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Chapter 15
Organic Chemistry
Life is chemistry.
Source: Science, Jan 9, 2009, “On the origins of life on earth”
Organic chemistry is enough to
drive one mad. - Friedrich Wöhler
What is an organic compound?
Organic compound – contains carbon, nearly always bonded to other C and H, and
often other elements
Vitalism
A major misconception that stifled organic chemistry research in early 19th Century.
Resulted in the basic distinction between organic and inorganic substances
An unobservable spiritual energy existed within organic compounds of living
things, making them impossible to synthesize and fundamentally different
from inorganic compounds (compounds of the “mineral world” – mostly, what
we have studied so far)
Wöhler’s experiment changed that
Vitalism: Organic molecules were
yhought to arise spontaneously
(Spontaneous Generation) and could
not be synthesized from inorganics
Urea synthesized from ammonium cyanate.
(2 compounds – same molecular formula)
Reading for today: Did life originally arise from inorganic chemicals?
Classes of organic molecules
I. Hydrocarbons – simplest type of organic compound
-functional groups & reactivity
-polymers
Classes of organic molecules
II. Biomolecules – natural polymers
-polysaccharides, proteins, nucleic acids
Section 15.1: What’s so special about Carbon?
Atomic properties of carbon (and bonding behavior) make it special.
Structural complexity of organic compounds
C’s location in the periodic table tells you a lot
Always bonds covalently – moderate
EN makes formation of C ions
energetically impossible under ordinary
conditions.
Why? – REVIEW
Ionization energy (IE) – Chap 8
Energy required for the complete
removal of 1 mole of e- from 1 mole of
gaseous atoms or ions (E to overcome
attraction between protons & e-)
Structural complexity of organic compounds
C is small and forms 4 covalent bonds
Ionization energy (IE) – Chap 8
As size decreases, more E to remove
an e-
Ionization energy (IE) – Energy required to get C4+ ion = IE1 + IE2 + IE3 + IE4
A lot of energy to remove an e- ……..and to add e-’s  Electron affinity (EA):
The energy change accompanying the addition of
1 mole of e-’s to 1 mole of gaseous atoms or ions.
Energy is required to get C4- ion = EA1 + EA2 + EA3 + EA4
EA1 is negative (exothermic): Energy released
EA2 – EA4 are positive (endothermic): Energy required
Structural complexity of organic compounds
Carbon has the ability to catenate – form chains of atoms (= large, complex molecules)
Due to the sp3 hybridization:
C forms 4 bonds in nearly all
of its compounds
C forms short, strong bonds:
Small size allows close approach
of another atom
Structural complexity of organic compounds
Carbon easily forms double and triple bonds: C – C bond is short enough to allow
side-to-side overlap
Double bond
Triple bond
Structural complexity of organic compounds
Double and triple bonds:
Restricts rotation
=
MORE variety
Structural complexity of organic compounds
C’s location in the periodic table tells you a lot – Periodic Trends
So why don’t Si, Ge and Sn also form organic compounds? In same Group 4A as C.
(1) Atomic size and bond strength
i.e. C – C bonds = 347 kJ/mol
Si – Si bonds = 226 kJ/mol
(2) ∆Hreaction
i.e. C – C (347), C – O (358)
Si – Si (226), Si – O (368)
(3) Orbitals available for reaction
i.e. C has s and p orbitals
Si has s, p, and d orbitals
d orbitals can be attacked by lone
e- pairs of incoming reactants
Ethane (CH3-CH3): Stable in water
and air
Disilane (SiH3-SiH3): Breaks down in
water, spontaneous ignition in air
Chemical diversity of organic compounds
CRC Handbook of Physics and Chemistry - # of C-based compounds dwarfs the # of
compounds formed from all of the other elements combined
Chemical diversity also a result of atomic and bonding behavior of carbon.
Bonding to heteroatoms: Organic compounds contain
atoms other than C and H (also N, O, S, P and halogens)
Example:
23 organic molecules
4 singley bonded C
1O
Filled in with H
Chemical diversity of organic compounds
Electron density and reactivity
Most chemical reactions start (and new bonds form) when a region of high e- density
on one molecule meets a region of low e- density on another
Regions of high e- density can be due to:
(1) Multiple bonds
(2) Partial charges
(3) Lone pairs
4 bonds commonly found in organic molecules:
C – C: Generally, unreactive – EN values equal and bond is nonpolar
C – H: Largely unreactive – EN values close (C = 2.5, H = 2.1)
– bond is short (strong)  C and H are both small atoms
C – O: Reactive – Highly polar (∆ EN = 1.0)  O end of bond is e- rich
Bonds to other heteroatoms (S, P, Br): Reactive – bonds longer (S, P, Br large
relative to H)
Chemical diversity of organic compounds
Functional Group – a specific combination of bonded atoms that reacts in a
characteristic way, no matter what organic molecule it occurs in
In fact, reactions in organic molecule nearly always take place at functional groups.
