Organic chemistry I.

Organic chemistry I.

for Pharmacist Students

László Juhász, PhD

Department of Organic Chemistry,

University of Debrecen

The

range

of

chemicals

studied

in

organic

chemistry

include hydrocarbons (compounds containing only

carbon

and

hydrogen), as

well as myriad compositions based always on

carbon

, but also containing

other elements, especially

oxygen, nitrogen, sulfur, phosphorus

(these,

included in many organic chemicals in biology)

halogens.

Organometallic compounds

, i.e., involving alkali (e.g., lithium, sodium, and

potassium) or alkaline earth metals (e.g., magnesium).

Organic chemistry

Organic chemistry is the study of the

structure,

properties,

composition,

reactions, and preparation of

carbon-containing compounds.

The bonding patterns open to carbon, with its valence of four—formal single,

double, and triple bonds, as well as various structures with delocalized

electrons—make the array of organic compounds structurally diverse, and their

range of applications enormous.

They either form the basis of, or are important constituents of, many commercial

products including

pharmaceuticals;

petrochemicals

and products made from

them (including

lubricants, solvents, etc.);

plastics; fuels

and

explosives; etc. As

indicated, the study of organic chemistry overlaps with

organometallic chemistry

and biochemistry

, but also with

medicinal chemistry,

polymer chemistry

, as well as many aspects of materials science.

The huge number and variety of different organic compounds -

with different

physical and chemical properties

- is due to the different ways in which (a relatively

small number of) elements can link to each other,

no

due to the involvement of

many different elements.

More than 90 % of these molecules are organic compounds.

Chemical Abstracts Service (CAS), a web resource that

contains CAS Registry

Numbers A CAS Registry Number is a unique numerical identifier assigned by Chemical Abstracts Service (CAS) to every chemical substance described in the open scientific literature

What is organic chemistry?

Organic chemistry is the study of carbon compounds.

The doctrine: Vitalism

Originally, the science of organic chemistry was the study of

compounds extracted from natural organisms and their natural

products. It was Jöns

Jacob Berzelius

who in 1807 coined the term

“organic chemistry” for the study of

compounds derived from

natural sources. Berzelius, like almost everyone else at the time,

subscribed to the doctrine known as vitalism. Vitalism

held that

living systems possessed a

“vital force”

which was absent in

nonliving systems. Compounds derived from natural sources

(organic) were thought to be fundamentally different from inorganic

compounds; it was believed inorganic compounds could be

synthesized in the laboratory, but organic compounds could not—at

least not from inorganic materials.

The first chemist who transformed an inorganic compound to organic

compound was Friedrich W

öhler

(1828).

The transformation observed by Wöhler

was one in which an

inorganic

salt, ammonium cyanate, was converted to urea, a known

organic

substance earlier isolated from urine.

This experiment is now recognized as a scientific milestone, the first

step toward overturning the philosophy of vitalism.

The main content of organic chemistry as a discipline are:

•

Synthesis

–

how to design new molecules and prepare them

•

Reaction mechanism

–

how to find out how these molecules react

with each other and how to predict their reaction

•

Structure determination

–

how to find out the structure of new

compounds

•

Theoretical organic chemistry

–

how to understand these

structures in terms of atoms and the electrons that bind them

together

•

Biological chemistry

–

how to find out Nature does and how the

structure of biologically active molecules are related to what they

do

Organic compounds

nonsteroidal anti-inflammatory drug lipid-lowering agent semisynthetic opioid

essential structural component of all animal cell membranes

What will you learn?

•

Reaction types in organic chemistry

•

Hybridisation of atom orbitals of carbon

•

Differentiation of organic compounds on the

basis of the hybridisation of carbon atom

•

Functional groups

•

Transformation of functional groups

•

Reaction mechanism

Chemical Bonds forming Organic Compounds

As the reason for the extent and diversity of organic chemistry is the flexibility in

the ways carbon can bond with other elements, it is useful to note the following

about the bonding of carbon and other atoms, especially in organic chemical

compounds:

•

Carbon forms strong bonds - both with itself (other carbon atoms) and with other

elements (e.g., hydrogen, halogens, oxygen, nitrogen and others. This is true of

carbon bonds generally, including those forming both organic and inorganic

compounds.

•

Organic molecules are bound by covalent bonds,

rather than ionic bonds

.

Understanding the ways in which atoms join to other atoms to form organic

compounds, i.e. the chemical "bonds" involved, is important in the study of organic

chemistry and is extremely useful in order to understand what is organic chemistry.

•

Covalency in Organic Compounds:

In organic compounds,

•

carbon exhibits a covalency of 4,

•

nitrogen (usually) 3,

•

oxygen 2 and

•

hydrogen 1 and the

•

halogens each exhibit a covalency of 1.

•

Carbon-Carbon Bonds: Two carbon atoms can be linked together by either:

•

a single bond (involving one shared electron pair),

•

a double bond (involving two shared electron pairs), or

•

a triple bond (involving three shared electron pairs).

•

Carbon can also form:

•

double bonds (involving two shared electron pairs) with oxygen, and

•

depending on the situation, double bonds (involving two shared electron

pairs) or triple bonds (involving three shared electron pairs) with nitrogen.

