chapter 1- chm 261 organic chemistry
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1-1
Organic Chemistry
1-2
Review of Molecular Structure
Chapter 1
1-3
Organic Chemistry
The study of the compounds of carbon Over 10 million compounds have been identified
• about 1000 new ones are identified each day! C is a small atom
• it forms single, double, and triple bonds• it is intermediate in electronegativity (2.5)• it forms strong bonds with C, H, O, N, and some metals
1-4
Schematic View of an Atom
• a small dense nucleus, diameter 10-14 - 10-15 m, which contains positively charged protons and most of the mass of the atom
• an extranuclear space, diameter 10-10 m, which contains negatively charged electrons
1-5
Electron Configuration of Atoms
Electrons are confined to regions of space called principle energy levels (shells)• each shell can hold 2n2 electrons (n = 1,2,3,4......)
Shell
Number of Electrons Shell
Can Hold
Relative Energiesof Electrons
in These Shells
3218 8 2
4321
higher
lower
1-6
Electron Configuration of Atoms
Shells are divided into subshells called orbitals, which are designated by the letters s, p, d, f,........• s (one per shell) • p (set of three per shell 2 and higher) • d (set of five per shell 3 and higher) .....
Shell Orbitals Contained in That Shell
3
2
1 1s
2s, 2px, 2py, 2pz
3s, 3px, 3py, 3pz, plus five 3d orbitals
1-7
Electron Configuration of Atoms
Aufbau Principle: • orbitals fill in order of increasing energy from lowest
energy to highest energy Pauli Exclusion Principle:
• only two electrons can occupy an orbital and their spins must be paired
Hund’s Rule: • when orbitals of equal energy are available but there
are not enough electrons to fill all of them, one electron is added to each orbital before a second electron is added to any one of them
1-8
Electron Configuration of Atoms
The pairing of electron spins
1-9
Electron Configuration of Atoms
Table 1.3 The Ground-State Electron Configuration of Elements 1-18
1-10
Lewis Dot Structures
Gilbert N. Lewis Valence shell:
• the outermost occupied electron shell of an atom Valence electrons:
• electrons in the valence shell of an atom; these electrons are used to form chemical bonds and in chemical reactions
Lewis dot structure: • the symbol of an element represents the nucleus and
all inner shell electrons• dots represent valence electrons
1-11
Lewis Dot Structures
Lewis Dot Structures for Elements 1-18
N OB
H
Li Be
Na
He
Cl
F
S
Ne
Ar
C
SiAl P
1A 2A 3A 4A 5A 6A 7A 8A
Mg ::
::
::
.
.
.
.
.
.
.
..
..
. .
.
.
.
:
:
:
::::::::
::::::.
:::
:
1-12
Lewis Model of Bonding
Atoms bond together so that each atom acquires an electron configuration the same as that of the noble gas nearest it in atomic number• an atom that gains electrons becomes an anion• an atom that loses electrons becomes a cation• the attraction of anions and cations leads to the
formation of ionic solids• an atom may share electrons with one or more atoms
to complete its valence shell; a chemical bond formed by sharing electrons is called a covalent bond
• bonds may be partially ionic or partially covalent; these bonds are called polar covalent bonds
1-13
Covalent Bonds
The simplest covalent bond is that in H2
• the single electrons from each atom combine to form an electron pair
• the shared pair functions in two ways simultaneously; it is shared by the two atoms and fills the valence shell of each atom
