Chapter 1: Introduction & Review
I) Origins of Organic Chemistry
A) Definition of Organic Chemistry
1) Organic chemistry is defined as the chemistry of carbon compounds.
2) Organic chemistry originally was defined as those compounds that were
derived from living organisms. (Vitalism)
B) Why Carbon?
1) Carbon readily undergoes catenation to form long chain and ring
structures. Catenation refers to carbon’s ability to form strong covalent
bonds with other carbon atoms, thereby forming the millions of different
organic compounds.
2) All life on Earth is carbon based.
3) Tremendous variety of organic compounds exists (e.g., amino acids,
proteins, fats, carbohydrates, vitamins, drugs, pesticides, synthetic fabrics,
polymers, etc.).
II) Principles of Atomic Structure
A) Two Structures found Inside Atom
1) Nucleus
- small highly dense (+) charged core where majority of atom’s
mass is found.
- consists of (+) charged protons & (0) neutral charged neutrons.
- mass of proton & neutron are similar.
2) Electron cloud
- majority of atom’s volume found outside the nucleus.
- consists of (-) charged electrons.
- mass of electron is much less than proton & neutron (1/1800).
- Electrons participate in chemical bonding & reactions.
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B) Chemical Elements: Atoms & Periodic Table
1) Fundamental building blocks found in ALL matter.
2) Elements constitute our “chemical alphabet.”
C) Isotopes: Two Definitions
1) Atoms that have the same atomic number but different mass (nucleon)
numbers.
Nucleon is another name for protons & neutrons found in atom’s
nucleus.
2) Isotopes have the same number of protons but different numbers of
neutrons.
Number of protons = Atomic Number (Z) for atom.
Sum of number of protons & neutrons = Mass Number (A)
3) In a neutral atom, the number of protons = number of electrons.
D) Ions
1) An ion is a charged atom.
2) Two classes of ions
Cations are (+) charged and result from metals losing electrons.
Anions are (-) charged and result from nonmetals gaining
electrons.
3) In an ion, the number of protons & electrons are different.
III) Electronic Structure of Atom
A) Chemical Reactivity: Atomic Orbitals
1) Chemical reactivity is a result of the interaction of electron clouds
between atoms.
Heisenberg Uncertainty Principle- impossible to know the exact
location of an electron. We refer instead to electron density,
(probability of finding an electron in a particular part of an orbital).
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2) Electrons are distributed, within the electron cloud, in certain allowable
energy levels. The energy levels found in an atom are identified by the
principal quantum number n and are referred to as shells.
n = { 1, 2, 3, ... } which correspond to the K shell, L shell, etc.
n is same as the row number in the periodic table.
most elements in organic chemistry are found in first 2
rows of periodic table (n= 1 & 2).
2n2 = maximum number of electrons in a given n shell
3) Within a given shell are various sublevels consisting of atomic
orbitals, regions of space where electrons reside.
Examples of atomic orbitals are 1s, 2s, 2p, 3s, 3p, etc.
Atomic Orbitals are characterized by four quantum numbers
Principal
Azimuthal
Magnetic
Spin
Energy level
Shape of orbital
Orientation in space
Electron spin
n
l
ml
ms
Each spherical s-orbital stands alone and is non-directional.
There are three degenerate dumbbell shaped p-orbitals that
differ only in their orientation in 3-D space.
Degenerate orbitals are orbitals that are equivalent in
energy.
There are five degenerate d-orbitals, & seven degenerate f-orbitals.
4) Individual orbitals can accommodate a maximum of two electrons.
Pauli Exclusion Principle: Only two electrons maximum are
allowed in an orbital. The paired electrons have opposite spins.
No two electrons in an atom will have the same set of four
quantum numbers (n, l, ml , ms ).
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Table 1.1. Distribution of Electrons Inside Electron Cloud
Principal Energy Level (n)
Maximum No. of Electrons (2n2)
Atomic Orbitals in Sublevel
Designation of Filled Orbitals
Maximum Electrons per Sublevel
Orbitals per Sublevel
1
2
1s
1s2
2
1
2
8
2s, 2p
2s2, 2p6
2, 6
1, 3
3
18
3s, 3p, 3d
3s2, 3p6, 3d10
2, 6, 10
1, 3, 5
4
32
4s,4p,4d,4f
4s2, 4p6, 4d10, 4f14
2, 6, 10, 14
1, 3, 5, 7
B) Electronic Configurations
1) Electrons are introduced into atomic orbitals via the Aufbau Principle
(a.k.a., Diagonal Rule).