Example:
Structure of amino acids
20 amino acids differ
only by functional group
Section 15.2: Hydrocarbons
Organic Molecule-Animal Analogy for Hydrocarbons:
• C – C bonds form the skeleton
• H atoms are the skin covering the skeleton
• Functional groups are limbs protruding from body ready to “grab” (react with) reactants
Hydrocarbons – a large group of organic compounds containing only H and C atoms
Example: Natural gas and gasoline are hydrocarbon mixtures
Section 15.2: Hydrocarbons
Carbon skeletons – What different possible arrangements exist for C atoms?
For example: If you have two carbon atoms, there is one possible arrangement
C–C
As the number of carbon atoms increases, the number of arrangements increases.
Section 15.2: Hydrocarbons
Practice Drawing Hydrocarbons
Purpose: Get a sense of the number of possibilities for a given formula (i.e. C6H14)
Steps:
#1: Are there single, double, or triple bonds? How many of each?
#2: Figure out the arrangement of C atoms
#3: Add the H skeleton
(1) Six C atoms, no multiple bonds, no rings
(2) Four C atoms, one double bond, no rings
(3) Four C atoms, no multiple bonds, one ring
Section 15.2: Hydrocarbons
Hydrocarbon classification – 4 main groups:
(1) Alkanes – single bonds
(2) Alkenes – double bonds
(3) Alkynes – triple bonds
(4) Aromatic Hydrocarbons - rings
Alkanes – CnH2n+2
Each carbon is sp3 hybridized
Each C is bonded to the maximum number of other atoms – saturated hydrocarbons
Naming: Each chain, branch or ring has a name based on the number of carbons
Prefix + root + suffix
Root: # of carbon atoms in the longest continuous chain in the molecule (Table 15.1)
Suffix: type of organic compound (identifies key functional group)  -ane for alkanes
Prefix: groups attached to the main chain
Example: Table 15.2
Section 15.2: Hydrocarbons
Different ways to depict molecules
Section 15.2: Hydrocarbons
Cyclic Hydrocarbons – Rings
Cycloalkanes – 2 H’s are lost when ring forms from straight chain – CnH2n
Section 15.2: Hydrocarbons
Isomers – two or more compounds with the same molecular formula but
with different properties
Constitutional Isomers – different arrangements of bonded atoms
Section 15.2: Hydrocarbons
Physical Properties of Alkanes
Why do we see this trend in boiling point?
Section 15.2: Hydrocarbons
Chiral Molecules and Optical Isomerism
Optical isomers – molecules are mirror images of each other
Often indicated with L and D:
L-alanine
D-alanine
Most naturally proteins are composed of L-amino acids: L-leucine, L-glutamine.
Opposite for naturally occuring carbohydrates: D-glucose metabolized, L-glucose excluded
Section 15.2: Hydrocarbons
Alkenes – CnH2n
Each carbon is sp2 hybridized
Each C is bonded to fewer than max # of other atoms – unsaturated hydrocarbons
Naming: Each chain, branch or ring has a name based on the number of carbons
Prefix + root + suffix
Root: # of carbon atoms in the chain that contains the double bonds (even if not longest)
Suffix: type of organic compound (identifies key functional group)  -ene for alkenes
Prefix: groups attached to the main chain
Name these alkenes:
Section 15.2: Hydrocarbons
Geometric Isomers: Cis-Trans Isomerism – because π bonds restrict rotation
Section 15.2: Hydrocarbons
Alkynes – CnH2n-2
Each carbon is sp hybridized
Alkanes
1 σ bond
Alkenes
1 σ bond
1 π bond
Alkynes
1 σ bond
2 π bonds
Section 15.2: Hydrocarbons
Aromatic hydrocarbons – one or more rings of 6 carbons atoms
Benzene is simplest example
Naming = attached groups + -benzene suffix
Section 15.3: Organic Reaction Types
Functional Group – a specific combination of bonded atoms that reacts in a
characteristic way, no matter what organic molecule it occurs in
Notation: R – CH2 – Br where R is an alkyl group (a saturated hydrocarbon chain)
Three main reaction types:
1) Addition reactions: unsaturated reactant  saturated product
Generic reaction
Example: Ethylene
Characteristics:
• common for double and triple bonded C’s, and C = O bonds
• π bonds break, σ bonds remain
• reaction occurs b/c it is energetically favorable
Show why is this reaction energetically favorable
Section 15.3: Organic Reaction Types
2) Elimination reactions: opposite of addition reactions
saturated reactant  saturated product
Generic reaction
Example
Characteristics:
• Typically eliminates:
2 halogens (i.e. Cl2), H and halogen (i.e. HBr), or H and –OH group (i.e. H2O)
• Driving force of this reaction is formation of small, stable molecules
Addition Reaction example
Reactants
Bond Energy
Products
Bond Energy
2693 kJ
3098 kJ
4410 kJ
4373 kJ
Elimination Reaction example
What is wrong with this picture?
Thermodynamics in a Nutshell
G – Gibbs free energy – in chemistry, the “force” that causes chemical reactions
– can tell us whether or not a reaction will occur
H – enthalpy – keeps track of the quantity of energy
– in chemical reactions, it is the energy change during a
reaction (∆Hreaction, ∆Hlattice)
You can ask: Will the reaction occur spontaneously?