•

Many carbon atoms can link together in very long chains forming huge

molecules. This type of organic molecule forms substances called polymers.

Polymer Chemistry is a large subject area that includes the structures,

properties, and behaviours of plastics.

•

Carbon atoms can link together forming rings. This type of structure results in

particular categories of organic compound, i.e. aromatic compounds.

bond E (kJ/mol) C-C 346 N-N 167 O-O 142 S-S 226 Si-Si 176 Si-O 368 C-H 411

Stable C-C bond



chains, ring

n_C C_nH_2n+2 C_nH_2n+1-X 3 1 2 5 3 9 7 9 74 10 75 1553 20 366.319 82.299.275 40 62.498.178.805.831

-Huge variety, complex molecules

Organic compounds: C, H + and other atoms

Organogen elements: N, O, S, P, Hlg (which are especially

characteristic ingredients of organic compounds)

Elem organic compounds: not organic elements in the C-X bond

Organometallic compounds: metal-carbon bond (C-M) rapidly developing research field

Organic chemistry

–

chemistry of functional groups

Functional group:

an atom or a group of atoms that determines the chemical and

physical properties of a given compound



R-X

Chemical bonds

Phenomena:

A • + • B → A ―

B

Properties:

Bond energy: the amount of energy per mole that releases when a bond is formed – reverse: bond

dissociation energy: The amount of energy required to dissociate a molecule to two separate atoms.

Bond length: distance between atoms

Atoms in a bond – homonuclear / heteronuclear bond

(C-C / C-Br)

Separated atoms Atoms are

bonded

Atoms combine with one another to give compounds

having properties different from the atoms they contain. The attractive force between atoms in a compound is a chemical bond.

Types of chemical bonds I.

Plot of potential energy versus distance for A and B atoms. At long distances, there is a weak attractive force. As the distance decreases, the potential energy decreases, and the system becomes more stable

because each electron now feels the attractive force of two protons

rather than one. The optimum

distance of separation corresponds to the normal bond distance of an A-B molecule. At shorter distances, nucleus–nucleus and electron–

electron repulsions are greater than electron–nucleus attractions, and the system becomes less stable.

One type of chemical bond, called an ionic bond, is the force of attraction between

oppositely charged species (ions). Ions that are positively charged are referred to as cations; those that are negatively charged are anions. Attracting force: electrostatic attraction.

Another type of chemical bond, called covalent bond, is the chemical bond between two atoms that results from their sharing of two electrons.

A coordinate covalent bond, also known as a dative bond or coordinate bondis a kind of 2-center, 2-electron covalent bondin which the two electrons derive from the same atom.

Types of chemical bonds II.

Carbon monoxide consists of one carbon atom and

one oxygen atom, connected by a triple bond that consists of two covalent bonds as well as one dative covalent bond.

Formation of anadductof ammonia and boron trifluoride, involving formation of a coordinate covalent bond.

The covalent, or shared electron pair, model of chemical bonding was first suggested

by G. N. Lewis of the University of California in 1916. Lewis proposed that a sharing of two electrons by two hydrogen atoms permits each one to have a stable closed-shell (noble gas) electron configuration analogous to helium.

Models for describing the chemical bond

The shared electron pair model of chemical bonding

Structural formulas of this type in which electrons are represented as dots are called Lewis structures.

Covalent bonding in F₂ gives each fluorine 8 electrons in its valence shell and a stable electron configuration equivalent to that of the noble gas neon:

The Lewis model limits second-row elements (Li, Be, B, C, N, O, F, Ne) to a total

of 8 electrons (shared plus unshared) in their valence shells. Hydrogen is limited to 2.

Octet rule: in forming compounds the elements gain, lose, or share electrons to give a stable electron configuration characterized by eight valence electrons.

When the octet rule is satisfied for carbon, nitrogen, oxygen, and fluorine, they have an electron configuration analogous to the noble gas neon.

Lewis model of the organic compounds methane and carbon tetrafluoride.

Carbon has 8 electrons in its valence shell in both methane and carbon tetrafluoride. By

forming covalent bonds to four other atoms, carbon achieves a stable electron configuration analogous to neon.

Representing a 2-electron covalent bond by a dash (—), the Lewis structures for hydrogen fluoride, fluorine, methane, and carbon tetrafluoride become:

DOUBLE BONDS AND TRIPLE BONDS

Lewis’s concept of shared electron pair bonds allows for 4-electron double bonds and

6-electron triple bonds. Carbon dioxide (CO₂) has two carbon–oxygen double bonds, and the octet rule is satisfied for both carbon and oxygen. Similarly, the most stable Lewis structure for hydrogen cyanide (HCN) has a carbon–nitrogen triple bond.

Ethylene (C

H

) contains a carbon

–

carbon double bond in its most stable Lewis

structure, and each carbon has a completed octet. The most stable Lewis structure

for acetylene (C

H

) contains a carbon

–

carbon triple bond. Here again, the octet

rule is satisfied.

POLAR COVALENT BONDS AND ELECTRONEGATIVITY

Electrons in covalent bonds are not necessarily shared equally by the two atoms that they connect. If one atom has a greater tendency to attract electrons toward itself than the other, we say the electron distribution is polarized, and the bond is referred to as a polar covalent bond.

