Transcript Orbitals

University of Isfahan
Department of Chemistry
90-91 ‫نيمسال اول‬
FUNDAMENTALS OF
ORGANIC CHEMISTRY
FOR BIOLOGY STUDENTS
Dr. Gholam Ali Koohmareh
Text Book:
Fundamentals of Organic Chemistry
By: John. Mc. Murry
‫مباني شيمي آلي‬
‫نوشته جان مك موري‬
‫ترجمه دكتر عيسي ياوري‬
2
‫نصايحي چند در مورد مطالعه شيمي آلي‬
‫ساختار دانش در شيمي‬
‫آلي‬
‫ماده مفقوده‬
‫ساختارها شبيه هرم‬
‫متعادلي است كه روي‬
‫نوك ايستاده است‬
‫اگر چيزي گم شده باشد‪.....‬‬
‫كل ساختار خراب ميشود!!!‬
‫براي درك ساختار مواد بايد مطالعه كنيد‪.‬‬
‫هرمطلبي را كه مسلط نميشويد چندين بار‬
‫تكرار كنيد‪.‬‬
‫ما از مباحث ساده شروع ميكنيم و وارد بحثهاي‬
‫پيچيده تر ميشويم‪.‬‬
‫اگر مطلبي را متوجه نميشويد هر چه زودتر و قبل از‬
‫انباشته شدن مطالب كمك بگيريد‪....‬‬
‫ميتوانيد بعد از كالس يا در طول ساعات اداري در اطاق من در گروه شيمي‬
‫و يا از طريق ايميل مشكل درسي خود را مطرح كنيد‬
‫‪Email: [email protected]‬‬
‫هر زماني كه وقت داشتيد تمرين حل كنيد‪.‬‬
‫يادگيري فعال بسيار موثرتر ازمطالعه اجباري است!‬
‫انجام تمرينات درك شما را بيشتر ميكند و اين مستلزم آن است‬
‫كه هر چه ياد ميگيريد بكار گيريد‪.‬‬
‫تمرينات آخر فصل بسيار مفيدند‪.‬‬
‫عجيب ولي واقعي‬
‫فعال‬
‫‪ 90%‬چيزي را كه ميگوييم و بكار ميگيريم‬
‫‪ 70%‬چيزي را كه ميگوييم و صحبت ميكنيم‬
‫‪ 50%‬چيزي را كه ميبينيم و ميشنويم‬
‫ما حفظ ميكنيم‬
‫‪ 30%‬چيزي را كه ميبينيم‬
‫‪ 20%‬چيزي را كه ميشنويم‬
‫اجباري‬
‫‪ 10%‬چيزي را كه ميخوانيم‬
‫يك مطالعه روانشناسي ‪…..‬‬
‫فصل اول‪ .‬ساختار و پيوند‬
Organic Chemistry


“Organic” – until mid 1800’s referred to compounds
from living sources (mineral sources were
“inorganic”)
Wöhler in 1828 showed that urea, an organic
compound, could be made from a non-living source:
10
Organic Chemistry

Today, organic compounds are those
based on carbon structures and
organic chemistry studies their
structures and reactions


Includes biological molecules, drugs,
solvents, dyes, food additives, pesticides,
and others.
Does not include metal carbonate salts and
other simple ionic compounds (inorganic)
11
Organic Chemistry
12
1.1 Atomic Structure

Structure of an atom



Positively charged nucleus (very dense, protons and
neutrons) and small (10-15 m)
Negatively charged electrons are in a cloud (10-10 m) around
nucleus
Diameter is about 2  10-10 m (200 picometers (pm))
[the unit angstrom (Å) is 10-10 m = 100 pm]
13
1.1 Atomic Structure
14
Atomic Number & Mass
Number



The atomic number (Z) is the
number of protons in the atom's
nucleus
The mass number (A) is the
number of protons plus neutrons
All the atoms of a given element
have the same atomic number
15
Atomic Number and Atomic
Mass



