Fundamentals of Organic Chemistry
McMurry / Simanek ; 6e
Chapter 1.
Structure and Bonding ; Acids and Bases
1.1
1.2
1.3
1.4
1.5
1.6
1.7
1.8
1.9
1.10
1.11
Atomic Structure 2
Electron Configuration of Atoms 4
Development of Chemical Bonding Theory 5
The Nature of Chemical Bonds 6
Forming Covalent Bonds: Valence Bond Theory 9
Hybridization: sp3 Orbitals and the Structure of Methane 10
Hybridization: sp3 Orbitals and the Structure of Ethane 11
Double and Triple Bonds 12
Polar Covalent Bonds: Electronegativity 16
Acids and Bases: The Brønsted–Lowry Definition 18
Acids and Bases: The Lewis Definition 22
3

Organic Chemistry : Constituents of natural material around human, such as
food, wood, fiber medicine, plastic from living organism.

cf, inorganic chemistry deals with all 103 element, while organic chemistry focuses
on carbon and 20 first elements including H2, O2, and N2.
Criterion
Elements
Bonding
Rate of rxn.
Conductivity
Tm
Volatility
Solubility in water
Solubility in organic solvent
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Inorganic
Organic
All 103
Ionic
Fast
Electrolyte
>700
Nonvolatile
Yes
No
a few (C,H,O,N,S,P,F,Cl)
Covalent
Slow
Nonelectrolyte
<300
Volatile(distillation)
No
Yes
4

Friedrich Wöhler discovered in 1828,

convert the “inorganic” salt to “organic” compound

The only distinguishing characteristic of organic compounds is that all contain the
element carbon.
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
Organic Chemistry : study of compounds of carbon

Basics of chemistry and also living things

Chemistry of carbon and hydrogen (mostly)

Carbon to bond together, forming ring and long chain [ex., diverse products from
methane to DNA(tens of billion)]

Not all organics compounds are derived from living organism, but most medicine,
dye, polymer, pesticide, etc.

Relation between molecular structure, properties, and their reactivity

Designing molecules and its understanding become more important today (NT,
BT, etc)
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1.1 Atomic Structure

Atom consists of dense, positively charged nucleus surrounded by negatively
charged electron

Description of atom ;

원자(Atoms): nucleus+ electrons

핵(Nucleus): positively charged, most of the mass of the atom

protons + neutrons

전자(Electrons)

양성자(Protons): positively charged

중성자(Neutrons): neutral

원자번호(Atomic number): equal to the number of protons

원자량(Atomic weight): ~ sum of protons & neutrons
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How Electrons are Arranged in Atoms

Atom’s electrons


Their number & arrangement determines how an atom reacts with other atoms to
form molecules
Electron distribution

Electrons are not free to move and they are confined to different region with
different energy level

Electrons belongs to different layer or shells around nulceus

The larger shell, the more electrons can be hold
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

Orbitals: certain region of space that electrons are located

Each orbital can contain a maximum of two electrons

The orbitals differ in shape: s, p, d…

Orbitals are grouped in shell

Shell: 1, 2, 3…

Each shell contains different types of orbitals
Numbers of orbitals and electrons in the first three shells
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Figure 1.3 The energy levels
of electrons in an atom.
Figure 1.4 Representations of s and p orbitals.
(a) The s orbitals are spherical, and (b) the p orbitals are dumbbell-shaped. The lobes of p orbitals are
often drawn for convenience as “teardrops,” but their true shape is more like that of a doorknob, as
indicated by the computer-generated representation. (c) The three p orbitals in a given shell are
oriented along mutually perpendicular directions.
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1.2 Electron Configuration of Atoms

Ground-state electron configuration



Rule 1 (aufbau 원리)

lowest energy are filled first,

according to the order 1s → 2s → 2p
→ 3s → 3p → 4s → 3d
Rule 2 (Pauli의 배타원리)


The lowest-energy arrangement the
atom’s electrons occupy
Only two electrons with opposite
spin occupy a obital
Rule 3 (Hund 규칙)

Until all orbitals are half-filled, one
electron is placed in each with spin
parallel
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Table 1.1 Ground-State Electron
Configuration of Some Elements
11
Electron arrangements of the first 18 elements
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Ref. Pauli exclusion principle

Each atomic orbital can contain two electrons

These electrons have different spin state

If electrons have same spins, they repulse each other as far as possible

전자가 같은 부호의 spin을 가지면, 가능한 한 멀리 떨어지려 함

분자형태와 성질을 결정하는 중요한 인자
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1.3 Development of Chemical Bonding Theory