Hydrogen fluoride, for example, has a polar covalent bond. Because fluorine attracts electrons more strongly than hydrogen, the electrons in the H-F bond are pulled

toward fluorine, giving it a partial negative charge, and away from hydrogen giving it a partial positive charge. This polarization of electron density is represented in various ways.

Electronegativity:

Electronegativityincreases across a row in the periodic table. The most electronegative of the second-row elements is fluorine; the most electropositive is lithium. Electronegativity decreases in going down a column. Fluorine is more electronegative than chlorine. The most commonly cited electronegativity scale was devised by Linus Pauling.

Centers of positive and negative charge that are separated from each other constitute a dipole.

The dipole moment (μ) of a molecule is equal to the charge e (either the positive or the negative charge, since they must be equal) multiplied by the distance between the centers of charge:

The dipole moment

[μ] = 1 D (Debye) = 3.346 x 10-30 _{C*m (coulomb*meter)}

μ =

e

x

d

A polar molecule has a dipole moment, a nonpolar one does not. Thus, all of the hydrogen halides are polar molecules. In order to be polar, a molecule must have polar bonds, but

can’t have a shape that causes all the individual bond dipoles to cancel.

The molecular dipole moment is the resultant of all of the individual bond dipole moments of a substance. Some molecules, such as carbon dioxide, have polar bonds, but lack a dipole moment because their shape causes the individual C=O bond dipoles to cancel.

CCl

The quantum mechanical model of the atom

•

Louis de Broglie proposed that all particles could be treated as matter

waves.

•

Erwin Schrödinger proposed the quantum mechanical model of the atom,

which treats electrons as matter waves.

•

Schrödinger's equation

H

ψ

=

Eψ

, can be solved to yield a series of wave

function

ψ

(psi), each of which is associated with an electron binding

energy, E.

•

The square of the wave function,

ψ

_{​​, represents the}

_probability

_{of finding an}

electron in a given region within the atom.

•

An atomic orbital is defined as

the region within an atom that encloses

where the electron is likely to be 90% of the time.

•

The Heisenberg uncertainty principle states that we can't know both the

energy and position of an electron. Therefore, as we learn more about the

electron's position, we know less about its energy, and vice versa.

•

Electrons have an intrinsic property called spin, and an electron can have

one of two possible spin values: spin-up or spin-down.

•

Any two electrons occupying the same orbital must have opposite spins.

Key points:

Schrödinger's equation

On a very simple level, we can think of electrons as standing matter waves

that have certain allowed energies. Schrödinger formulated a model of the

atom that assumed the electrons could be treated at matter waves. The basic

form of Schrödinger's wave equation is as follows:

H

ψ

=

Eψ

ψ is called a wave function, His known as the Hamiltonian operator; and E is the binding energy of the electron.

Solving Schrödinger's equation yields multiple wave functions as solutions,

each with an allowed value for

E.

The wave functions that are derived from Schrödinger's equation for a

specific atom are also called

atomic orbitals. Chemists define an atomic

orbital as

the region within an atom that encloses where the electron is likely

to be 90% of the time.

The square of a wave function is proportional to the

probability

of finding an

electron in a particular volume of space within an atom. The function

ψ

_is

often called the

probability density.

Shapes of atomic orbitals

•

The s orbitals are spherical

•

The p orbitals are shaped like dumbbells that are oriented along one of the

axes

—

x, y, z.

•

The d orbitals can be described as having a clover shape with four possible

orientations

—

with the exception of the d orbital that almost looks like a p

orbital with a donut going around the middle.

A

nodal plane

is always a plane where electrons cannot be.

The wave function changes sign at the nucleus. The yz-plane is a nodal surface

for the 2p

orbital. The probability of finding a 2p

electron in the yz-plane is

zero. Analogously, the xz-plane is a nodal surface for the 2p

orbital, and the

xy-plane is a nodal surface for the 2p

orbital.

Electron orbitals that have the same energy levels are called degenerate

orbitals. (e. g. p

-p

- p

)

The Pauli Exclusion Principle states that, in an atom or molecule, no two

electrons can have the same four electronic quantum numbers. As an orbital

can contain a maximum of only two electrons, the two electrons must have

opposing spins. This means if one is assigned an up-spin ( +1/2), the other must

be down-spin (-1/2).

Aufbau Principle

The electrons are added one by one to the various orbitals in order of their

increasing energy starting with the orbital of lowest energy. The increasing order

of energy of various orbitals is

1s<2s<2p<3s<3p<4s<3d<3d<5s<5p<6s<4f<5d<6p<7s<5f<6d<7p

……

Important Rules for Filling Orbitals

Hund’s Rule of maximum multiplicity

This rule deals with the filling of electrons in the orbitals having equal energy

(degenerate orbitals). According to this rule,

“Electron

pairing in p, d and f

orbitals cannot occur until each orbitals of a given subshell contains one

electron each or is singly

occupied”

.