All the atoms of a given element have the
same atomic number
Isotopes are atoms of the same element
that have different numbers of neutrons and
therefore different mass numbers
The atomic mass (atomic weight) of an
element is the weighted average mass in
atomic mass units (amu) of an element’s
naturally occurring isotopes
16
1.2 Atomic Structure: Orbitals

Quantum mechanics: describes electron
energies and locations by a wave equation





Wave function is a solution of the wave equation
Each Wave Function describes an orbital, 
A plot of  2 describes the probabililty of
finding the electron
The electron cloud has no specific boundary
so we show the most probable volume.
The boundary is called a “probablility surface”
17
Shapes of Atomic Orbitals for
Electrons



For the existing elements, there are
four different kinds of orbitals for
electrons based on those derived for a
hydrogen atom.
Denoted s, p, d, and f
s and p orbitals most important in
organic chemistry (carbon atoms have
no d or f orbitals)
18
Shapes of Atomic Orbitals for
Electrons


s orbitals: spherical, nucleus at center
p orbitals: dumbbell-shaped, nucleus at
middle
19
Orbitals and Shells



Orbitals are grouped in shells of
increasing size and energy
As shells increase in size, they contain
larger numbers and kinds of orbitals
Each orbital can be occupied by two
electrons (Pauli Exclusion Principle)
20
Orbitals and Shells



First shell contains one s orbital,
denoted 1s, holds only two electrons
Second shell contains one s orbital (2s)
and three p orbitals (2p): eight
electrons
Third shell contains an s orbital (3s),
three p orbitals (3p), and five d orbitals
(3d): 18 electrons
21
Orbitals and Shells
22
p-Orbitals


In each shell there
are three
perpendicular p
orbitals, px, py, and
pz, of equal energy
Lobes of a p orbital
are separated by
region of zero
electron density, a
node
23
1.3 Atomic Structure: Electron
Configurations




Ground-state electron configuration of an atom
lists orbitals occupied by its electrons.
1. Lowest-energy orbitals fill first: 1s  2s  2p 
3s  3p  4s  3d (Aufbau (“build-up”) principle)
2. Electron spin can have only two orientations, up 
and down . Only two electrons can occupy an
orbital, and they must be of opposite spin (Pauli
exclusion principle) to have unique wave equations
3. If two or more empty orbitals of equal energy are
available, electrons occupy each with spins parallel
until all orbitals have one electron (Hund's rule).
24
Electron Configurations
25
Electron Configurations
26
1.4 Development of Chemical
Bonding Theory


Kekulé and Couper independently observed that
carbon always has four bonds
van't Hoff and Le Bel proposed that the four bonds of
carbon have specific spatial directions

Atoms surround carbon as corners of a tetrahedron
Note that a dashed line
indicates a bond is behind
the page
Note that a wedge indicates a
bond is coming forward
27
Tetrahedral Carbon:
28
1.5 The Nature of the Chemical
Bond



Atoms form bonds because the compound
that results is more stable than the separate
atoms
Ionic bonds in salts form as a result of
electron transfers
Organic compounds have covalent bonds
from sharing electrons (G. N. Lewis, 1916)
29
1.5 The Nature of the Chemical
Bond

Lewis structures shown valence
electrons of an atom as dots



Hydrogen has one dot, representing its 1s
electron
Carbon has four dots (2s2 2p2)
Stable molecule results when the
outermost (valence) shell is completed:
octet (eight dots) for main-group atoms
(two for hydrogen)
30
Number of Covalent Bonds to an
Atom


Atoms with one, two, or three valence
electrons form one, two, or three bonds
Atoms with four or more valence
electrons form as many bonds as they
need electrons to fill the s and p levels
of their valence shells to reach a stable
octet
31
Number of Covalent Bonds to
an Atom
32
Valences of Carbon

Carbon has four valence electrons (2s2 2p2),
forming four bonds (CH4)
33
Valences of Oxygen

Oxygen has six valence electrons (2s2 2p4)
but forms two bonds (H2O)
34
Valences of Nitrogen

Nitrogen has five valence electrons (2s2
2p3) but forms only three bonds (NH3)
35
Non-bonding electrons