In 1858, August Kekulé and Archibald Couper, proposed that carbon has four
“affinity units.”  tetravalent  carbon forms four bonds to form stable element


Kekulé, carbon atoms can bond to one another to form extended chains, and chains
can double back on themselves to form rings. (still viewed in a two-dimensional
way until 1874)
Jacobus van’t Hoff and Joseph Le Bel ; 3rd dimensional ideas (1874)
Figure 1.5 Representation of a tetrahedral
carbon atom. The heavy wedged line comes
out of the plane of the paper, and the dashed
line goes back into the plane.
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1.4 The Nature of Chemical Bonds

Why chemical bonding formation ?

When chemical bond is formed, energy is released and the bonding state is more
stable than separate one


Make bonds release energy, break bond  absorb energy
Octet rule ;

Group 8A : neon (2+8), argon (2+8+8), krypton(2+8+18+8).

Eight electrons in the outermost shell, or valence shell, gives more stability to form
noble-gas (rare gas)

Valence electrons: located in the outermost shell, mainly involved in chemical bonding

Group numbers correspond to the number of valence electrons (except He)
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Ionic and Covalent Bonding


Lewis’ theory of chemical bonding

Inert gas: very stable electron arrangement because they don’t combine with other
atoms

Other atoms might react in such a way as to achieve these stable arrangement

Complete transfer of electrons from one atom to another

Sharing of electrons between atoms
Ionic bonds

Formed by the transfer of valence electrons

Forms “cation” and “anion” NaCl

Na: 1valence electron; donate electrons(cation) ; electropositive(metals)

Cl: 7 valence electron; accept electrons(anion) ; electronegative(nonmetals)

Electropositive atom + electronegative atoms

Attractive force between opposite charges
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Carbon and the Covalent Bond


Carbon ; 4 valence electrons

Valence shell is half filled.

Neither electropositive nor electronegative.

Usually covalent bond

Methane: sharing electron with 4 hydrogens

By sharing electron pairs, the atoms complete their valence shells, forms molecule ;

In 1916 by G.N. Lewis, proposed Covalent bond

Neutral collection of atom by covalent bond : Molecules
Notation of covalent bond

Lewis structure : electron-dot structures

Kekulé structures : line-bond structures
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1.5 Forming Covalent Bonds: Valence Bond Theory

Covalent Bond

Image of covalent bond :
overlapping two atomic orbital

Bond between elements that have
similar electronegativity

H2


Bond energy (binding energy,
BE) : 104 kcal/mol release

Bond length (balance between
repulsion & attraction) :
0.74Å(=74pm)
Sigma(σ) bonds ; for H-H bond,
elongated gas shape orbital

Formation by head-on overlap of
two atomic orbital circular cross
section
숭실대학교 환경화학공학과
Figure 1.6 A plot of energy versus internuclear
distance for two hydrogen atoms.
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1.6 Hybridization: formation of sp3 Orbitals , structure of Methane

CH4 (2s2 2p2)
 two kinds of orbital for bonding C-H in fact, all four bonding is identical


Answer : s orbital and p orbital combine or hybridize to form four equivalent atomic orbital that
are spatially oriented toward for corner of tetrahedron
sp3 orbital ; hybrid orbital by combination of one s and three p orbital

in 1931 by Linus Pauling

Unsymmetrical orbital

Can overlap better to form bond and more stronger bond than unhybridized sp
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1.7 Hybridization: sp3 Orbitals and the Structure of Ethane
Some representations of ethane

Image of ethane molecules
420 kJ/mol
376 kJ/mol
sp3 carbon
sp3 carbon
숭실대학교 환경화학공학과
sp3–sp3 σ bond
20
1.8 Double and Triple Bonds
1. Double and triple bonds are far more reactive than single bonds.
2. A CC bond is not twice as strong as a C–C bond and a CC bond is not three times
as strong as a single C–C bond.
3. Double bonds lead to a flat (or planar) shape. Triple bonds lead to a linear shape.