Quantum mechanical description of electrons and orbitals

Phenomena: - particle-wave dualism of electron (de Broglie, 1923) - wave equation (Schrödinger, 1926)

Electrons as waves :

- wave function: ψ = f(x,y,z) and

| ψ |2 _{~ the probabilityof being in a certain place (in x,y,z point) for electrons}

- univalent and continuous function

-possible values: positive, negative or zero (node)

Schrödinger equation: Hψ = Eψ H is called Hamiltonian operator

Atomic orbital (AO) – an orbital represents the space where an electron spends 90%-95% of its time. Empty orbital ~ unoccupied orbital; orbital with electron ~ occupied orbital

Rulesfor occupying orbitals:

the lowest E orbitalis filled first (1s vs 2s), Pauli exclusion principle (2e-_{/orbital with opposite}

spin) and Hund’s rule (the greatest number of unpaired electrons on degenerate orbitals)

Atomic orbitals (Summary)

Lewis proposed his shared electron-pair model of bonding in 1916, almost a decade

before de Broglie’s theory of wave–particle duality. De Broglie’s radically different view of

an electron, and Schrödinger’s success in using wave equations to calculate the energy of an electron in a hydrogen atom, encouraged the belief that bonding in molecules could be

explained on the basis of interactions between electron waves. This thinking produced two widely used theories of chemical bonding: one is called the valence bond model, the other the molecular orbital model.

The valence bond and molecular orbital theories differ in how they use the orbitals of two atoms to describe the orbital that contains the electron pair in the molecule.

One property of waves that is important here is called “interference” in physics. Constructive interference occurs when two waves combine so as to reinforce each other (“in phase”);

destructive interference occurs when they oppose each other(“out of phase”). In the valence

bond model constructive interference between two electron waves is seen as the basis for the shared electron-pair bond. In the molecular orbital model, the wave functions of

molecules are derived by combining wave functions of atoms.

BONDING IN H

: THE VALENCE BOND MODEL

The characteristic feature of valence bond theory is that it describes a covalent bond between two atoms in terms of an in-phase overlap of a half-filled orbital of one atom with a half-filled orbital of the other, illustrated for the case of H₂.

Two hydrogen atoms, each containing an electron in a 1s orbital, combine so that their orbitals overlap to give a new orbital associated with both of them. In-phase orbital

overlap (constructive interference) increases the probability of finding an electron in the region of overlap.

Valence bond picture of bonding in H₂. Overlap of half-filled 1s orbitals of two hydrogen atoms gives a new orbital encompassing both atoms. This new

orbital contains the two original electrons. The electron density (electron probability) is highest in the region between the two atoms. The black dots correspond to the nuclei, and the + signs to the signs of the wave functions. When the wave functions are of the same sign, constructive

interference leads to an increase in the probability of finding an electron in the region where the two orbitals overlap.

Generating MO-s from AO-s - OVERLAPPING

Settlement principals:

✓ maximal overlapping, combined AO-s must have the same phase

✓ forbidden: overlapping of AO-s with opposite phase

✓ different AO-s (e.g. s and p) can form MO-s

π

-bond:

(perpendicular) to the line connecting the nuclei, there is a node between the nuclei!

s-bond:

Overlap of two p orbitals along internuclear axis gives a σbond.

rotational symmetry, no

node between the nuclei!

BONDING IN H

: THE MOLECULAR ORBITAL MODEL

The molecular orbital approach to chemical bonding is based on the notion that, as

electrons in atoms occupy atomic orbitals, electrons in molecules occupy molecular orbitals. In the molecular orbital method this is accomplished by representing molecular orbitals as combinations of atomic orbitals, the linear combination of atomic orbitals-molecular

orbital (LCAO-MO) method.

Take H₂ for example.

Two molecular orbitals (MOs) are generated by combining the 1s atomic orbitals (AOs) of two hydrogen atoms. In one combination, the two wave functions are added; in the other they are subtracted.

The two new orbitals that are produced:

The additive combination generates a bonding orbital; the subtractive combination generates an antibonding orbital. Both the bonding and antibonding orbitals have rotational symmetry around the line connecting the two atoms; they have σ

symmetry. The two are differentiated by calling the bonding orbital σ and the

antibonding orbital σ* (“sigma star”). The bonding orbital is characterized by a region of high electron probability between the two atoms, and the antibonding orbital has a nodal surface between them.

A molecular orbital diagram for H

When assigning electrons to MOs, the same rules apply as for writing electron

configurations of atoms. Electrons fill the MOs in order of increasing orbital energy, and the maximum number of electrons in any orbital is 2. The 2 electrons of H₂ occupy the bonding orbital, have opposite spins, and both are held more strongly than they would be in separated hydrogen atoms. There are no electrons in the antibonding orbital.

The customary format shows the starting AOs at the left and right sides and the MOs in the middle. It must always be true that the number of MOs is the same as the number of AOs that combine to produce them. Thus, when the 1s AOs of two hydrogen atoms combine, two MOs result. The bonding MO (σ) is lower in energy and the antibonding MO (σ *) higher in energy than either of the original 1s orbitals.

For a molecule as simple as H₂, it is hard to see much difference between the valence bond and molecular orbital methods. The most important differences appear in molecules with more than two atoms—a very common situation indeed. In those cases, the valence bond method continues to view a molecule as a collection of bonds between connected atoms. The molecular orbital method, however, leads to a picture in which the same electron can be associated with many, or even all, of the atoms in a molecule. (Basic proposal: In a

molecule all electron moves in the field of all nuclei of the molecule  polycentric MO-s (delocalized electrons))

In molecular orbital theory, molecular orbitals (MOs) can be calculated as the linear combination of atomic orbitals. In MO theory, the molecular orbitals (MOs) are

approximated as sums of all the atomic orbitals (AOs) on all the atoms; there are as many MOs as AOs. Each AO_i has a weightingcoefficient c_i that indicates the AO's contribution to a particular MO.

