Chemical Bonds 2.ppt

Chemical Bonding
and
Molecular Structure

Matter ??
made up of one or different type of elements
Does element exist independently??
Under normal conditions no other element
exists as an independent atom in nature, except
noble gases.
Group of atoms is found to exist together as
one species having characteristic properties.
Such a group of atoms is called a molecule.

How do atoms exist together??
Obviously there must be some force which
holds these constituent atoms together in the
molecules.
The attractive force which holds various
constituents (atoms, ions, etc.) together in
different chemical species is called a
Chemical Bond.

If so, now it raises following questions…..
1. Why do atoms combine? ?
2. Why are only certain combinations possible? ?
3. Why do some atoms combine while certain
others do not? ?
4. Why do molecules possess definite shapes? ?

To answer such questions different theories and concepts
have been put forward from time to time. These are
1.Kössel-Lewis approach
2.Valence Shell Electron Pair Repulsion (VSEPR)
Theory
3.Valence Bond (VB) Theory
4.Molecular Orbital (MO) Theory

Why I need to study about
Bonding????

Why I need to study about Bonding????
Every system tends to be more
stable and
bonding is nature’s way of lowering
the energy of the system to attain
stability

KÖSSEL-LEWIS APPROACH TO
CHEMICAL BONDING
To explain the formation of chemical
bond in terms of electrons, it was
only in 1916 when
Kössel and Lewis succeeded
independently in giving a satisfactory
explanation

Lewis pictured the atom in terms of
a positively charged ‘Kernel’ (the
nucleus plus the inner electrons)
and the outer shell that could
accommodate a maximum of
eight electrons.
This octet of electrons, represents
a particularly stable electronic
arrangement.
Lewis postulated that atoms
achieve the stable octet when they
are linked by chemical bonds.
1. Na+ and Cl– ions
2. Other molecules like Cl2 , H2 , F2 , etc.

Octet rule
atoms can combine either by
transfer of valence electrons from one atom to
another (gaining or losing)
or
by sharing of valence electrons in order to
have an octet in their valence shells

only the outer shell electrons take
part in chemical combination - Valence
Electrons
simple notations to represent valence
electrons in an atom – are called Lewis
symbols
The dots represent electrons. Such structures are referred to as
Lewis dot structures

formation of NaCl from
sodium and chlorine, according to
the said scheme

The bond formed, as a result of the
electrostatic attraction between the
positive and negative ions is termed
as the electrovalent bond

Langmuir (1919) refined the Lewis
Postulations by introducing the term covalent
bond.
The Lewis-Langmuir theory can be
understood by considering the formation of
the chlorine molecule (Cl2)
sharing of a pair of electrons
between the two chlorine
atoms,
each chlorine atom
contributing one electron to
the shared pair.

The important conditions being
that:
• Each bond is formed as a result of
sharing of an electron pair between the
atoms.
• Each combining atom contributes at
least one electron to the shared pair.
• The combining atoms attain the
outershell noble gas configurations as
a result of the sharing of electrons.

when two atoms share one
electron pair they are said to be
joined by a single covalent bond
If two atoms share two
pairs of electrons, the covalent
bond between them is called a
double bond
When combining atoms share
three electron pairs a triple
bond is formed

Problem
Write the Lewis dot structure
of CO molecule.
HINT
Step 1. Count the total number of valence
electrons of carbon and oxygen atoms.
Step 2. Write the skeletal structure of CO
Step 3. Draw bond between C and O
Step 4. Complete the octet on O

Formal Charge
Lewis dot structures, in general, do not represent the actual
shapes of the molecules.
Polyatomic ions→the net charge is possessed by the ion
as a whole and not by a particular atom.
The formal charge of an atom in a polyatomic molecule or
ion may be defined as
the difference between the number of valence
electrons of that atom in an isolated or free state
and the number of electrons assigned to that atom
in the Lewis structure.

Limitations of the Octet Rule
The octet rule, though useful, is not universal.
There are three types of exceptions to the octet rule.
1. The incomplete octet of the central atom
• In some compounds, the number of electrons
surrounding the central atom is less than eight.
• This is especially the case with elements having less
than four valence electrons.
• Examples are LiCl, BeH2 and BCl3.