Imagine combination the carbon 2s orbital with only two of three available 2p
orbital to form sp2 orbital
∴Three sp2 hybride orbital and one unhybridized 2p orbital remain unchanged
 three sp2 orbitals lie in at 120° angle
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Figure 1.10 The hybridizations of carbon. Varying the number of 2p orbitals that are
hybridized with the 2s orbital provides sp3, sp2, and sp hybrid orbitals.
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π bond : double bond

π bond ; covalent bond by sideway overlap of two p-orbital

For ethylene,

The combination of sp2-sp2 sigma overlap and 2p-2p pi (π) overlap result in net
sharing of four electron and four electron and formation of C=C double bond

Four hydrogen atom form sigma bond to remaining four sp2 orbital

Planar structure of ethylene with H-C-H and H-C-C angle of 120°


C-H : 1.076Å, 103kcal/mol
C=C is more stronger and shorter than C-C bond
∵ double bond results from four electron sharing rather than two electron

C=C : 133pm, 610 kJ/mol

C-C ; 154pm, 376 kJ/mol
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π bond : triple bond

Hybridization ; sp orbital and acetylene structure by sharing 6 electron
 form

triple bond
Image the combination of 2s carbon orbital with a single 2p orbital
 two sp hybride orbitals result and two p orbital remain unchanged
 two sp orbital and linear on perpendicular to y-axis and z-axis

Triple bond formation ;

sp-hybrid orbital overlap to form strong

sp-sp sigma bond and pz-pz π bond and py-py π bond formation by sideway overlap

One sigma bond and two pi bond  carbon triple bond
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Implication of π bond between C=C and C≡C
1. 4 electron show in C=C vs. 6 electron show in C ≡ C

π electron are kept in distance instead crowding nuclei
2. C=C bond strength : 270 kJ/mol < C-C(370 kJ/mol), C ≡ C (245 kJ/mol)

π electron do not hold tightly as σ bond
3. C=C or C ≡ C more reactive
∵ π -bond are weaker and π -electron are available for rxn instead of hidden ebetween two nuclei
4. Bond length : C-C > C=C > C ≡ C


To provide best overlap between p-orbital, atoms squeeze closer together
sp orbital is shorter and rounder than sp2, sp3  make σ bond shorter
5. Four sp3 orbital  tetrahedral shape : max. distance & min. repulsion

But sp2, p orbital to form bond  planar geometry

sp, linear geometry
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Table 1.3 Summary of Bonding for Carbon
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1.9 Polar Covalent Bonds: Electronegativity

CH3-CH3 : electronically symmetric, tow bonding electron are equally shared


Na+Cl- : electron is transferred from sodium to chloride to give Na+ & Cl

Full covalent Y : Y
Full ionic
M+M-
Most majority chemical bond, shared electrons are attracted more strongly by
one and the other

Polar covalent bond Xδ+ :Yδ- (δ+ , partial positive)
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
Bonding polarity

Due to difference in electronegativity the intrinsic ability of atom to attract electron
in covalent bond
Figure 1.12 Electronegativity values and trends. Elements in red are the most electronegative, those in green are medium, and those in yellow are the least electronegative.
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
Bond of atoms with similar electronegativity(EN) are covalent

ΔEN < 2 , polar covalent bond

ΔEN > 2 , ionic bond

Bonds between carbon and less EN element


Tetraethyl lead (organometallic compound)
Inductive effect : atom’s ability to polarize bond
 shift electrons in response to electronegative difference


Inductively donate electron : Li, Mg

Inductively withdraw electron : O, Cl
Most chemical reaction  due to inductive effect
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Ref. Formal Charges

Formal Charges(형식 전하) - 공유결합을 이루고 있는 원자들의 전하분포
를 알아내는 방법

formal charge = (valence electrons in the atom) – (number of unshared
electrons) – 1/2(number of shared electron)

형식전하 = 최 외각 전자 수 – 비 공유전자 수 – 1/2공유전자 수
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Ref. Formal Charges
HCN vs HNC
•
•
•
•

Lewis structure
Charge
Formal charge
Stability
H:C:::N:
0
00 0
>
:C:::N:H
0
-1 +1 0
Fewer formal charge  More stable
Two oxygens are different according to the structure,
While they are the same according the experiments!
Why?
Resonance!
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Ref. Resonance

원소의 위치는 변하지 않고 전자의 이동으로 생기는 현상

일반적으로 공명구조가 많을수록 안정한 화합물이 된다

Double-headed 화살표로 표시
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Ref. Resonance

The actual structure of CH3NO2
 Any
Lewis structure cannot represent the real molecule (experiment)
 the actual structure is a resonance hybrid
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Ref. Shapes of molecules


Shapes of molecules are determined by actual experiments not by theoretical
considerations!

It is Chemistry and Science.