∆H is negative  exothermic (energy lost) = more stable = YES
∆H is positive  endothermic (energy required) = less stable = NO
Addition Reaction example
Reactants
Bond Energy
Products
Bond Energy
2693 kJ
3098 kJ
4410 kJ
4373 kJ
Elimination Reaction example
Thermodynamics in a Nutshell
G – Gibbs free energy – in chemistry, the “force” that causes chemical reactions
– can tell us whether or not a reaction will occur
S – entropy – keeps track of the distribution of energy in a system
Rule: Energy becomes distributed more uniformly (more disordered) with time
Hot
Cold
Heat flow
Dissolution (Chap12)
Diffusion
Proton Pump
(non-spontaneous)
Summary: Thermodynamics in a Nutshell
G – Gibbs free energy – in chemistry, the “force” that causes chemical reactions
– can tell us whether or not a reaction will occur
H – enthalpy – keeps track of the quantity of energy
– in chemical reactions, it is the energy change during a
reaction (∆Hreaction, ∆Hlattice)
You can ask: Will the reaction occur spontaneously?
∆H is negative  exothermic (energy loss as heat) = more stable = YES
∆H is positive  endothermic (energy needs to be added) = less stable = NO
S – entropy – keeps track of the distribution of energy in a system
– energy becomes distributed more uniformly (more disordered) with time
You can ask: Will the reaction occur spontaneously?
uniformity/disorder increases  YES
uniformity/disorder decreases  NO
Addition Reaction example
Reactants
Bond Energy
Products
Bond Energy
2693 kJ
3098 kJ
4410 kJ
4373 kJ
Elimination Reaction example
Section 15.3: Organic Reaction Types
3) Substitution reactions:
Generic reaction
Example
Characteristics:
• C involved in bonding can be saturated or unsaturated (involved in double, triple bonds)
Section 15.3: Redox Process in Organic Reactions
Oxidation-reduction reactions in O-chem:
Do NOT monitor change in O.N. of various C atoms in a compound.
Rather, note movement of e- density around C based on # of more/less EN atoms
More EN atom takes e- density from C (oxidation)
Example: C – C bonds replaced with C – O bonds
2 CH3-CH3 + 7 O2  4 CO2 + 6 H2O
Nature’s Redox:
Respiration
(Oxidation)
Less EN atom gives e- density to C (reduction)
Example: C – H bonds replaces a C – O bond
CH3O  CH4
Photosynthesis
(Reduction)
In O Chem: Focus is usually on the organic reactant only.
Oxidation: C forms more bonds to O, Br, F, etc or fewer to H
Reduction: C forms fewer bonds to O, Br, F, etc or more bonds to H
Section 15.4: Properties & Reactivities of Functional Groups
The distribution of e- density in the functional group affects the reactivity
(1) Functional groups with single bonds only
alcohols, haloalkanes, amines
(2) Functional groups with double bonds
alkenes, carbonyl group (aldehydes & ketones)
(3) Functional groups with both single and double bonds
carboxylic acid, ester, amide
(4) Functional groups with triple bonds
nitrile, alkynes
Section 15.5: Monomers & Polymers – Synthetic Macromolecules
Polymers – many monomer units bonded together
Section 15.5: Monomers & Polymers – Synthetic Macromolecules
Petroleum-based products – there will be a shortage of raw materials soon
bisphenol A (BPA)
- used in synthesizing DGEBA,
a building block for an epoxy resin
Section 15.5: Monomers & Polymers – Synthetic Macromolecules
Addition polymers – as each monomer adds to the chain, it forms a new reactive site.
Section 15.6: Monomers & Polymers – Biological Macromolecules
Section 15.6: Polysaccharides
Glucose is a monosaccharide – alcohol and aldehyde groups react to make cyclic forms
Polysaccharide chains formed from cyclic forms that undergo dehydration reactions.
Different disaccharides formed from different monosaccharides:
sucrose (table sugar): glucose (C-1) + fructose (C-2)
lactose (milk sugar): glucose (C-1) + galactose (C-4)
maltose (beer): glucose (C-1) + glucose (C-4)
Section 15.6: Polysaccharides
3 main groups of polysaccharides:
Cellulose – most abundant organic chemical on earth, structural function (plant cell
walls), long chains of glucose, humans cannot digest this (cows, sheep, termites)
(C6H10O5)n
Starch – energy storage in plants (amylose and amylopectin)
Glycogen – energy storage in animals
Section 15.6: Amino Acids and Proteins
Section 15.6: Amino Acids and Proteins
tyrosine
valine
leucine
tyrosine
Ionic bonds
Hydrogen bonds
Disulfide bonds
(covalent)
+
hydrophobic
interactions
between –CH3
Section 15.6: Nucleic Acids, DNA, and RNA
Pyrimidines:
Thymine (T) [Uracil (U)], cytosine (C)
Purines:
Guanine (G), Adenine (A)
A – T(U) , G – C