2. Odd-electron molecules
• In molecules with an odd number of electrons
like nitric oxide, NO and nitrogen dioxide, NO2,
the octet rule is not satisfied for all the atoms

3. The expanded octet
• In a number of compounds there are more than
eight valence electrons around the central atom.
• This is termed as the expanded octet.
• Obviously the octet rule does not apply in such
cases.
• Some of the examples of such compounds
are: PF5, SF6, H2SO4 and a number of
coordination compounds.

Other drawbacks of the octet theory
• This theory does not account for the shape of
molecules.
• It does not explain the relative stability of the
molecules being totally silent about the
energy of a molecule.

IONIC OR ELECTROVALENT BOND
From the Kössel and Lewis treatment of the formation of
an ionic bond, the formation of ionic compounds would
primarily depend upon:
• The ease of formation of the positive and
negative ions from the respective neutral
atoms
• The arrangement of the positive and negative
ions in the solid, that is, the lattice of the
crystalline compound.

IONIC OR ELECTROVALENT BOND
• The formation of a positive ion involves
ionization, i.e., removal of electron(s) from the
neutral atom and that of the negative ion
involves the addition of electron(s) to the
neutral atom.

• The electron gain enthalpy, DegH, is
the enthalpy change, when a gas phase atom
in its ground state gains an electron.
• The electron gain process may be exothermic
or endothermic.
• The ionization, on the other hand, is always
endothermic.

• Electron affinity, is the negative of the
energy change accompanying electron
gain.

Obviously ionic bonds will be
formed more easily between
elements with comparatively
low ionization enthalpies and
elements with comparatively
high negative value of electron
gain enthalpy.
Most ionic compounds have
cations derived from metallic
elements and anions from non-
metallic elements.

In ionic solids,
the sum of the electron gain
enthalpy and the ionization
enthalpy
may be positive but still the crystal
structure gets stabilized
due to
the energy released in the
formation of the crystal lattice.

The energy released in the
processes is more than the energy
absorbed.
Stability of an ionic compound is
provided by its enthalpy of lattice
formation and not simply by achieving
octet of electrons around the ionic
species in gaseous state.

Lattice Enthalpy
The Lattice Enthalpy of an ionic solid is
defined as the energy required to
completely separate one mole of a solid
ionic compound into gaseous constituent
ions.
This process involves both the attractive
forces between ions of opposite charges and
the repulsive forces between ions of like
charge.

BOND PARAMETERS
Bond Length
Bond length is defined as the equilibrium
distance between the nuclei of two bonded
atoms in a molecule.
Each atom of the bonded pair contributes to
the bond length.
In the case of a covalent bond, the
contribution from each atom is
called the covalent radius of that atom.

The covalent radius is the radius of an atom’s
core which is in contact with the core of an
adjacent atom in a bonded situation.
The covalent radius is half of the distance
between two similar atoms joined by a
covalent bond in the same molecule.

The van der Waals radius represents
the overall size of the atom which
includes its valence shell in a
nonbonded situation.
Further, the van der Waals radius is half
of the distance between two similar
atoms in separate molecules in a
solid.

Bond Angle
It is defined as the angle between the orbitals
containing bonding electron pairs around the
central atom in a molecule/complex ion.
• It gives some idea regarding the
distribution of orbitals around the central
atom in a molecule/complex ion
• It helps us in determining its shape. For
example H–O–H bond angle in water can be
represented as under :

Bond Enthalpy
It is defined as the amount of energy
required
to break one mole of bonds of a
particular type between two atoms in a
gaseous state.
The unit of bond enthalpy is kJ mol–1.
It is important that larger the bond
dissociation enthalpy, stronger will be
the bond in the molecule.

Bond Order
In the Lewis description of covalent
bond,
the Bond Order is given by the number
of bonds between the two atoms in a
molecule.
A general correlation useful for
understanding the stabilities of molecules
is that: with increase in bond order, bond
enthalpy increases and bond length
decreases.

Polarity of Bonds
The existence of a hundred percent ionic or
covalent bond represents an ideal situation.
In reality no bond or a compound is either
completely covalent or ionic.
Even in case of
covalent bond between two hydrogen atoms,
there is some ionic character.