Valence electrons not used in bonding are called
nonbonding electrons, or lone-pair electrons
Nitrogen atom in ammonia (NH3)Shares six valence
electrons in three covalent bonds and remaining two
valence electrons are nonbonding lone pair
36
Some simple molecules:
37
1.6 Valence Bond Theory

Covalent bond forms when two atoms
approach each other closely so that a
singly occupied orbital on one atom
overlaps a singly occupied orbital on the
other atom
38
1.6 Valence Bond Theory

Electrons are paired in the
overlapping orbitals and
are attracted to nuclei of
both atoms


H–H bond results from the
overlap of two singly
occupied hydrogen 1s
orbitals
H-H bond is cylindrically
symmetrical, sigma (s)
bond
39
Bond Energy


Reaction 2 H·  H2 releases 436 kJ/mol
Product has 436 kJ/mol less energy than two atoms:
H–H has bond strength of 436 kJ/mol. (1 kJ =
0.2390 kcal; 1 kcal = 4.184 kJ)
40
Bond Length



Distance between
nuclei that leads to
maximum stability
If too close, they
repel because both
are positively
charged
If too far apart,
bonding is weak
41
What is a Bond? An Energy
Minimum
42
1.7 Hybridization: sp3 Orbitals
and the Structure of Methane



Carbon has 4 valence electrons (2s2
2p2)
In CH4, all C–H bonds are identical
(tetrahedral)
sp3 hybrid orbitals: s orbital and
three p orbitals combine to form four
equivalent, unsymmetrical, tetrahedral
orbitals (sppp = sp3), Pauling (1931)
43
3
sp
Hybridization
44
Tetrahedral Structure of
Methane



sp3 orbitals on C overlap with 1s
orbitals on 4 H atom to form four
identical C-H bonds
Each C–H bond has a strength of
438 kJ/mol and length of 110 pm
Bond angle: each H–C–H is
109.5°, the tetrahedral angle.
45
Tetrahedral Structure of
Methane
46
1.8 Hybridization: sp3 Orbitals
and the Structure of Ethane


Two C’s bond to each other by s overlap of an sp3
orbital from each
Three sp3 orbitals on each C overlap with H 1s
orbitals to form six C–H bonds
47
1.8 Hybridization: sp3 Orbitals
and the Structure of Ethane



C–H bond
strength in
ethane 420
kJ/mol
C–C bond is 154
pm long and
strength is 376
kJ/mol
All bond angles
of ethane are
tetrahedral
48
1.9 Hybridization: sp2 Orbitals
and the Structure of Ethylene
49
1.9 Hybridization: sp2 Orbitals
and the Structure of Ethylene



sp2 hybrid
orbitals: 2s orbital
combines with two
2p orbitals, giving 3
orbitals (spp = sp2)
sp2 orbitals are in a
plane with120°
angles
Remaining p orbital
is perpendicular to
the plane
90
120
50
Bonds From sp2 Hybrid
Orbitals





Two sp2-hybridized orbitals overlap to form a
s bond
p orbitals overlap side-to-side to formation a
pi () bond
sp2–sp2 s bond and 2p–2p  bond result in
sharing four electrons and formation of C-C
double bond
Electrons in the s bond are centered between
nuclei
Electrons in the  bond occupy regions are on
either side of a line between nuclei
51
Bonds From sp2 Hybrid Orbitals
52
Structure of Ethylene




H atoms form s bonds with four sp2 orbitals
H–C–H and H–C–C bond angles of about 120°
C–C double bond in ethylene shorter and stronger
than single bond in ethane
Ethylene C=C bond length 133 pm (C–C 154 pm)
53
1.10 Hybridization: sp Orbitals
and the Structure of Acetylene

C-C a triple bond sharing six electrons
54
1.10 Hybridization: sp Orbitals
and the Structure of Acetylene



Carbon 2s orbital hybridizes with a single p orbital
giving two sp hybrids (two p orbitals remain
unchanged)
sp orbitals are linear, 180° apart on x-axis
Two p orbitals are perpendicular on the y-axis and
the z-axis
55
Orbitals of Acetylene