Mathematical equitation: Ψ = c₁ψ₁ + c₂ψ₂ + c₃ψ₃ + … + c_iψ_i

c_i : weighting coefficient; ψ: wave function

BUT! It works perfectly for diatomic molecules – for more atoms approximate methods has to be applied (searching E minimum by iterative path)

Polycentric molecules: e.g. conjugate

molecules

• different E levels with hierarchical orientation

• filling up MO-s with electrons –

according to rising E levels

• Excitation, radical formation: HOMO

→ LUMO transition

e. g. UV spectra p-p* transition in conjugate molecules

(1,3-butadiene)

Possible simplifying and neglects:

1. inner, completed shells can be described with AO-s

2. molecules that could be described with Lewis-Langmuir formula: the non-bonding, unshared electrons MO’s ~ isolated atom AO’s

3. molecules that could be described with Lewis-Langmuir formula: bonding, shared electrons could be well approximated by delocalized, bicentered MO-s

All other cases polycentric MO-s has to be applied (conjugation, aromatic systems, M-effect)

BONDING IN METHANE (CH

) AND ORBITAL HYBRIDIZATION

Since covalent bonding requires the overlap of half-filled orbitals of the connected atoms, carbon with an electron configuration of 1s2_2s2_2p

x1 2py1 has only two half-filled orbitals.

How can it have bonds to four hydrogens?

(a) Electron configuration of carbon in its most stable state.

(b) An electronis “promoted”

from the 2s orbital to the vacant 2p orbital.

(c) The 2s orbital and the three 2p orbitals are combined to give a set of four equal-energy sp3

-hybridized orbitals, each of which contains one electron.

In the 1930s Linus Pauling offered an ingenious solution to it.

He began with a simple idea: “promoting” one of the 2s electrons to the empty 2p_z orbital gives four half-filled orbitals and allows for four C-H bonds. The electron configuration that results (1s2 _2s1 _2p

x 1 2py 1 2pz 1), however, is inconsistent with the fact that all of these bonds

are equivalent and directed toward the corners of a tetrahedron.

(The formed bonds are not equivalent and the bond angles are different These are not in correlation with the experimental results

(4 equivalent bonds, 109.5o₎₎

The second part of Pauling’s idea was novel: mix together (hybridize) the four valence orbitals of carbon (2s, 2p_x, 2p_y, and 2p_z) to give four half-filled orbitals of equal energy

The four new orbitals in Pauling’s scheme are called sp3 _{hybrid orbitals because they come} from one s orbital and three p orbitals.

(a) Electron configuration of carbon in its most stable state.

(b) An electronis “promoted”

from the 2s orbital to the vacant 2p orbital.

(c) The 2s

orbital and the three 2p orbitals are combined to give a set of four equal-energysp 3-hybridized orbitals, each of which contains one electron.

A tetrahedral arrangement of four σbonds is characteristic of sp3_{-hybridized carbon.}

The sp3 _{hybrid orbitals are arranged in a tetrahedral fashion around carbon. Each orbital}

contains one electron and can form a bond with a hydrogen atom to give a tetrahedral methane molecule.

bond angle

Mixing of ones orbital with three p orbitals generates foursp3 _{hybrid orbitals. Each sp}3

hybrid orbital has 25% scharacter and 75% p character. The four sp3 _{hybrid orbitals have}

their major lobes directed toward the corners of a tetrahedron, which has the carbon atom at its center. The bond angles are 109.5°.

s + 3p  sp3

sp

_{HYBRIDIZATION AND BONDING IN ETHYLENE (C}

H

)

sp

_{HYBRIDIZATION AND BONDING IN ETHANE (C}

2

H

)

The orbital hybridization model of covalent bonding is readily extended to carbon–

carbon bonds. Ethane is described in terms of a carbon–carbon σ bond joining two CH₃ (methyl) groups. Each methyl group consists of an sp3_{-hybridized carbon attached to three}

hydrogens by sp3_{–1s bonds. Overlap of the remaining half-filled orbital of one carbon with}

that of the other generates a σbond between them. Here is a third kind of σ bond, one that has as its basis the overlap of two sp3_{-hybridized orbitals.}

In general, you can expect that carbon will be sp3_{-hybridized when it is directly bonded to}

four atoms.

Ethylene is a planar molecule. Because sp3 _{hybridization is associated with a tetrahedral}

geometry at carbon, it is not appropriate for ethylene, which has a trigonal planar geometry at both of its carbons.

All the atoms of ethylene lie in the same plane. All the bond angles are close to 120°, and the carbon–carbon bond distance is significantly shorter than that of ethane.

The hybridization scheme is determined by the number of atoms to which the carbon is directly attached.

In ethane, four atoms are attached to carbon by bonds, and so four equivalent sp3 _hybrid orbitals are required.