Experiment First, then Theory

There are some rules for the prediction of the shape.
One is VSEPR (valence shell electron pair repulsion) theory

Rule 1: Pairs of electrons in the valence shell repel each other

Rule 2: Unshared electron pairs repels more

Rule 3: Double and triple bonds as one electron pair
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Ref. VSEPR (valence shell electron pair repulsion) theory
Rule 1:
Pairs of electrons in the valence
shell repel each other
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Ref. VSEPR (valence shell electron pair repulsion) theory
Rule 2:
Unshared electron pairs
repels more
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Ref. VSEPR (valence shell electron pair repulsion) theory
Rule 3:
Double and triple bonds as one electron pair
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1.10 Acids and Bases: The Brønsted–Lowry Definition


Acidity & basicity – similar concept – electronegativity & polarity

The Brønsted–Lowry definition and the Lewis definition
Brønsted–Lowry definition acid & base

Acid : a substance to donate a proton (hydrogen ion, H+)

Base : a substance to accept a proton (hydrogen ion, H+)
Acid

Base
Conjugate acid
conjugate base
The ability to donate proton ∝ acidity constant

Range of Ka : 1015~10-60 (strong acid  week acid)

Inorganic acid 102~109, most organic acid10-5~10-15

pKa : acid strength
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Understanding of proton performance

Strong acid lose proton more easily  conjugate base hold proton less tight
 weak base

Acid strength of acid  inverse form  Base strength of conjugate base

HCl is strong acid means Cl- does not hold proton tightly and thus weak base

Inversely water(H2O) is weak acid = OH- holds proton tightly and strong base
Table 1.4 Relative Strength of Some Common Acids
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Acetic acid
(pKa 4.76)

Acetate ion
Water
(pKa 15.74)
Hydroxide ion has more strong affinity for proton than acetate ion
 water
 for

Hydroxide ion
is weaker acid than acetic acid (inversely, acetic acid is more strong acid)
conjugate base, pKa ↑ more affinity to proton ↑
Reaction element stability ;

Predict acid-base reactivity

Product should be more stable, if rxn occurs , product acid & base is less reactive
than stating material (acid & base)
Stronger
acid
숭실대학교 환경화학공학과
Stronger
base
Weaker
acid
Weaker
base
1-40

Two main groups of organic acid
(1) those with acidic hydrogens attached to oxygen atoms in compounds like methyl
alcohol and acetic acid and
(2) those where the acidic hydrogen
atom is attached to a carbon atom.
∵ C–H hydrogen is not acidic
as O–H hydrogens

Organic base ; atom with lone pair electron that can bond to H+

Nitrogen-containing compounds such as trimethylamine are the most common

but oxygen-containing compounds can also act as bases when reacting with a
sufficiently strong acid.
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1.11 Acids and Bases: The Lewis Definition


Lewis acid : a substance that accept an
electron pair
Lewis base : a substance that donates an
electron pair
cf) Donated electron pair is shared

Proton : Lewis acid since it accept a pair of electron to fill vacant is orbital


AlCl3 : too accepts an electron pair from a Lewis base to fill a vacant valence orbital
Lewis base definition : a compound with a pair of nonbonding electrons to form
a bond to Lewis acid

similar to Brønsted–Lowry definition.
H2O act as Lewis base by donating electron
pair to proton in forming H3+O ion
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
In general, most oxygen & nitrogen containing compound : Lewis base
∵ lone pair of electron

Alcohol & carboxylic acid

Acts as acid : lose OH proton
as base : oxygen accept proton
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Summary and Key Words

Acidity constant, Ka, p. 19

Lewis acid, p. 22

Bond angle, p. 11

Lewis base, p. 22

Bond length, p. 9

Lewis structure, p. 7

Bond strength, p. 9

Line-bond structure, p. 8

Brønsted–Lowry acid, p. 18

Lone-pair electrons, p. 7

Brønsted–Lowry base, p. 18

Molecule, p. 7

Conjugate acid, p. 18

Nonbonding electrons, p. 7

Conjugate base, p. 18

Orbital, p. 3

Covalent bond, p. 7

Organic chemistry, p. 2

Electron shell, p. 3

Pi (π) bond, p. 13

Electronegativity, p. 16

Polar covalent bond, p. 16

Ground-state electron

Sigma (σ) bond, p. 11

configuration, p. 4

sp Hybrid orbital, p. 12

Inductive effect, p. 17

sp2 Hybrid orbital, p. 12

Ionic bond, p. 7

sp3 Hybrid orbital, p. 10

Isotope, p. 3

Valence bond theory, p. 9

Kekule structure, p. 8

Valence shell, p. 6
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To solve the problems
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To solve the problems
숭실대학교 환경화학공학과