Two sp hybrid orbitals from each C form sp–sp s
bond
pz orbitals from each C form a pz–pz  bond by
sideways overlap and py orbitals overlap similarly
56
Bonding in Acetylene


Sharing of six electrons forms C C
Two sp orbitals form s bonds with hydrogens
57
Comparison of C-C & C-H bonds:
58
Comparison of C-H bonds:
Molecule
Bond
Energy (kcal)
Length (pm)
Ethane
Csp3-H
100
110
Ethylene
Csp2-H
106
108
Acetylene
Csp-H
132
106
59
Comparison of C-C bonds:
Molecule
Bond
Energy (kcal)
Length (pm)
Ethane
Csp3-Csp3
90
154
Ethylene
Csp2-Csp2
146
133
Acetylene
Csp-Csp
200
120
60
1.11 Hybridization of Nitrogen
and Oxygen




Elements other than C can
have hybridized orbitals
H–N–H bond angle in
ammonia (NH3) 107.3°
N’s orbitals (sppp) hybridize to
form four sp3 orbitals
One sp3 orbital is occupied by
two nonbonding electrons, and
three sp3 orbitals have one
electron each, forming bonds
to H
61
Hybridization of Oxygen in
Water



The oxygen atom is sp3-hybridized
Oxygen has six valence-shell electrons but forms only
two covalent bonds, leaving two lone pairs
The H–O–H bond angle is 104.5°
62
1.12 Acids and Bases: The
Brønsted–Lowry Definition


Brønsted–Lowry theory defines acids and
bases by their role in reactions that transfer
protons (H+) between donors and
acceptors.
“proton” is a synonym for H+ - loss of an
electron from H leaving the bare nucleus—a
proton. Protons are always covalently bonded
to another atom.
63
Brønsted Acids and Bases



“Brønsted-Lowry” is usually shortened to
“Brønsted”
A Brønsted acid is a substance that donates
a hydrogen ion, or “proton” (H+): a proton
donor
A Brønsted base is a substance that accepts
the H+: a proton acceptor
64
The Reaction of HCl with H2O



When HCl gas dissolves in water, a
Brønsted acid–base reaction occurs
HCl donates a proton to water
molecule, yielding hydronium ion
(H3O+) and Cl
The reverse is also a Brønsted acid–
base reaction of the conjugate acid and
conjugate base
65
The Reaction of HCl with
H2 O
66
67
Quantitative Measures of Acid
Strength



The equilibrium constant (Ke) for the reaction of an acid
(HA) with water to form hydronium ion and the conjugate
base (A-) is a measure related to the strength of the acid
Stronger acids have larger Ke
Note that brackets [ ] indicate concentration, moles per
liter, M.
68
Ka – the Acidity Constant



The concentration of water as a solvent does not change
significantly when it is protonated in dilute solution.
The acidity constant, Ka for HA equals Ke times 55.6 M
(leaving [water] out of the expression)
Ka ranges from 1015 for the strongest acids to very small
values (10-60) for the weakest
69
Ka – the Acidity Constant
70
1.13 Acid and Base Strength



The ability of a Brønsted acid to donate a
proton to is sometimes referred to as the
strength of the acid.
The strength of the acid can only be
measured with respect to the Brønsted base
that receives the proton
Water is used as a common base for the
purpose of creating a scale of Brønsted acid
strength
71
pKa – the Acid Strength Scale

pKa = -log Ka (in the same way that
pH = -log [H+]


The free energy in an equilibrium is
related to –log of Keq (DG = -RT log Keq)
A larger value of pKa indicates a
stronger acid and is proportional to the
energy difference between products
and reactants
72
pKa – the Acid Strength Scale

The pKa of water is 15.74
73
pKa – the Acid Strength Scale
74
1.14 Predicting Acid–Base Reactions
from pKa Values


pKa values are related as
logarithms to equilibrium constants
The difference in two pKa values is
the log of the ratio of equilibrium
constants, and can be used to
calculate the extent of transfer
75
Predicting Acid–Base Reactions from pKa
Values
76
Predicting Acid–Base Reactions from pKa
Values
77
Prob’s 2.14 & 2.15: Will these
reactions take place as written?
pKa = 36
pKa = 19
78
1.15 Organic Acids and Organic
Bases


