welcome to organic chemistry
DESCRIPTION
Chemistry of Carbon CompoundsTRANSCRIPT
Organic Chemistry• Jons Jakob Berzelius – (1779 – 1848)
– organic – contained the vital force and were obtained from living things
– Inorganic – do not contain the vital force and were from non living things
Scientists believed that is was impossible to create compounds that contained the vital force.
• Problem arose when Friedrich Wohler produced urea
NH4+NCO- NH2 C NH2
OHeat
ammonium cyanate
Urea
Review
• Atomic number• Atomic mass• Isotope• Atomic weight – amu = 1/12 mass of C12 • Molecular weight – sum of atomic
weights of all atoms present
Review
• Waves and Particles– Particles have wave properties – waves have particle properties– What do you do when you can’t see any more?????????? …………….. You rely on mathematics
• Schrodinger equation tells us the most likely energy and electron possesses
• The closer the shell is to the nucleus the lower the energy
• Practice an electronic configuration
• Re-visit this section in Gen. Chem. if needed.
Covalent Bond
Review
• Chemical Bonding– Ionic– Covalent– Polar Covalent
Electronegativity is my new Organic Chemistry Friend – I know everything and anything about it……………….
Electronegativity
• Attraction an atom has for a pair of shared electrons
• Dipole – molecule with a positive and negative end
reported in debye (D)
Electrostatic Potential Maps
Lewis Structures
• Count valence electron
• Place least electronegative in center
• Draw dashes to form bonds
• Fill valence on outer atoms
• Put extra electrons in pairs on center atom
• Make sure center atom has octet
• Form double or triple bonds if needed
Lewis Structures
Basic Shapes
# of atoms bonded # of lone pair electrons
to center on center atom Shape
4 0Tetrahedral
3 0 Triangular
3 1 Pyramidal
2 0 Linear
2 1 or 2 Bent
Lewis Structures
Practice the Following
CO2
CH2O
SiCl4
PCl3 SO2
Formal Charge
FC = # valence – ( lone pairs + ½ bonded e- )
Practice
NO2- NH4+ CH2O
When two structures are possible, the one with more negative
numbers on the electronegative atoms is correct
CHN , CS2
Atomic Orbitals
Traveling wave – wave that travels through space (light – water)
Standing wave – wave that is confined to a limited space (guitar string)
Electrons acts like a standing wave
but in three dimensions
Atomic Orbital Theory
• Wave equation predicts probability of e- location
Atomic Orbital Theory
• Orbitals are actually solutions to the Wave Equation
Degenerate Orbitals• Simultaneous Roots (equal in energy)
"Many people tell you that an expert is someone who knows a great deal about his subject. To this I would object that no one can ever know very much about any subject. I would much prefer the following definition: an expert is someone who knows some of the worst mistakes of his subject, and how to avoid them."
- Werner Heisenberg
Molecular Orbital Theory
• An overlap of atomic orbitals yields … – Molecular orbitals
Bonding• Bond distance occurs at an energy
minimum• Direct overlap of orbitals results in a (sigma)
bond.– Cylindrically symmetrical– e- attracted to both nuclei
MO Theory
• Overlap of wave equation shows– reinforcement and…– cancellation
MO Theory• Molecular Orbital Theory predicts a
bonding orbital and a or antibonding orbital.
• Results: one 1s orbital and one 1s
orbital.
2p overlapMolecular Orbitals
Atomic Orbitals Atomic Orbitals
2p overlap
Bond order = ½ (# e- in bond MOs - # e- in anti bond MOs)
e-
e-
e-
e-e-
e-
1s
2s
2p
E
EN
RG
Y
Carbon Atom
CH
HH
H
1sE
EN
RG
Y
Carbon sp3
sp3
C C
H
H
H
H
1sE
EN
RG
Y
Carbon sp2
sp2
2p
C C HH
2p
1sE
EN
RG
Y
Carbon sp
sp
Bonding in Methane
• 4 Covalent bonds– All identical bonds (4 bonds)– Bond angles of 109o – Non-polar molecule since e- density uniform
Bonding in Ethane - CH3CH3
• Ethane has sp3 hybridization about each C orbitals– Nearly tetrahedral shape about each C
(All single bonds in organic chem are bonds!)
Bonding in Ethene - CH2CH2
• Overlap of the unhybridized, parallel orbitals yields a bond (along with the bonds you would expect).
Bonding in Ethyne - CHCH
• Each C forms a bond and 2 bonds
Representation of Acetylene
• Shape is referred to as “linear” (180o)
Bonding in Cations and Anions
• Cations have too few electrons• Anions have extra electrons• Unpaired e- are unhybridized
• Consider the methyl cation, CH3+
CH3.