In ethylene, three atoms are attached to each carbon, so three equivalent hybrid orbitals are required. These three orbitals are generated by mixing the carbon 2s orbital with two of the 2p orbitals and are called sp2 _{hybrid orbitals. One of the 2p orbitals is left unhybridized.}

(c) The 2s orbital and two of the three 2p orbitals are combined to give a set of three equal-energy sp2_-hybridized

orbitals. One of the 2p orbitals remains unchanged.

(a) Electron configuration of carbon in its most stable state.

(b) An electron is “promoted”

from the 2s orbital to the vacant 2p orbital.

The three sp2 _{orbitals are of equal energy; each has one-third s character and two-thirds p}

character. Their axes are coplanar, and each has a shape much like that of an sp3 _orbital.

Each carbon of ethylene uses two of its sp2 _{hybrid orbitals to form σ bonds to two}

hydrogen atoms. The remaining sp2 _{orbitals, one on each carbon, overlap along the}

internuclear axis to give a σ bond connecting the two carbons.

Each carbon atom still has, at this point, an unhybridized 2p orbital available for bonding. These two half-filled 2p orbitals have their axes perpendicular to the framework of σ

bonds of the molecule and overlap in a side-by-side manner to give what is called a pi (π) bond. According to this analysis, the carbon–carbon double bond of ethylene is viewed as a combination of a σ bond plus a π bond. The additional increment of bonding makes a carbon–carbon double bond both stronger and shorter than a carbon–carbon single bond. Electrons in a πbond are called πelectrons. The probability of finding a π electron

is highest in the region above and below the plane of the molecule. The plane of

the molecule corresponds to a nodal plane, where the probability of finding a πelectron is zero.

In general, you can expect that carbon will be sp2_{-hybridized when it is directly}

s + 2p + p_z sp2_{+ p} z1

sp HYBRIDIZATION AND BONDING IN ACETYLENE (C₂H₂)

One more hybridization scheme is important in organic chemistry. It is called sp

hybridizationand applies when carbon is directly bonded to two atoms, as it is in acetylene.

Since each carbon in acetylene is bonded to two other atoms, the orbital hybridization model requires each carbon to have two equivalent orbitals available for the formation of σ bonds. According to this model the carbon 2s orbital and one of the 2p orbitals

combine to generate a pair of two equivalent sp hybrid orbitals. Each sphybrid orbital has 50% s character and 50% p character. These two sp orbitals are oriented at an angle of

180° to each other. Two of the original 2p orbitals remain unhybridized. Their axes are perpendicular to each other and to the common axis of the pair of sp hybrid orbitals.

(c) The 2s orbital and one of the three 2porbitals are combined to give a set of two equal-energy

sp-hybridized orbitals. Two of the 2p orbitals remain unchanged. (b) An electronis “promoted”

from the 2s orbital to the vacant 2p orbital.

(a) Electron configuration of carbon in its most stable state.

s + p + p_y+ p_z  sp + p_y1_{+ p} z1

The two carbons of acetylene are connected to each other by a 2sp–2sp σ bond, and each is attached to a hydrogen substituent by a 2sp–1s σ bond. The unhybridized 2p orbitals on one carbon overlap with their counterparts on the other to form two π bonds. The carbon–

carbon triple bond in acetylene is viewed as a multiple bond of the σ + π+ π type.

In general, you can expect that carbon will be sp-hybridized when it is directly bonded to two atoms.

bond d (nm) bond d (nm) bond d (nm) C-C (s) sp3_-sp3 sp3_-sp2 sp3_-sp sp2_-sp2 sp2_-sp sp-sp 1.54 1.51 1.47 1.48 1.43 1.38 C=C (p) sp2_-sp2 sp2_-sp C≡C (p) sp-sp 1.32 1.31 1.18 C-H (s) sp3_-H sp2_-H sp-H 1.09 1.08 1.08

Hybridisation – describing single, double and triple bonds (primary for the elements of the second row of the period!!!) In the third row of the periodic table weaker multiple bonds (smaller overlapping).

Bond distance (bond order) depends on hybridisation of the connected atoms.

WHICH THEORY OF CHEMICAL BONDING IS BEST?

Organic chemists use all three - Lewis structures, orbital hybridization, and molecular orbital descriptions - emphasizing whichever one best suits a particular feature of structure or reactivity.

RESONANCE

When writing a Lewis structure, we restrict a molecule’s electrons to certain well-defined locations, either linking two atoms by a covalent bond or as unshared electrons on a single atom. Sometimes more than one Lewis structure can be written for a molecule, especially those that contain multiple bonds.

However both bond distances in ozone are exactly the same (128 pm) — somewhat shorter than the single bond distance and somewhat longer than the double bond distance. The structure of ozone requires that the central oxygen must be identically bonded to both terminal oxygens.

The following Lewis structure for ozone satisfies the octet rule; all three oxygens have 8 electrons in their valence shell.

This Lewis structure, however, doesn’t accurately portray the bonding in ozone, because the two terminal oxygens are bonded differently to the central oxygen. The central oxygen is depicted as doubly bonded to one and singly bonded to the other. Since it is generally true that double bonds are shorter than single bonds.

According to the resonance concept, when more than one Lewis structure may be written for a molecule, a single structure is not sufficient to describe it. Rather, the true structure has an electron distribution that is a “hybrid” of all the possible Lewis structures that can

be written for the molecule. In the case of ozone, two equivalent Lewis structures may be written. We use a double-headed arrow to represent resonance between these two Lewis structures.