When covalent bond is formed between
two similar atoms, for example in H2, O2, Cl2,
N2 or F2, the shared pair of electrons is
equally attracted by the two atoms.
As a result electron pair is situated exactly
between the two identical nuclei.
The bond so formed is called
nonpolar covalent bond.

Contrary to this in
case of a heteronuclear molecule like HF, the
shared electron pair between the two atoms
gets displaced more towards fluorine since the
electronegativity of fluorine is far greater than that
of hydrogen.
The resultant
covalent bond is a polar covalent bond.

As a result of polarisation, the molecule
possesses the dipole moment which can be
defined as the
Product of the magnitude of the charge and the
distance between the centres of positive and
negative charge.
It is usually designated by a
Greek letter ‘μ’. Mathematically, it is expressed
as follows :

THE VALENCE SHELL ELECTRON
PAIR REPULSION (VSEPR) THEORY
Lewis concept is unable
to explain the shapes of molecules.
This theory provides a simple procedure to
predict the
shapes of covalent molecules.
Based on the repulsive interactions of the
electron pairs in the valence shell of the atoms.

The main postulates of VSEPR
theory are as follows:
• The shape of a molecule depends upon
the number of valence shell electron pairs
(bonded or nonbonded) around the central
atom.
• Pairs of electrons in the valence shell repel
one another since their electron clouds are
negatively charged.

• These pairs of electrons tend to occupy
such positions in space that minimise
repulsion and thus maximise distance
between them.
• The valence shell is taken as a sphere with
the electron pairs localising on the
spherical surface at maximum distance
from one another.

• A multiple bond is treated as if it is a single
electron pair and the two or three electron
pairs of a multiple bond are treated as a
single super pair.
• Where two or more resonance structures
can represent a molecule, the VSEPR
model is applicable to any such structure.

Valence Bond Theory
• Knowledge of atomic orbitals
• Electronic configurations of elements
• Overlap criteria of atomic orbitals
• Principles of variation and superpositions

Consider
the formation of hydrogen
molecule
Consider two hydrogen atoms A and B
approaching each other having
nuclei NA and NB and
electrons present in them are
represented by eA and eB.

When the two atoms
are at large distance from each other, there is
no interaction between them.
As these two
atoms approach each other, new attractive
and
repulsive forces begin to operate.
Attractive forces arise between:
(i) nucleus of one atom and its own electron
that is NA – eA and NB– eB.

(ii) nucleus of one atom and electron of other
atom i.e., NA– eB, NB– eA.
Similarly repulsive forces arise between
(i) electrons of two atoms like eA – eB,
(ii) nuclei of two atoms NA – NB.
Attractive forces tend to bring the two
atoms close to each other whereas repulsive
forces tend to push them apart

Orbital Overlap Concept
• In the formation of hydrogen molecule, there is
a minimum energy state when two hydrogen
atoms are so near that their atomic orbitals
undergo partial interpenetration.
• This partial merging of atomic orbitals is called
overlapping of atomic orbitals which results in
the pairing of electrons.

• The extent of overlap decides the
strength of a covalent bond.
• Greater the overlap the stronger is the
bond formed between two atoms.

Types of Overlapping and Nature
of Covalent Bonds
The covalent bond may be classified into
two types depending upon the types of
overlapping:
(i) Sigma(σ) bond, and (ii) pi(π) bond

(i) Sigma(σ) bond :
This type of covalent bond is formed by the
end to end (head-on)
overlap of bonding orbitals along the
internuclear axis.
This is called as head on overlap or axial
overlap.

s-s overlapping :
In this case, there is overlap of
two half filled s-orbitals along
the internuclear axis

s-p overlapping:
This type of overlap occurs between
half filled s-orbitals of one
atom and half filled p-orbitals of
another atom.

p–p overlapping :
This type of overlap takes place
between half filled p-orbitals
of the two approaching atoms.

(ii) pi(π ) bond :
In the formation of π bond
the atomic orbitals overlap in such a
way that their axes remain parallel
to each other and perpendicular to
the internuclear axis

To explain the characteristic
geometrical shapes of polyatomic molecules
like CH4, NH3 and H2O etc.,
the concept of hybridisation is considered.
Considerations are:
the atomic orbitals combine to form new set
of
equivalent orbitals known as hybrid orbitals.
The phenomenon is
known as hybridisation

Hybridization can be defined
as the process of intermixing of the orbitals
of slightly different energies so as to
redistribute
their energies, resulting in the formation of
new set of orbitals of equivalent energies
and shape.