• Consider the Methyl Radical, CH3.
CH3-
• Consider the Methyl Anion, CH3-
sp3 Water H2O
e-
e-
e-
e-e-
e- H O
H
e-
e-
1s
2s
2p
EN
RG
Y
Oxygen Atom
1sE
EN
RG
Y
Water sp3
sp3
Ammonia• Ammonia, NH3 has one pair of “unbonded” e-
QuickTime™ and aTIFF (Uncompressed) decompressor
are needed to see this picture.
1s
2s
2p
EN
RG
Y
Nitrogen Atom
1sE
EN
RG
Y
Ammonia
sp3
Ammonium Ion
• Ammonium ion, NH4+, has 8 valence e-
• Final shape is tetrahedron
Bonding - HX
• Hydrogen halides– F is in second series, Cl in third, etc.– Longer bonds are weaker bonds
Determining Bond Hybridization
• Count e- pairs as one “thing”
• Count a bond as one “thing”– Double bonds count as only one thing– Triple bonds count as one only thing
Acid / Base
Acid• Anything that increases the hydronium ion
concentration• Anything that donates a hydrogen• Anything that accepts a pair of electrons
Base • Anything that increases the hydroxide ion
concentration• Anything that accepts a hydrogen• Anything the donates a pair of electrons
Bronsted-Lowry Definitions
• Proton donors/acceptors– HCl + H2O H3O+ + Cl-
acid base acid base
conjugateconjugate
acid base
Acid / Base
What is?????
pH
Ka
pKa
Acids and Bases
• pH is a measure of [H+]
• The more loosely held H+ , the more acidic
Acid Dissociation Constants
• Weak acids don’t dissociate completely!
• We measure dissociation using a Ka.
HA H+ + A-
Ka = [H+] [A- ] [HA]
(a weak acid)
Acid Dissociation Constants
• Stronger acids have a LARGER Ka
• Weaker acids have a smaller Ka
HA H+ + A-
Ka = [H+] [A- ] [HA]
(a weak acid)
• Name Formula Ka
phosphoric acid H3PO4 7.5 x 10-3
hydroflouric acid HF 3.5 x 10-4 formic acid HCOOH 1.8 x 10-4
lactic acid CH3CHOHCOOH 1.4 x 10-4
acetic acid CH3COOH 1.8 x 10-5
carbonic acidH2CO3 4.3 x 10-7
boric acid H3BO3 7.3 x 10-10
ammonium ion NH4+ 5.6 x 10-10
phenol C6H5OH 1.3 x 10-11
bicarbonate ion HCO3- 2.2 x 10-13
Acid Dissociation Constants
pKa
• pKa = - log (Ka)– The larger the pKa, the weaker the acid
– The smaller the pKa, the stronger the acid
Strength of Acids
• Equilibrium favors the weaker acid / base
Strength of Acids
weakest acid - CH4 < NH3 < H2O < HF - strongest acid
weakest acid - HF < HCl < HBr < HI - strongest acid
Atoms Connected To H Are More Acidic
Electronegativity
Size
1.20 - How Structure Affects AcidityWhy does size affect base stability more so than electronegativity?
Identifying trends - valence electrons of F-, Cl-, Br-, and I- are found in 2sp3, 3sp3, 4sp3, and 5sp3 shells respectively.
There is significantly more volume in a 5sp3 orbital than a 2sp3 orbital
The negative charge in I- is thus spread out over a much larger volume of space (see potential map), making it more stable, even though it is the least electronegative of the halide series.
Organic Acid (Carboxylic Acid)
R O
O
H
acidic H
1.18 - Organic Acids and BasesAcid and base behavior of carboxylic acids:
As an acid, a proton is donated (in reality (as indicated with arrows), the basic hydroxide ion takes the proton, and leaves the electrons from the O-H bond).
As a base, a lone pair of electrons from the sp2 oxygen of the carboxylic acid takes a proton from a hydronium ion (note also the electrons from the bond returning as a lone pair to oxygen)
1.21 - How Substituents Affect Acidity
Why the change in acidity?
Hydrogen is replaced with a more electronegative atom, which pulls the bonding electrons towards itself
Inductive electron withdrawal - pulling electrons through sigma () bonds
Inductive electron withdrawal stabilizes the conjugate base by decreasing the electron density about the oxygen atom
RULE - acid strength is (still) determined by stability of the conjugate base…but new atoms substituted for others (substituents) play a role in this stability
1.21 - How Substituents Affect AcidityRULE - the effect of a substituent on the acidity of a compound decreases as the distance between the substituent and the acidic proton increases