It is important to remember that the double-headed resonance arrow does not indicate a process in which the two Lewis structures interconvert.

Resonance attempts to correct a fundamental defect in Lewis formulas. Lewis formulas show electrons as being localized; they either are shared between two atoms in a

covalent bond or are unshared electrons belonging to a single atom.

In reality, electrons distribute themselves in the way that leads to their most stable

arrangement. This sometimes means that a pair of electrons is delocalized, or shared by several nuclei.

What we try to show by the resonance description of ozone is the delocalization of the lonepair electrons of one oxygen and the electrons in the double bond over the three atoms of the molecule.

Organic chemists often use curved arrows to show this electron delocalization.

Alternatively, an average of two Lewis structures is sometimes drawn using a dashed line

to represent a “partial” bond.

In the dashed-line notation the central oxygen is linked to the other two by bonds that are halfway between a single bond and a double bond, and the terminal oxygens each bear one half of a unit negative charge.

The Most Important Rules of Resonance

1. Atomic positions

(connectivity) must be the same in all resonance

structures; only the electron positions may vary among the various contributing structures.

2. Lewis structures in which second-row elements own or share more than 8 valence electrons are especially unstable and make no contribution to the true structure. (The octet rule may be exceeded for

elements beyond the second row.)

3. When two or more structures satisfy the octet rule, the most stable one is the one with the smallest separation of oppositely charged atoms.

4. Among structural formulas in which the octet rule is satisfied for all atoms and one or more of these atoms bears a formal

charge, the most stable

resonance form is the one in

which negative charge resides on the most electronegative atom (or positive charge on the most electropositive one).

5. Each contributing Lewis structure must have the same number of electrons and the same net charge, although the formal charges of individual atoms may vary among the various Lewis structures.

6. Each contributing Lewis structure must have the same number of unpaired electrons.

7. Electron delocalization

stabilizes a molecule. A molecule in which electrons are delocalized is more stable than implied by any of the individual Lewis structures that may be written for it. The degree of stabilization is greatest when the contributing Lewis

structures are of equal stability.

Electron delocalization lowers the potential energy of the substance and thus makes it more stable than any of the contributing structures. The difference between the potential energy of the actual structure and that of the contributing structure with the lowest

Resonance in Valence Bond (VB) Theory

Mathematical formula: Ψ = c₁φ₁ + c₂φ₂ + c₃φ₃ + … + c_iφ_i

Mathematically similar to LCAO-MO but in this case

Φ_i: wave function for each resonance structure, c_i weighting coefficient for each resonance structure

POLAR COVALENT BONDS AND ELECTRONEGATIVITY

Electrons in covalent bonds are not necessarily shared equally by the two atoms that they connect. If one atom has a greater tendency to attract electrons toward itself than the other, we say the electron distribution is polarized, and the bond is referred to as a polar covalent bond.

Electronegativity:

Different scales – the most known: Pauling (arbitrary base: EN_F = 4.0) Mostly for diatomic molecules but it is extensible for groups.

The electronegativity of C depends on the hybridisation state Larger s character  stronger nuclear attraction

EN_C(sp)> EN_C(sp2

) > ENC(sp3)

s character: 50% > 33% > 25% Csoport EN Csoport EN CH₃ CH₃CH₂ CH₂Cl 2.472 2.482 2.538 C₆H₅ CF₃ CN 2.717 2.985 3.208 CCl₃ 2.666 NO₂ 3.412 Group Group

Inductive effect (-I, +I), conjugation

Different EN  charge separation, dipole moment appears

Molecules with more atoms connected with polarized bonds: vector addition  permanent dipole moment.

Effect of an atom / a group to the σ-bond connected to the carbon – inductive (I) effect

Types: electron donating group (EDG) (+I), electron withdrawing group (EWG) (-I)

Most of the atoms/groups have -I effect (DEN)

Exceptions: metals, negatively charged atoms/ groups, alkyl groups (hyper conjugation)

+I -I O -COO -CH₃ CR₃ NH₃+ NR₃+ NH₂ NO₂ CN COOH Hlg OH OR COR SH C≡CR Ar CH=CR₂ SO₂R

The electron-donating or electron-withdrawing effect of a group that is transmitted through σ bonds is called an inductive effect.

Conjugation

Possibility of side-to-side overlapping of all p_z orbitals, electron delocalisation

Result of conjugation: lower energy level, energy gain

Mezomeric effect

Interaction of p_z orbitals of sp2 _{atoms with another p}

z orbital or NBMO, a p-bond and a pz

orbital of a substituent – mesomeric (M) effect (always spreads above the σ-bond system!) Types: electron donating (+M), electron withdrawing (-M). The sign (+/-) belongs to the substituent!

Interaction of p_z orbitals of sp2 _{carbon atoms with the nonbonding electron pair of a}

substituent (+M effect for the substituent)

LCAO-MO: 7 centered MO (7 MOs), 8 delocalised electrons

• Interaction of p_z orbitals of sp2 _{atoms with the empty p}

z orbital of a substituent (-M

effect for the substituent)

In both cases electrons moved from their original position and spread in a

larger space that results energy gain –

it is a stabilizing factor!