Salient features of
hybridisation:
1. The number of hybrid orbitals is equal
to the number of the atomic orbitals
that get hybridised.
2. The hybridised orbitals are always
equivalent in energy and shape.

3. The hybrid orbitals are more effective
in forming stable bonds than the pure
atomic orbitals.
4. The hybrid orbitals are directed in
space in some preferred direction to
have
minimum repulsion between electron
pairs and thus a stable arrangement.

Important conditions for
hybridisation
1.The orbitals present in the valence shell
of the atom are hybridised.
2.The orbitals undergoing hybridisation
should have almost equal energy.
3.It is not necessary that only half filled
orbitals participate in hybridisation.

Various types of hybridisation
involving s, p and d orbitals.
(I) sp hybridisation:
This type of
hybridisation involves the mixing of one s
and one p orbital resulting in the
formation of two equivalent sp hybrid
orbitals.

(I) sp2 hybridisation:
This type of
and two p orbital resulting in the
formation of three equivalent sp2 hybrid
orbitals.

(I) sp3 hybridisation:
This type of
and three p orbital resulting in the
formation of four equivalent sp3 hybrid
orbitals.

Molecular Orbital Theory
The salient features of this theory are :
1. The electrons in a molecule are present
in the various molecular orbitals as the
electrons of atoms are present in the
various atomic orbitals.
2. The atomic orbitals of comparable
energies and proper symmetry combine to
form molecular orbitals.

3. An electron in an atomic orbital is
influenced by one nucleus, in a
molecular orbital it is influenced by
two or more nuclei depending
upon the number of atoms in the
molecule.
4. Thus, an atomic orbital is
monocentric while a molecular
orbital is polycentric.

5. The number of molecular orbital
formed is equal to the number of
combining atomic orbitals.
6. When two atomic orbitals combine,
two molecular orbitals are formed.
One is known as bonding
molecular orbital while the other is
called antibonding molecular
orbital.

7. The bonding molecular orbital
has lower energy and hence
greater stability than the
corresponding antibonding
molecular orbital.

Hydrogen Bonding
Nitrogen, oxygen and fluorine are the highly
electronegative elements.
When they are attached to a hydrogen atom
to form covalent bond, the electrons of the
covalent bond are
shifted towards the more electronegative
atom.

This partially positively charged hydrogen
atom forms a bond with the other more
electronegative atom.
This bond is known as
hydrogen bond and is weaker than the
covalent bond.
For example, in HF molecule,
the hydrogen bond exists between hydrogen
atom of one molecule and fluorine atom of
another molecule

Types of H-Bonds
There are two types of H-bonds
1. Intermolecular hydrogen
bond
2. Intramolecular hydrogen
bond

1) Intermolecular hydrogen
bond :
It is formed between two different
molecules of the same or different
compounds.
For example, H-bond
in case of HF molecule, alcohol or
water molecules, etc.

2) Intramolecular hydrogen
bond :
It is formed when hydrogen atom is in
between the two highly electronegative
(F, O, N) atoms present within the
same molecule.
For example, in o-nitrophenol
the hydrogen is in between
the two oxygen atoms.

The elements with a half-filled or fully-filled outer
s orbital
comprise the s-block elements.
Similarly, the elements with a partly filled or fully
filled outer p orbitals comprise the p-block
elements.
The elements between these two blocks that
is, have at most two electrons in the
outermost s orbital, and incompletely filled d
orbitals next to outermost orbital.

These elements in which successive addition
of electrons takes place progressively in the
inner d orbitals are called as d- block
elements.
Similarly, the elements in which filling up of
electrons takes place in inner to inner
f orbitals are known as f-block elements.
These two blocks of elements are generally
called transition elements.

transition metals
form a large number of complex
compounds in which
the metal atoms are bound to a
number of anions or
neutral molecules.
In modern terminology such
compounds are called coordination
compounds.