1s
2s 2p
3s 3p 3d
4s 4p 4d 4f
5s 5p 5d 5f
6s 6p 6d
7s 7p
2) Hund’s Rule: Electrons will fill into degenerate orbitals one at a time,
until all are half-filled before pairing up.
C) Valence Electrons
1) Electrons that participate in chemical bonding.
2) Valence electrons are found in the highest n-level of an atom.
Example: What is electronic configuration for Carbon?
Answer: Carbon (Z=6) is the first member of column IVA and is
found in the second row of the periodic table. (valence level = 2)
The electronic configuration for C is 1s22s22px12py12pz . We see
that there are four valence electrons in a carbon atom.
3) For main group elements (s & p block), the column number is the
same as the number of valence electrons.
Group IA has 1 valence electron, IIA has 2 valence electrons, IIIA
has 3 valence electrons, etc.
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IV) Chemical Bonding
A) Types of Chemical Bonding
1) G. N. Lewis proposed the concept of bond formation.
Octet Rule: Atoms either transfer or share valence electrons to
attain a noble gas electronic configuration (filled shell of eight
valence electrons).
2) Ionic Bonding
- Transfer of valence electrons from one atom to another.
- Occurs between metal and nonmetal to form ions (metal cation &
nonmetal anion) that orient themselves into crystal
lattices.
- Rarely encountered in organic chemistry.
-
[Ne] 3s1
Example: NaCl
[Ne] 3s23p5
[Ne]
[Ne] 3s23p6
3) Covalent Bonding
- Sharing of valence electrons by nonmetal atoms.
- Commonly seen in organic chemistry.
- Example: H2
H:H
H-H (Dash = shared electron pair)
B) Bond Polarity
1) Nonpolar bond - bond with electrons shared equally between two
atoms.
2) Polar bond - unequally shared pair of bonding electrons.
- Most “covalent” bonds actually have some bond polarity or
separation of charge.
- Charge separation measured using dipole moment (  ) where
 =  d ( is the amount of charge separation and d is bond length)
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- Electronegativities are used to determine bond polarities where
the more electronegative atom has the partial negative charge.
- Rule: Electronegativity increases going across a row of periodic
table and decreases down a column (group). Values are based on
Pauling scale (F defined as 4.0, the largest value).
- The greater the electronegativity difference, the more polar
the bond. If the difference is very large (> 2.0), the bond is
considered ionic.
+ - Example:
CF bond is written as
CF
- Usually assume C and H have same electronegativity, thus
CH bonds are considered as nonpolar.
V) Lewis Structures
1) Lewis Structures are 2-D representations of the valence electrons found in a
molecule or polyatomic ion.
2) Two types of valence electrons found in a Lewis structure.
Bonding electrons are the valence electrons that form a covalent bond.
single bond
double bond
triple bond
1 pair of shared electrons
2 pairs of shared electrons
3 pairs of shared electrons
Nonbonding (lone pair) electrons are the valence electrons that are found
on an individual atom and do not participate in chemical bonding.
3) The number of bonds an atom usually forms is its valence. To draw Lewis
structures correctly, it is important to become familiar with common bonding
patterns for uncharged atoms often encountered in organic chemistry.
Table 1.2.
Element
Common Bonding Patterns for Uncharged Atoms
C
N
O
H
X = Halogens
# bonds (Valence) 4
Number Lone pairs 0
3
1
2
2
1
0
1
3
7
Examples of the most common bonding patterns found in organic
compoundsa
Atom
Valence
C
4
N
5
O
6
halogen
7
Positively charged
Neutral
Negatively charged
aThe
formal charge on the central atom depends on the number of bonds
(shared pairs) and the number of nonbonding electrons
A) The Rules for Writing Lewis Structures:
1) Need the molecular formula and the atomic positions (connectivity of
atoms) which are hopefully given in the problem. Consider the
following example:
CH3NO2  Is connectivity H3CONO or H3CNO2 ?
It can be either one depending upon whether it is methyl nitrite
(H3CONO) or nitromethane (H3CNO2).