The number of shared pairs• one shared pair forms a single bond• two shared pairs form a double bond• three shared pairs form a triple bond
H H H-H+ • H0 = -435 kJ (-104 kcal)/mol•
1-14
Lewis Structures
To write a Lewis structure• determine the number of valence electrons
• Atom by atom analysis• determine the arrangement of atoms
• Different arrangements are almost always possible - isomers• connect the atoms by single bonds• arrange the remaining electrons so that each atom has
a complete valence shell• Use common bonding patterns as your guide
• show a bonding pair of electrons as a single line• show a nonbonding pair of electrons as a pair of dots
1-15
Stable Bonding in Neutral Molecules
In neutral molecules• hydrogen has one bond• carbon has 4 bonds and no lone pairs• nitrogen has 3 bonds and 1 lone pair• oxygen has 2 bonds and 2 lone pairs• halogens have 1 bond and 3 lone pairs• 3rd row and transition metal bonding patterns are more
complicated but we don’t encounter a lot of them here
1-16
Multiple Bonding
Bonded atoms can share more than one pair of electrons to form multiple bonds
Molecular oxygen (O2) has a double bond between the oxygen atoms
+ givesO O O O
1-17
Lewis Structures -
H2O (8) NH3 (8) CH4 (8) HCl (8)
C2H4 (12) C2H2 (10) CH2O (12) H2CO3 (24)
H-O-H H-N-HH
H-C-HH
HH-Cl
H-C C-HH
HC O
H
HC C
H
HO O
CH H
O
Ethylene
Hydrogen chlorideMethaneAmmoniaWater
Carbonic acidFormaldehydeAcetylene
1-18
Electronegativity
Electronegativity: • a very crude measure of an atom’s attraction for the
electrons it shares with another atom in a chemical bond
Pauling scale• generally increases left to right in a row• generally increases bottom to top in a column
1-19
Polar and Nonpolar Covalent Bonds
Although all covalent bonds involve sharing of electrons, they differ widely in the degree of sharing
We divide covalent bonds into• nonpolar covalent bonds• polar covalent bonds
Difference in ElectronegativityBetween Bonded Atoms Type of Bond
Less than 0.5
0.5 to 1.9Greater than 1.9
Nonpolar covalentPolar covalent
Ions form
1-20
Polar and Nonpolar Covalent Bonds
• an example of a polar covalent bond is that of H-Cl• the difference in electronegativity between Cl and H is
3.0 - 2.1 = 0.9• we show polarity by using the symbols d+ and d-, or
by using an arrow with the arrowhead pointing toward the negative end and a plus sign on the tail of the arrow at the positive end
H Cl+ -
H Cl
1-21
Polar Covalent Bonds
Bond dipole moment (m): • a measure of the polarity of a covalent bond• the product of the charge on either atom of a polar
bond times the distance between the nuclei• average bond dipole moments of selected covalent
bonds
H-OH-NH-C
H-S C-I
C-FC-ClC-Br C-N
-
C-O
C=N
C=O
-
Bond Dipole (D) Bond
1.4
1.51.51.41.2
Bond
Bond Dipole (D)
1.30.3 0.7
0.20.7
2.3
3.5
Bond Dipole (D)Bond
1-22
Polar and Nonpolar Molecules
To determine if a molecule is polar, we need to determine • if the molecule has polar bonds• the arrangement of these bonds in space
Molecular dipole moment (): the vector sum of the individual bond dipole moments in a molecule• reported in debyes (D)
1-23
Polar and Nonpolar Molecules
these molecules have polar bonds, but each has a zero dipole moment
O C O
Carbon dioxide = 0 D
B
F
F
F
Boron trifluoride = 0 D
C
Cl
ClClCl
Carbon tetrachloride = 0 D
1-24
Polar and Nonpolar Molecules
these molecules have polar bonds and are polar molecules
N
HH
H
O
H H
Water = 1.85D
Ammonia = 1.47D
directionof dipolemoment
directionof dipolemoment
1-25