Coordination Compounds
The transition metals and their ions have
much higher tendency to from coordination
compounds as compared to the s- and p- block
elements.
It is because of their relatively
smaller sizes, higher ionic charges and the
availability of d orbitals for bond formation.

insights into the functioning of
vital components of biological
systems-Chlorophyll,
haemoglobin and vitamin B12
are coordination
compounds of magnesium,
iron and cobalt respectively.

Coordination entity
A coordination entity constitutes a central
metal atom or ion bonded
to a fixed number of ions or molecules.
For example, [CoCl3(NH3)3]
is a coordination entity in which the cobalt ion
is surrounded by
three ammonia molecules and three chloride
ions.
Other examples are [Ni(CO)4], [PtCl2(NH3)2], [Fe(CN)6]4–,
[Co(NH3)6]3+ .

Central atom/ion
In a coordination entity, the atom/ion to which
a fixed number
of ions/groups are bound in a definite
geometrical arrangement
around it, is called the central atom or ion.
These central atoms/ions are also referred to as
Lewis acids.

Ligands
The ions or molecules bound to the
central atom/ion in the
coordination entity are called ligands.
These may be simple ions
such as Cl–, small molecules such as
H2O or NH3, or even macromolecules,
such as proteins.

When a ligand is bound to a metal ion
through a single donor
atom, as with Cl–, H2O or NH3, the ligand
is said to be unidentate.
When a ligand can bind through two
donor atoms as in
H2NCH2CH2NH2 (ethane-1,2-diamine) or
C2O4
2– (oxalate), the ligand is said to be
didentate

when several donor atoms are
present in a single ligand as in
N(CH2CH2NH2)3, the ligand is
said
to be polydentate.

Coordination number
The coordination number (CN) of a metal ion
in a complex can be
defined as the number of ligand donor
atoms to which the metal is
directly bonded.
For example, in the complex ions, [PtCl6]2–
and [Ni(NH3)4]2+, the coordination number of
Pt and Ni are 6 and 4 respectively.

The main postulates are:
1. In coordination compounds metals show
two types of linkages
(valences)-primary and secondary.
2. The primary valences are normally
ionisable and are satisfied by
negative ions.

3. The secondary valences are non
ionisable. These are satisfied by
neutral molecules or negative ions. The
secondary valence is equal to
the coordination number and is fixed for a
metal.
4. The ions/groups bound by the secondary
linkages to the metal have
characteristic spatial arrangements
corresponding to different
coordination numbers.

Bonding in Coordination
Compounds

basic questions like:
(i) Why only certain elements possess the
remarkable property of forming coordination
compounds?
(ii) Why the bonds in coordination compounds
have directional properties?
(iii) Why coordination compounds have
characteristic magnetic and
optical properties?

Describe the electronic and magnetic
structure of crystalline solids.
Crystal Field Theory describes the
interaction between a central metal ion
that is surrounded by anions.
A quantum mechanical description of
the metal ion is employed, with
attention focused on the valence shell
d, s, and p orbitals.

The surrounding anions are typically
treated as point charges.
The essential insight of Crystal Field
Theory is that the geometry of the
negatively charged point charges
influences the energy levels of the
central metal ion.

effects on a metal d orbital energies of moving a
set of negative point charges close to a metal
ion.
The energies of the d orbitals rise as the
negative charges approach the metal ion,
owing to the repulsions between the d orbital
electrons and the surrounding charge.
If the surrounding negative charge is spherically
symmetric, all five d orbitals are equally affected.

In practice, the surrounding negative
charge is never spherically
distributed, because the charge is
associated with specific ions that
occupy specific positions.
The consequence is each d orbital is
affected differently, and how a
particular d orbital is affected
depends upon the geometry of the
surrounding point charges.

CFT provides valuable insights into the electronic
structure of transition metals in crystal lattices, it
has serious limitations.
The most severe limitation is its inability to
account for chemical bonding.
Complexes may also be formed between neutral
metal atoms and neutral or cationic ligands.
Crystal Field Theory is poorly suited to explain
such interactions.

A more detailed description of bonding in
coordination compounds is provided
by Ligand Field Theory.
In coordination chemistry, the ligand is a
Lewis base, which means that the ligand is
able to donate a pair of electrons to form a
covalent bond.

