University of IsfahanUniversity of IsfahanDepartment of Department of

ChemistryChemistry

FUNDAMENTALS OF FUNDAMENTALS OF ORGANIC CHEMISTRYORGANIC CHEMISTRY

FOR BIOLOGY STUDENTSFOR BIOLOGY STUDENTS

Dr. Gholam Ali Koohmareh

اول 90-91نيمسال

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Text Book:

Fundamentals of Organic Chemistry

By: John. Mc. Murry

مباني شيمي آلي

نوشته جان مك موري

ترجمه دكتر عيسي ياوري

مورد در چند نصايحيآلي شيمي مطالعه

مفقوده ماده

در دانش در ساختار دانش ساختارآلي آلي شيمي شيمي

شبيه ساختارهااست متعادلي هرم

روي كهاست ايستاده نوك

ميشود خراب ساختار !!!كل

باشد شده گم چيزي .....اگر

كنيد مطالعه بايد مواد ساختار درك .براي

بار چندين نميشويد مسلط كه را هرمطلبيكنيد .تكرار

بحثهاي وارد و ميكنيم شروع ساده مباحث از ماميشويم تر .پيچيده

از قبل و زودتر چه هر نميشويد متوجه را مطلبي اگربگيريد كمك مطالب شدن ....انباشته

شيمي گروه در من اطاق در اداري ساعات طول در يا كالس از بعد ميتوانيدكنيد مطرح را خود درسي مشكل ايميل طريق از يا و

Email: Koohmareh@gmail.com

كنيد حل تمرين داشتيد وقت كه زماني .هر

است آن مستلزم اين و ميكند بيشتر را شما درك تمرينات انجامگيريد بكار ميگيريد ياد چه هر .كه

مفيدند بسيار فصل آخر .تمرينات

است اجباري ازمطالعه موثرتر بسيار فعال !يادگيري

ميگيريم %90 بكار و ميگوييم كه را چيزي

ميكنيم %70 صحبت و ميگوييم كه را چيزي

ميشنويم %50 و ميبينيم كه را چيزيميكنيم حفظ ما

ميبينيم %30 كه را چيزي

ميشنويم %20 كه را چيزي

ميخوانيم %10 كه را چيزي

روانشناسي ..… مطالعه يك

واقعي ولي واقعي عجيب ولي عجيب

فعال

اجباري

. پيوند و ساختار اول فصل

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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:

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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)

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Organic Chemistry

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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]

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1.1 Atomic Structure

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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

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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

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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”

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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)

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Shapes of Atomic Orbitals for Electrons s orbitals: spherical, nucleus at center p orbitals: dumbbell-shaped, nucleus at

middle

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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)

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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

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Orbitals and Shells

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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

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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).

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Electron Configurations

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Electron Configurations

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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 wedge indicates a bond is coming forward

Note that a dashed line indicates a bond is behind the page

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Tetrahedral Carbon:

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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)

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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)

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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

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Number of Covalent Bonds to an Atom

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Valences of Carbon Carbon has four valence electrons (2s2

2p2), forming four bonds (CH4)

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Valences of Oxygen Oxygen has six valence electrons (2s2 2p4) but forms

two bonds (H2O)

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Valences of Nitrogen Nitrogen has five valence electrons (2s2

2p3) but forms only three bonds (NH3)

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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

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Some simple molecules:

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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

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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 () bond

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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)

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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

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What is a Bond? An Energy Minimum

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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)

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sp3 Hybridization

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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.

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Tetrahedral Structure of Methane

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1.8 Hybridization: sp3 Orbitals and the Structure of Ethane

Two C’s bond to each other by overlap of an sp3 orbital from each

Three sp3 orbitals on each C overlap with H 1s orbitals to form six C–H bonds

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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

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1.9 Hybridization: sp2 Orbitals and the Structure of Ethylene

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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

90120

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Bonds From sp2 Hybrid Orbitals Two sp2-hybridized orbitals overlap to form a bond p orbitals overlap side-to-side to formation a pi () bond sp2–sp2 bond and 2p–2p bond result in sharing four

electrons and formation of C-C double bond Electrons in the bond are centered between nuclei Electrons in the bond occupy regions are on either side

of a line between nuclei

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Bonds From sp2 Hybrid Orbitals

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Structure of Ethylene H atoms form 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)

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1.10 Hybridization: sp

Orbitals and the Structure of Acetylene

C-C a triple bond sharing six electrons

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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

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Orbitals of Acetylene Two sp hybrid orbitals from each C form sp–sp bond pz orbitals from each C form a pz–pz bond by sideways

overlap and py orbitals overlap similarly

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Bonding in Acetylene Sharing of six electrons forms C C Two sp orbitals form bonds with hydrogens

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Comparison of C-C & C-H bonds:

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Comparison of C-H bonds:

Molecule BondEnergy (kcal)

Length (pm)

Ethane Csp3-H 100 110

Ethylene Csp2-H 106 108

Acetylene Csp-H 132 106

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Comparison of C-C bonds:

Molecule BondEnergy (kcal)

Length (pm)

Ethane Csp3-Csp3 90 154

Ethylene Csp2-Csp2 146 133

Acetylene Csp-Csp 200 120

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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

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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°

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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.

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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

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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

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The Reaction of HCl with H2O

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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.

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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

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Ka – the Acidity Constant

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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

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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 (G = -RT log Keq) A larger value of pKa indicates a

stronger acid and is proportional to the energy difference between products and reactants

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pKa – the Acid Strength Scale

The pKa of water is 15.74

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pKa – the Acid Strength Scale

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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

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Predicting Acid–Base Reactions from pKa Values

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Predicting Acid–Base Reactions from pKa Values

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Prob’s 2.14 & 2.15: Will these reactions take place as written?

pKa = 19

pKa = 36

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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

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Some organic acids:

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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)

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Organic Acids

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Conjugate bases:

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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

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Organic Bases

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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

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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

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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

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Illustration of Curved Arrows in Following Lewis Acid-Base

Reactions

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BF3 as a Lewis Acid:

BF3 O(CH3)2 F3B-O(CH3)2

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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

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Lewis Bases

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Prob. : Identify acids & bases

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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

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Summary Sigma () 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

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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