Hyperconjugation

Hyperconjugation– the explanation of the electron donating effects of alkyl groups, this

effect can be observed in experiments. The interaction of the electron pair of s-bond of C-H bond with an adjacent π-electron pair or empty p_z orbitals to give an extended MO that

increases the stability of the system– stabilising effect!! Carbenium ion: electron sextet, sp2 _hybridised

carbon with empty p_zorbital (carbocation)

Hyperconjugation plays role in the enhancement of the stability of electron deficient carbons e.g.: carbocations, radicals.

The interaction of the electron pair of s-bond of C-H bond with an adjacent π-electron pair

The interaction of the electron pair of s-bond of C-H bond with an empty p_z orbitals

We use the same terminology for carbocations. A primary carbocation is attached to one other carbon, a secondary to two, and a tertiary to three.

The terminology of carbon-containing functional groups: primary, secondary,

tertiary, quaternary

•Primary carbons, are carbons attached to oneother carbon. (Hydrogens – although usually

3 in number in this case – are ignored in this terminology).

•Secondary carbons are attached to two other carbons.

•Tertiary carbons are attached to three other carbons.

•Finally, quaternary carbons are attached to four other carbons. The name depends on the number of carbons directly attached to the red carbon The name depends on the number of carbons directly attached to the carbocations

Properties:min. 100 times weaker bonds compare to the first order bonds (ionic, covalent, dative), intermolecular (sometimes intramolecular!) interactions

Responsible for the physical properties of the compounds (melting point (mp), boiling point (bp), solubility, and also for the structures of macromolecules (e.g. DNA).

Hydrogen bond (hydrogen bridge) – an attractive interaction

between a hydrogen bonded to an electronegative atom (X = N, O, F, Cl) and an unshared electron pair on another electronegative atom ( Y = N, O, F, Cl.) X and Y atoms with high electronegativity

Orinteraction of the hydrogen of a polarised X-H bond with a nonbonding e-pair of Y atom.

Typical values: E: 12-30 kJ/mol (exception: HF…FƟ_{, 162 kJ/mol)}

Noncovalent interactions

Van der Waals-forces – Attractive forces between neutral species (atoms or molecules, but not ions) are referred to as van der Waals forces

Interactions without electron transfer

Properties: „non-saturable”, nonspecific, work on greater distance, low E (5-15 kJ/mol), E ~ r-6

These forces are electrical in nature, and in order to vaporize a substance, enough energy must be added to overcome them.

Three types of it

dipole-dipole

dipole-induced dipole

Dipole-dipole interaction

Polar molecules engage in dipole–dipole and dipole/induced-dipole attractions.

The dipole–dipole attractive force is easiest to visualize and is illustrated as follows

Two molecules of a polar substance are oriented so that the positively polarized region of one and the negatively polarized region of the other attract each other.

Dipole-induced dipole interaction

(inductive effect) –Deformation of the electron cloud of a molecule in the presence of a dipole – polarisability!

The dipole/induced-dipole force combines features of both the induced-dipole/induced-dipole and induced-dipole/induced-dipole–dipole attractive forces. A polar region of one molecule alters the electron distribution in a nonpolar region of another in a direction that produces an attractive

force between them.

A force of attraction that results when a species with a permanent dipole induces a complementary dipole in a second species.

Induced dipole-induced dipole interaction

Quvantummechanical type, elemental dipoles are formed by electron movements In organic compounds it has important role e.g. – attraction between alkyl chains, lipophilicity

Induced-dipole/induced-dipole attractions are very weak forces individually, but a typical organic substance can participate in so many of them that they are collectively the most important of all the contributors to intermolecular attraction in the liquid state. They are the only forces of attraction possible between nonpolar molecules such as

alkanes.

Most alkanes have no measurable dipole moment, and therefore the only van der Waals force to be considered is the induced- dipole/induced dipole attractive force.

The neighboring molecule B “feels” the dipolar electric field of A and undergoes a spontaneous

adjustment in its electron positions, giving it a temporary dipole moment that is complementary to that of A.

The electric fields of both A and B fluctuate, but always in a way that results in a weak attraction between them.

Extended assemblies of induced-dipole/induced-dipole attractions can accumulate to give substantial intermolecular attractive forces. An alkane with a higher molecular weight has more atoms and electrons and, therefore, more opportunities for intermolecular attractions and a higher boiling point than one with a lower molecular weight.

The electric field of a molecule, however, is not static, but fluctuates rapidly. Although, on average, the centers of positive and negative charge of an alkane nearly coincide, at any

instant they may not, and molecule A can be considered to have a temporary dipole moment. It might seem that two nearby molecules A and B of the same nonpolar substance would be unaffected by each other.

A dipole–dipole attraction between the positively polarized proton of the OH group of one ethanol molecule and the negatively polarized oxygen of another. The term hydrogen

bonding is used to describe dipole–dipole attractive forces of this type. The proton involved must be bonded to an electronegative element, usually oxygen or nitrogen.

Protons in C―H bonds do not participate in hydrogen bonding.

Hydrogen bonding (in ethanol)

Hydrogen bonding in ethanol involves the oxygen of one molecule and the proton of an

―OH group of another. Hydrogen bonding is much stronger than most other types of dipole–dipole attractive forces.