Polar and Nonpolar Molecules
• formaldehyde has polar bonds and is a polar molecule
Formaldehyde = 2.33 D
directionof dipolemoment H H
C
O
1-26
Formation of Ions
A rough guideline: • ions will form if the difference in electronegativity
between interacting atoms is 1.9 or greater• example: sodium (EN 0.9) and fluorine (EN 4.0)• we use a single-headed (barbed) curved arrow to show
the transfer of one electron from Na to F
• in forming Na+F-, the single 3s electron from Na is transferred to the partially filled valence shell of F
Na + F Na+ F -
• •• •••
••
• •••
••
+ F-(1s22s22p6)Na+(1s22s22p6)F(1s22s22p5)+Na(1s22s22p63s1)
1-27
Formal Charge
Formal charge: the charge on an atom in a molecule or a polyatomic ion
To derive formal charge1. write a correct Lewis structure for the molecule or ion
2. assign each atom all its unshared (nonbonding) electrons and one-half its shared (bonding) electrons
3. compare this number with the number of valence electrons in the neutral, unbonded atom
Number of valence electrons in the neutral, unbonded atom
All unsharedelectrons
One half of all shared electrons
+Formalcharge
=
1-28
Formal Charge
Example: Draw Lewis structures, and show which atom in each bears the formal charge
NH2- HCO3
- CO32-
CH3NH3+ HCOO- CH3COO-
(b) (c)
(d) (e) (f)
(a)
C O
H
HC
H
H
H
C N
1-29
Exceptions to the Octet Rule
Molecules containing atoms of Group 3A elements, particularly boron and aluminum
Aluminum chloride
:
:
:
F B
F
F
Cl Al
Cl
Cl
6 electrons in the valence shells of boron
and aluminum
Boron trifluoride
: :
: :
: :
::
::
:
::
::
1-30
Exceptions to the Octet Rule
Atoms of third-period elements have 3d orbitals and may expand their valence shells to contain more than 8 electrons• Sulfur, phosphorus may have up to 10
1-31
Resonance
For many molecules and ions, no single Lewis structure provides a truly accurate representation
Ethanoate ion(acetate ion)
C
O
O
H3CC
O
O
H3C
-
-
and
1-32
Resonance
Linus Pauling - 1930s• many molecules and ions are best described by
writing two or more Lewis structures• individual Lewis structures are called contributing
structures• connect individual contributing structures by double-
headed (resonance) arrows• the molecule or ion is a hybrid of the various
contributing structures
1-33
Resonance
Examples: equivalent contributing structures
Acetate ion(equivalent contributing
structures)
Nitrite ion(equivalent contributing
structures)
::
::
: : :
N
O -
O
: N
O
O -
:
: : :
:
:
: :
: :
C
O -
O
CH3 C
O
O -
CH3
: :
:
:
1-34
Resonance
Curved arrow: a symbol used to show the redistribution of valence electrons
In using curved arrows, there are only two allowed types of electron redistribution:• from a bond to an adjacent atom• from an atom to an adjacent bond
Electron pushing is a survival skill in organic chemistry• learn it well!
1-35
Resonance
All contributing structures must1. have the same number of valence electrons
2. obey the rules of covalent bonding• no more than 2 electrons in the valence shell of H • no more than 8 electrons in the valence shell of a 2nd
period element• 3rd period elements, such as P and S, may have up to
12 electrons in their valence shells
3. differ only in distribution of valence electrons; the position of all nuclei must be the same
4. have the same number of paired and unpaired electrons
1-36
Resonance The carbonate ion, for example
• a hybrid of three equivalent contributing structures• the negative charge is distributed equally among the