Let's look at methyl nitrite, written as CH3ONO to illustrate the atomic
positions. The atoms are arranged as shown below.
H
H
C
O
N
O
H
2) Determine the total number of valence electrons. This is obtained by
multiplying the number of atoms by the group number for each atom. If
one has an ion, the overall charge on the species has to be incorporated
into the total.
CH3NO2  1(4) + 3(1) + 1(5) + 2(6) = 24 valence electrons
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Note: Group numbers are in parenthesis.
3. Connect bonded atoms by a dash (  ) which represents the shared
electron pair in the bond.
H
|
H  C  O  N  O
|
H
Structure contains 6 single bonds (12 valence electrons). We need
to account for 24 total valence electrons.
4. Add remaining valence electrons as lone pairs to complete the octets of
those atoms (except hydrogen) that have fewer than eight.
5. If one atom does not have an octet, then use unshared pairs on an
adjacent atom to form a double or triple bond to complete the octet.
(Recall from Table 1.2 that O usually forms two bonds and N has three
bonds).
6. Check for formal charges (FC) on each atom:
FC = Group number - # of bonds - # of unshared electrons
Formal charges should always add up to the overall charge
found on the molecule or ion.
Ideally we want to write Lewis structures that has no formal
charge separation.
Note: It is not always possible to eliminate formal charges.
H
|



H  C  O  N == O
|


H
Carbon:
4-4-0=0
Hydrogen:
1-1-0=0
Oxygen (-O-):
6-2-4=0
Oxygen (=O):
6-2-4=0
Nitrogen:
5-3-2=0
Sum Total
=0
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Example 2:
Draw the Lewis structure for CH3O - (methoxide ion).
(Note: -1 charge on ion)
Challenge Problem: Draw Lewis structure for nitromethane (H3CNO2).
7. Rule: Starting with n=3 (the third row of the periodic table), an element
can expand its valence shell of electrons beyond an octet (usually to 10 or
12). This phenomenon mainly occurs with P, S, As, Se, etc.
Example: H2SO4
(sulfuric acid)
Poor Lewis Structure (large charge separation)
Calculate formal charges for all atoms.
Best Lewis Structure (zero charge separation)
Calculate formal charges for all atoms.
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VI) Resonance
1) Resonance - Compounds that have more than one possible Lewis structure,
differing only in the placement of electrons. The actual structure is a resonance
hybrid of its resonance forms.
A) Rules for Drawing Resonance Structures
1) Atomic positions (connectivity of atoms) must be the same in all
resonance structures, only the electron positions can vary (i.e., lone pairs
and multiple bonds).
2) The stable resonance structures must satisfy the octet rule and have the
smallest separation of oppositely charged atoms.
3) We want any negative formal charge on the most electronegative atom.
4) Each Lewis structure must have the same number of electrons and the
same net charge.
5) Each must have the same number of unpaired electrons.
6) Electron delocalization stabilizes a molecule and is best when all the
resonance structures are of the same energy.
7) Frequently we will talk about major versus minor contributors:
H2C == N+H2

H2C+  NH2
This structure is best since
C and N both have an octet |
even though plus charge is on
the N.
This structure is poor even though
plus charge is on the C because C
does not have an octet.
(major contributor)
(minor contributor)
Note: See Table 1.2 on page 7 of notes packet & work enough
problems to commit the bonding patterns to memory!
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VII) Formula Writing in Organic Chemistry
A) Chemical Formulas
1) Empirical Formulas
- An empirical formula is simply the relative ratios of the
elements present.
- Ratios are determined through elemental analysis and are
expressed as percent composition. Sum of percentages for elements
should add up to 100 %.
Note: Cannot obtain % O directly.
% O is assumed to be 100.00% - sum total of percents
listed in problem.
- Review how to calculate empirical formulas (general chemistry).
2) Molecular Formulas
- Actual chemical formula and is represented by the number of
atoms of each element in one mole of the compound.
- Very often is found to be a multiple of empirical formula.
- To find the molecular formula one must have the empirical
formula & the molecular weight (MW) of the compound.
EMM (x) = MW
EMM = Empirical Molecular Mass
x = multiple (used to find molecular formula)
MW = molecular weight of compound (given in problem).