The reaction patterns of organic
compounds often are acid-base
combinations
The transfer of a proton from a
strong Brønsted acid to a Brønsted
base, for example, is a very fast
process and will always occur along
with other reactions
79
Some organic acids:
80
Organic Acids


Those that lose a proton from O–H,
such as methanol and acetic acid
Those that lose a proton from C–H,
usually from a carbon atom next to
a C=O double bond (O=C–C–H)
81
Organic Acids
82
Conjugate bases:
83
Organic Bases



Have an atom with a lone pair of
electrons that can bond to H+
Nitrogen-containing compounds
derived from ammonia are the
most common organic bases
Oxygen-containing compounds can
react as bases when with a strong
acid or as acids with strong bases
84
Organic Bases
85
1.16 Acids and Bases: The Lewis
Definition


Lewis acids are electron pair acceptors; Lewis
bases are electron pair donors
The Lewis definition leads to a general description of
many reaction patterns but there is no quantitatve
scale of strengths as in the Brønsted definition of pKa
86
Lewis Acids and the Curved
Arrow Formalism

The Lewis definition of acidity includes metal
cations, such as Mg2+



They accept a pair of electrons when they
form a bond to a base
Group 3A elements, such as BF3 and AlCl3,
are Lewis acids because they have unfilled
valence orbitals and can accept electron pairs
from Lewis bases
Transition-metal compounds, such as TiCl4,
FeCl3, ZnCl2, and SnCl4, are Lewis acids
87
Lewis Acids and the Curved
Arrow Formalism


Organic compounds that undergo
addition reactions with Lewis bases
(discussed later) are called
electrophiles and therefore Lewis
Acids
The combination of a Lewis acid and a
Lewis base can shown with a curved
arrow from base to acid
88
Illustration of Curved Arrows in
Following Lewis Acid-Base Reactions
89
BF3 as a Lewis Acid:
BF3
O(CH3)2
F3B-O(CH3)2
90
Lewis Bases



Lewis bases can accept protons as well
as other Lewis acids, therefore the
definition encompasses that for
Brønsted bases
Most oxygen- and nitrogen-containing
organic compounds are Lewis bases
because they have lone pairs of
electrons
Some compounds can act as either
acids or bases, depending on the
reaction
91
Lewis Bases
92
Prob. : Identify acids & bases
93
Summary



Organic chemistry – chemistry of carbon compounds
Atom: positively charged nucleus surrounded by negatively
charged electrons
Electronic structure of an atom described by wave equation


Electrons occupy orbitals around the nucleus.
Different orbitals have different energy levels and different shapes



s orbitals are spherical, p orbitals are dumbbell-shaped
Covalent bonds - electron pair is shared between atoms
Valence bond theory - electron sharing occurs by overlap of
two atomic orbitals
94
Summary



Sigma (s) bonds - Circular cross-section and are formed by
head-on interaction
Pi () bonds – “dumbbell” shape from sideways interaction of
p orbitals
Carbon uses hybrid orbitals to form bonds in organic molecules.




In single bonds with tetrahedral geometry, carbon has four sp3
hybrid orbitals
In double bonds with planar geometry, carbon uses three
equivalent sp2 hybrid orbitals and one unhybridized p orbital
Carbon uses two equivalent sp hybrid orbitals to form a triple
bond with linear geometry, with two unhybridized p orbitals
Atoms such as nitrogen and oxygen hybridize to form strong,
oriented bonds

The nitrogen atom in ammonia and the oxygen atom in water are
sp3-hybridized
95
Summary





A Brønsted(–Lowry) acid donatea a proton
A Brønsted(–Lowry) base accepts a proton
The strength Brønsted acid is related to the 1 times the logarithm of the acidity constant,
pKa. Weaker acids have higher pKa’s
A Lewis acid has an empty orbital that can
accept an electron pair
A Lewis base can donate an unshared
electron pair
96