The metal is a Lewis acid, which means it has
an empty orbital that can accept a pair of
electrons from a Lewis base to form a covalent
bond.
This bond is sometimes called a coordinate
covalent bond or adative covalent
bond to indicate that both electrons in the bond
come from the ligand.

The following list summarizes the
key concepts of Ligand Field Theory
• One or more orbitals on the ligand overlap with
one or more atomic orbitals on the metal.
• If the metal- and ligand-based orbitals have
similar energies and compatible symmetries, a
net interaction exists.
• The net interaction produces a new set of
orbitals, one bonding and the other antibonding
in nature.

• Where no net interaction exists, the original
atomic and molecular orbitals are
unaffected and are nonbonding in nature as
regards the metal-ligand interaction.
• Bonding and antibonding orbitals are of
sigma (σ) or pi (π) character, depending
upon whether the bonding or
antibonding interaction lies along the line
connecting the metal and the ligand.

Biological role of transition metals
Metal ions play essential roles in about
one third of enzymes
These ions can
modify electron flow in a substrate or
enzyme, thus effectively controlling an
enzyme catalyzed
reaction.

Metal ions may be bound by main-chain amino
and carbonyl groups,
but specific binding is achieved by the amino
acid side chains, particularly the
carboxylate groups of aspartic and glutamic
acid, and the ring nitrogen atom of
histidine.
Other side chains that bind metals ions include
tryptophan (ring nitrogen),
cysteine (thiol), methionine (thioether), serine,
threonine, tyrosine (hydroxyl groups),
and asparagine and glutamine (carbonyl groups, less
often amino groups).

Copper-based protein assays, including the BCA and Lowry methods, depend on the
well-known "biuret reaction", whereby peptides containing three or more amino acid
residues form a colored chelate complex with cupric ions (Cu2+) in an alkaline
environment containing sodium potassium tartrate.
It is chemically similar a complex that forms with the organic compound biuret
(NH2-CO-NH-CO-NH2) and the cupric ion. Biuret, a product of excess urea
and heat, reacts with copper to form a light blue
tetradentate complex.

Extraction of
sample from the
Biological Matrix

Bioanalysis is a sub-discipline of
analytical chemistry covering the
quantitative measurement drugs and
their metabolites in biological
systems.
Bioanalytical assays to accurately and
reliably determine the drugs at lower
concentrations

Some techniques commonly used in
bioanalytical studies include:
Hyphenated techniques
• LC–MS (liquid chromatography–mass spectrometry)
• GC–MS (gas chromatography–mass spectrometry)
• LC–DAD (liquid chromatography–diode array detection)
• CE–MS (capillary electrophoresis–mass spectrometry)
Chromatographic methods
• HPLC (high performance liquid chromatography)
• GC (gas chromatography)
• UPLC (ultra performance liquid chromatography)
• Supercritical fluid chromatography

Because biological samples are extremely
complex matrices comprised of many
components
that can
interfere with good separations and or good
mass spectrometer signals, sample
preparation is an important aspect of
bioanalytical estimation.
This is important whether samples originate
as tissue extracts, plasma, serum or urine.

Sample preparation is a difficult step, and
generally involves
Filtration
solid phase extraction
precipitation and desalting.
Sample preparation prior to chromatographic
separation is performed to dissolve or dilute the
analyte in a suitable solvent, removing the
interfering compounds and pre-concentrating the
analyte.

The principle objectives of
sample preparation from
biological matrix are
a. Isolation of the analytes of interest from
the interfering compounds
b. Dissolution of the analytes in a suitable
solvent and pre-concentration.

Goal and objectives of sample
preparation
Quantitative recovery - A minimum number of
steps.
Successful sample preparation has a threefold
objective
• In solution
• Free from interfering matrix elements
• At a concentration appropriate for
detection and measurement

Extraction is the withdrawing of an active
agent or a waste substance from a solid or
liquid mixture with a liquid solvent.
The solvent is not or only partially miscible with
the solid or the liquid.
By intensive contact between analyte and the
extraction medium this leads the analyte
transfers from the solid or liquid mixture
into the extraction medium (extract).

The most common approach in analyte
separation involves a two phase
system where the analyte and
interferences are distributed between
the two phases.
Distribution is an equilibrium process and
is reversible.