three oxygens
1-37
Resonance
Preference 1: filled valence shells• structures in which all atoms have filled valence shells
contribute more than those with one or more unfilled valence shells
••
••••
Greater contribution; both carbon and oxygen have complete valence shells
Lesser contribution;carbon has only 6 electrons in its valence shell
+ +CCH3 OCH3 O
H
HC
H
H
1-38
Resonance
Preference 2: maximum number of covalent bonds• structures with a greater number of covalent bonds
contribute more than those with fewer covalent bonds
••
••••
Greater contribution(8 covalent bonds)
Lesser contribution(7 covalent bonds)
+ +CCH3 OCH3 O
H
HC
H
H
1-39
Resonance
Preference 3: least separation of unlike charge• structures with separation of unlike charges contribute
less than those with no charge separation
Lesser contribution(separation of unlike
charges)
CH3-C-CH3 CH3-C-CH3
Greater contribution(no separation of unlike charges)
O -O
::
:::
1-40
Resonance
Preference 4: negative charge on the more electronegative atom • structures that carry a negative charge on the more
electronegative atom contribute more than those with the negative charge on the less electronegative atom
CH3CH3H3CCH3H3CC
O
C
O
H3CC
O
(b)Greater
contribution
(c)Should notbe drawn
(a)Lesser
contribution
(1) (2)
1-41
Structural Formulas
Lewis structures indicate atom connectivity• Do not show 3-D structure, but it is there
Drawing Lewis structures of large molecules can get cumbersome• Use shortcuts to simplify things• Rely on normal bonding properties of atoms
Types include• Condensed structural formulas• Line-angle formulas
All subsets of molecular formula
1-42
Condensed Structural formulas
Based on typical bonding patterns• Carbon makes 4 bonds in stable molecules• Oxygen makes 2 bonds and has 2 lone pairs• Nitrogen makes 3 bonds and has 1 lone pair
Show molecule by “groups” centered on carbon
C C
H
H
H
H
H
H
C
H
HH
H
CH4
CH3CH3
C C
H
H
H
H
C
H
H
H
H
CH3CH2CH3
or
(CH3)2CH2
C
C
C C
H H
H
HH
H
H
H
HH
CH3CHCH3
CH3
CH3CH(CH3)CH3
(CH3)3CH
1-43
Which is the correct Lewis structure?
(CH3)2CH(CH2)2CH(CH3)2
C
C CH
H
H
H H
C C
C
H
H
H
H
H
H
H H
H
CC
C
C
H
H
H
H
H
H
H
HH
C
C
H
H
C CH
H
H
HH
HH
CC C
C
C
H
H
H
H H
H
HH
H
H
C
C
H
H
H
HH
H
CC C
C
C
H
H
H
H H
H
HH
H
C
C
C
H
H
H
HH
H
HH
H
AB
CD
1-44
Can add heteroatoms (N, O, halogens) and multiple bonds
CH3CH2OH
CH3CH2NHCH3
CH3CHClCH3
CH3CH=CHCH3
1-45
Line-Angle structures
Lines to represent bonds • “skeletal structures”
Carbons are not drawn, but occur at• Intersections• Vertices (ends of lines)
Hydrogens are not usually shown• assumed to be as many as needed to fulfill carbon
valency
C
C CH
H
H
H H
C C
C
H
H
H
H
H
H
H H
H
CH3(CH2)4CH3
1-46
A. 0 B. 1 C. 2 D. 3 E. 4
Heteroatoms are shown explicitly WITH Hydrogens
OH
O
NH2
How many hydrogens are present at marked position?
*
*
1-47
Combined structures are allowed BUT…
…if you show C atoms, you MUST show the 4 bonds to it (including H)
H3C OH
H2CO
NH2
OK`
OH
O
NH2
O
CH(CH3)2
NOT OK
O
CCH3
CH3
C C C C
(C14H30)
Trivalent carbon
1-48
Find the acceptable structure
C
Cl2C
NH2
Cl
C
ONH2
Cl
C
ONH2
Cl
C
ONH2
Cl
CH3
ONH2
Cl
CH3
A B
C D
E
1-49
Trivalent carbon can be possible for ions and radicals
O
CCH3
CH3
O
CCH3
CH3
+
_
one missing electron around carbon(3 net electrons)
one extra electron around carbon(5 net electrons)
One for each bond (3) – no more (electron deficient)
3 from the bonds, must have a non-bonded lone pair (note that it has an octet even!)
