organic Acids

• Those that lose a proton from OH, such as
  methanol and acetic acid
• Those that lose a proton from CH, usually from a
  carbon atom next to a CO double bond (OCCH)

129
Organic Bases

• Have an atom with a lone pair of electrons that
  can bond to H
• Nitrogen-containing compounds derived from
  ammonia are the most common organic bases
• Oxygen-containing compounds can react as bases
  when with a strong acid or as acids with strong
  bases

130
Acids and Bases The Lewis Definition

• Lewis acids are electron pair acceptors and Lewis
  bases are electron pair donors
• The Lewis definition leads to a general
  description of many reaction patterns but there
  is no scale of strengths as in the Brønsted
  definition of pKa

131
Illustration of Curved Arrows in Following Lewis
Acid-Base Reactions
132
Lewis Acids and the Curved Arrow Formalism

• The Lewis definition of acidity includes metal
  cations, such as Mg2
• They accept a pair of electrons when they form a
  bond to a base
• Group 3A elements, such as BF3 and AlCl3, are
  Lewis acids because they have unfilled valence
  orbitals and can accept electron pairs from Lewis
  bases
• Transition-metal compounds, such as TiCl4, FeCl3,
  ZnCl2, and SnCl4, are Lewis acids
• Organic compounds that undergo addition reactions
  with Lewis bases (discussed later) are called
  electrophiles and therefore Lewis Acids
• The combination of a Lewis acid and a Lewis base
  can shown with a curved arrow from base to acid

133
Lewis Acids and Bases

• Acids accept electron pairs electrophile
• Bases donate electron pairs nucleophile

134
Lewis Bases

• Most oxygen- and nitrogen-containing organic
  compounds are Lewis bases because they have lone
  pairs of electrons
• Some compounds can act as both acids and bases,
  depending on the reaction

135
Drawing Chemical Structures

• Chemists use shorthand ways for writing
  structures
• Condensed structures C-H and C-C and single
  bonds aren't shown but understood
• If C has 3 Hs bonded to it, write CH3
• If C has 2 Hs bonded to it, write CH2 and so
  on. The compound called 2-methylbutane, for
  example, is written as follows
• Horizontal bonds between carbons aren't shown in
  condensed structuresthe CH3, CH2, and CH units
  are simply but vertical bonds are added for
  clarity

136
Skeletal Structures

• Cs are not shown. They are assumed to be at
  each intersection of any two lines (bonds) and at
  end of each line
• Hs bonded to Cs aren't shown Since carbon
  always has a valence of 4, we mentally supply the
  correct number of Hs by subtracting the of
  bonds shown from 4.
• All atoms other than C and H are shown
• See next slide for examples

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139
Multiple Bonds

• In condensed formulas double and triple bonds are
  drawn as they would be in a Lewis structure
  showing two dashes for a double bond and three
  dashes for a triple bond