B) Structural Formulas
1) Structural Formulas show which atoms are bonded to one another.
2) Two types of Structural Formulas
- Complete Lewis structures which show bonding & nonbonding
electrons.
- Condensed structural formulas do not show all the individual
bonds, instead each “central” atom is shown together with the
atoms bonded to it.
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Examples:
CH3CH2CH2CH3 or CH3(CH2)2CH3
CH3CH2OCH3
Note: Lone pairs are usually not shown , but you have to know that
they are there. Multiple bonds (double, triple, etc.) are often shown
in condensed structural formula.
C) Line Angle Formula
1) Line-angle formulas are usually used for ring systems. Bonds are
represented by lines.
2) Assume that carbon atoms are found at the end of a line or the
intersection of two or more lines.
3) Atoms other than carbon are shown at their respective locations in the
molecule. Hydrogens are not shown unless the hydrogen is bound to a
drawn atom.
VIII) Acids & Bases
A) Arrhenius System
1) An acid is a substance that liberates H3O+ ions into solution. Acids
contain H+ in chemical formula. (HCl, HNO3, H2SO4, etc.).
2) A base is a substance that liberates OH- ions into solution. Bases
contain OH- ions in chemical formula (LiOH, NaOH, Ca(OH)2, etc.).
3) Acids & Bases classified as strong or weak depending upon degree of
dissociation.
4) Mathematical Relationships (Acids/Bases) in aqueous solution:
pH = - log [H+]
pOH = - log [OH-]
pH + pOH = 14.0
pKa = - log Ka
pKb = - log Kb
KaKb = 1 x 10-14
B) Bronsted-Lowry
1) An acid is a proton donor. A proton is equivalent to a H+ ion.
2) A base is a proton acceptor. Base must have an unshared pair of
electrons that are used to make bond with the proton.
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3) Conjugate acid/base pairs differ from each other by H+.
Examples: NH4+ / NH3 (conjugate acid / base)
C) Acid Strength
1) Acid strength defined by comparing pKa values. Weak acids have large
pKa values. Strong acids have pKa values  0.
2) From general chemistry comes this useful axiom. "The stronger the
acid the weaker is its conjugate base or the weaker the acid, the
stronger is its conjugate base."
3) Acid-base reactions favor the formation of the weaker acid and the
weaker base.
D) Predictions of Acidity or Basicity
1) For hydrogens attached to atoms, the acidity depends on the bond
strength.
a) Vertical column of periodic table, the strength of the bond to
the proton is the determining effect. Therefore, acidity
increases as we descend a vertical column.
HF < HCl < HBr < HI
F - > Cl - > Br - > I -
Conjugate Base Strength
b) Horizontal row of periodic table, acidity depends on
electronegativity of the atom. Thus,
H-F > HO-H > H2N-H > H3C-H
H3C - > H2N - > HO - > F -
Base Strength
c) Also have resonance and inductive effects.
Resonance - when a molecule or ion can be represented by two
or more equivalent resonance structures, we have resonance
stabilization.
Inductive Effects - these are through the bonds and involve
electron-withdrawing effects which fall off as the distance from
the center increases.
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Example: Why is CH3COOH a stronger acid than CH3OH ?
Due to resonance CH3CO2 - is a weaker base than CH3O which means that CH3COOH is a stronger acid than CH3OH.
Rule: The more equivalent resonance structures that we can
write for the conjugate base of an acid, the stronger is the acid.
IX) Lewis Acids & Bases
A) Definitions
1) Lewis acid - an electron pair acceptor.
2) Lewis base - an electron pair donor.
3) Lewis acids are called electrophiles (electron loving) and Lewis bases
are termed nucleophiles.
B) Curved Arrow Formalism
1) Method that shows flow of an electron pair from the electron donor
(nucleophile) to the electron pair acceptor (electrophile).
2) It is a symbolic way of keeping track of the flow of electrons in a
chemical reaction and is a pivotal concept in organic chemistry.
3) Examples:
Now we have a new bond formed.
Chapter 1 Problem Notebook Questions:
Assignment #1
Assignment #2
Assignment #3
Note: Items found in Wade, 8th Ed.
1-21, 1-23, 1-25 through 1-27
1-31, 1-33 (a,b), 1-34 through 1-37
1-38, 1-42 through 1-45