Pentavalent carbon (5 bonds to C) is NEVER acceptable in this class
1-50
Acids and Bases
1-51
Arrhenius Acids and Bases
In 1884, Svante Arrhenius proposed these definitions • acid: a substance that produces H3O+ ions aqueous
solution• base: a substance that produces OH- ions in aqueous
solution• this definition of an acid is a slight modification of the
original Arrhenius definition, which was that an acid produces H+ in aqueous solution
• today we know that H+ reacts immediately with a water molecule to give a hydronium ion
H+(aq) + H2O(l) H3O+(aq)Hydronium ion
1-52
Brønsted-Lowry Definitions
Acid: a proton donor Base: a proton acceptor
+
Proton donor
Protonacceptor
-O H
H
OH
H
O HH + +O H H
+
Protonacceptor
Protondonor
+O H
H
OH
H
N H
H
H
H
H
H + +N H H
: ::
: ::
::
: :
::
1-53
Conjugate Acids & Bases
• conjugate base: the species formed from an acid when it donates a proton to a base
• conjugate acid: the species formed from a base when it accepts a proton from an acid
• acid-base reaction: a proton-transfer reaction• conjugate acid-base pair: any pair of molecules or ions
that can be interconverted by transfer of a proton
HCl(aq) H2O(l) Cl-(aq) H3O+(aq)+ +WaterHydrogen
chlorideHydronium
ionChloride
ion(base)(acid) (conjugate
acid of H2O)(conjugate
base of HCl)
conjugate acid-base pair
conjugate acid-base pair
1-54
Conjugate Acids & Bases
• Brønsted-Lowry definitions do not require water as a reactant
• consider the following reaction between acetic acid and ammonia
NH4+CH3COOH CH3COO-
NH3+ +
Acetic acid Ammonia
(acid)
conjugate acid-base pair
Acetate ion
Ammoniumion
(base) (conjugate baseacetic acid)
(conjugate acidof ammonia)
conjugate acid-base pair
1-55
Conjugate Acids & Bases
• we can use curved arrows to show the flow of electrons in an acid-base reaction
CH3-C-OO
H N HH
H CH3-C-O -
OH-N-H
H
H+ +
Acetic acid(proton donor)
Acetate ion
:
:: ::
:: :: :
Ammonia(proton acceptor)
Ammoniumion
1-56
Conjugate Acids & Bases
Many organic molecules have two or more sites that can act as proton acceptors in these molecules, the favored site of protonation is
the one in which the charge is more delocalized question: which oxygen of a acetic acid is protonated?
CH3-C-O-H
O+ H2SO4 CH3-C-O-H
HO+
CH3-C-O-HH
O+
+ HSO4-or
A(protonation
on the carbonyl oxygen)
B(protonation
on thehydroxyl oxygen)
acetic acid
1-57
Conjugate Acids & Bases
for protonation on the carbonyl oxygen, we can write three contributing structures
two place the positive charge on oxygen, one places it on carbon
A-1 and A-3 make the greater contribution because all atoms have complete octets
the positive charge is delocalized over three atoms with the greater share on the two equivalent oxygens
++
CH3-C-O-H
OH
CH3-C-O-H
OH
CH3-C
OH
++CH3-C-O-H
O
CH3-C=O-H
OH
CH3-C=O-H
OH+
A-1(C and O have
complete octets)
A-2(C has incomplete
octet)
A-3(C and O have
complete octets)
1-58
Conjugate Acids & Bases
for protonation on the hydroxyl oxygen, we can write two contributing structures
B-2 makes only a minor contribution because of charge separation and adjacent positive charges
therefore, we conclude that protonation of a carboxylic acid occurs preferentially on the carbonyl oxygen
CH3-C-O-H
O
H
+
B-1
CH3-C-O-H
O
H+
+
B-2(charge separation and
adjacent positive charges)
1-59
Conjugate Acids & Bases
Does proton transfer to acetamide occur preferentially on the oxygen or the nitrogen?
CH3-C-N-H
H
O+ HCl
+
CH3-C-N-H
OH
H
H
CH3-C-N-H
H
O+
+ Cl -or
A(protonation
on theamide oxygen)
B(protonation
on theamide nitrogen)
Acetamide(an amide)
1-60
Acids & Base Strengths
The strength of an acid is expressed by an equilibrium constant the acid dissociation of acetic acid is given by the
following equation
H3O+
+CH3CO-
OH2O+CH3COH
O
Acetic acid Water Acetateion
Hydroniumion
1-61
Weak Acids and Bases
We can write an equilibrium expression for the dissociation of any uncharged acid, HA, as:
water is a solvent and its concentration is a constant equal to approximately 55.5 mol/L
we can combine these constants to give a new constant, Ka, called an acid dissociation constant
HA + H2O
[H3O+][A
-]
[HA][H2O]Keq
A- + H3O
+
=
Keq[H2O]Ka[H3O+][A-]
[HA]= =
1-62HI I-
HBr Br-
HCl Cl-
H2SO4 HSO4-
H2OH3O+
H3PO4 H2PO4-
C6H5COOH C6H5COO-
CH3COOH CH3COO-
H2CO3 HCO3-
H2S HS-
NH3NH4+
C6H5OH C6H5O-
HCO3-
CO32-
CH3NH2CH3NH3+
H2O HO-
CH3CH2OH CH3CH2O-
HC CH HC C-H2 H-NH3 NH2
-CH2=CH2 CH2=CH
-CH3CH3 CH3CH2
-Acid Formula pKa Conjugate BaseEthane
Ammonia
EthanolWater
Bicarbonate ionPhenol
Ammonium ion
Carbonic acidAcetic acid
3525
Benzoic acidPhosphoric acid
Sulfuric acidHydrogen chlorideHydrogen bromideHydrogen iodide
51
38
10.33
15.715.9
4.766.36
9.24
9.95
-5.2-7
-9-8
4.19
2.1-1.74Hydronium ion
Strongerconjugate
base
Weakerconjugate
base
Weaker acid
Stronger acid
Methylammonium ion 10.64
Hydrogen sulfide 7.04
AcetyleneHydrogen
Ethylene 44
1-63
Acid-Base Equilibria
Equilibrium favors reaction of the stronger acid and stronger base to give the weaker acid and weaker base
CH3COOH NH3 CH3COO-
NH4++ +
Ammonia(stronger base)
Acetate ion(weaker base)
Acetic acidpKa 4.76
(stronger acid)
Ammonium ionpKa 9.24
(weaker acid)
pKeq = 4.76 - 9.24 = -4.48
Keq = 3.0 x 104
1-64
Acid-Base Equilibria
Consider the reaction between acetic acid and sodium bicarbonate we can write the equilibrium as a net ionic equation we omit Na+ because it does not undergo any
chemical change in the reaction
equilibrium lies to the right carbonic acid forms, which then decomposes to
carbon dioxide and water
CH3COH
OHCO3
- CH3CO-
O
H2CO3
Bicarbonate ion Acetate ionAcetic acidpKa 4.76
(stronger acid)
Carbonic acidpKa 6.36
(weaker acid)
+ +
1-65
Lewis Acids and Bases
Lewis acid: any molecule of ion that can form a new covalent bond by accepting a pair of electrons
Lewis base: any molecule of ion that can form a new covalent bond by donating a pair of electrons
Lewis base
Lewis acid
+ BA new covalent bondformed in this Lewisacid-base reaction
:+-
A B
1-66
Lewis Acids and Bases
• examples
:
: :
Br -CH3-C C-CH3
H H
H
CH3-C C-CH3
H H
H Br
Bromideion
sec-Butyl cation(a carbocation)
2-Bromobutane
++
: :
::
+ -
Diethyl ether(a Lewis base)
Boron trifluoride (a Lewis acid)
O
CH3CH2
CH3CH2
F
F
O
CH3CH2
CH3CH2
B-F
F
F
+ B
A BF3-ether complex
F:: :
1-67
Lewis Acids include Bronsted Acids• A proton (H+) is a Lewis acid
• Proton donors therefore “accept a pair of electrons” to form a new covalent bond – between the base and proton
• A covalent bond is also broken but that doesn’t matter• The definition of Lewis acid doesn’t say anything about not
breaking bonds • All that matters is that it can ACCEPT a pair of electrons
• Hydroxide (OH) is clearly a Lewis base• Other things are bases in Lewis approach (ammonia, halide
ions)
• Most neutral molecules are acids and bases• However, they can be very different in acid/base strength NH2
+ H3O+ :NH3 + H2O NH4+ + OH
pKa(NH3) = 38 pKa(NH4+) = 9