If the sample is distributed between
two immiscible liquid phase, the
techniques is called liquid-liquid
extraction.
If the sample is distributed between a
solid and a liquid phase, the
technique is called liquid-solid
adsorption

Molecular phenomena for
solubility and miscibility
To dissolve a drug, a solvent must break
the bonds like ionic bond, hydrogen
bond and Van der Waals forces which
inter links the compound to its neighbours
and must not break substantial
intermolecular bonds of the solvent
without replacing them with drug solvent
interaction.

Extractions
1.) Definition
 The transfer of a compound from one chemical phase to another
- The two phases used can be liquid-liquid, liquid-solid, gas-solid, etc
- Liquid-liquid is the most common type of extraction
- The partitioning of solute s between two chemical phases (1 and
2) is described by the equilibrium constant K
1
2
]
S
[
]
S
[
K 
K is called the partition coefficient
Immiscible
liquids

2.) Extraction Efficiency
 The fraction of moles of S remaining in phase 1 after one extraction
can be determined
- The value of K and the volumes of phases 1 and 2 need to be known
 The fraction of S remaining in phase 1 after n extractions is
 
2
1
1
KV
V
V
q


where: q = fraction of moles of S remaining in phase 1
V1 = volume of phase 1
V2 = volume of phase 2
K = partition coefficient
 
n
n
KV
V
V
q 







2
1
1 Assumes V2 is constant

Example #1:
 Solute A has a K = 3 for an extraction
between water (phase 1) and benzene
(phase 2).
If 100 mL of a 0.01M solution of A in
water is extracted one time with 500
mL benzene, what fraction will be
extracted?

Solution:
First determine fraction not extracted (fraction still in phase 1, q):
 
%
.
.
)
mL
(
)
(
mL
mL
KV
V
V
q
n
n 2
6
062
0
500
3
100
100
1
2
1
1 


















The fraction of S extracted (p) is simply:
%
.
.
.
q
p 8
93
938
0
062
0
1
1 






Example #2:
 For the same example, what fraction will be
extracted if 5 extractions with 100 mL
benzene each are used (instead of one 500
mL extraction)?

Solution:
Determine fraction not extracted (fraction still in phase 1, q):
 
%
.
.
)
mL
(
)
(
mL
mL
KV
V
V
q
n
n 98
0
00098
0
100
3
100
100
5
2
1
1 


















The fraction of S extracted (p) is:
%
.
.
.
q
p 902
99
99902
0
00098
0
1
1 






Selection of the solvent
• Practical concerns when choosing
extraction solvents.
• the two solvents must be immiscible.
• The properties of an ideal solvent is that it
should withdraw the active agent from
a mixture

Selectivity: Only the active agent has to
be extracted and no further substances which
mean that a high selectivity is required.
Capacity: To reduce the amount of
necessary solvent, the capacity of the solvent
has to
be high.
Miscibility: To achieve simple regeneration
of the solvent the miscibility of solvent has
to be low.

Difference in density: After extraction, the two
phases have to be separated in a separator and for
this a high positive difference in density is required.
Optimal surface tension: σ low → low amount of
energy for dispersing required; if
surface tension < 1 mN/m stable emulsions are
produced; σ > 50 mN/m → high
amount of energy for dispersing and high tendency
to coalesce.
Recovery: The solvent has to be separated from
the extract phase easily to produce solvent free
active agents.

• Corrosion: If the solvent is corrosive prices for
construction increases.
• Low price
• No or low toxicity and not highly flammable
• Flame temperature: 25°C higher than operating
temperature
• Vapour pressure: To prevent loss of solvent by
evaporation a low vapour pressure at operating
temperature is required.
• Viscosity: A low viscosity of the solvent leads to
low pressure drop and good heat and mass
transfer.

• Chemical and thermal stability
• Environmentally acceptable or easily
recoverable
• Convenient specific gravity
• Suitable volatility
• High chemical stability and inertness
• Not prone to form an emulsion
• Dissolves the neutral but not the
ionized form of the analyte

One of the most useful techniques for
isolating desired components from a mixture
is liquid-liquid extraction (LLE).
LLE is a method used for the separation of a
mixture using two
immiscible solvents.
In most LLEs, one of the phases is aqueous
and the other is an
immiscible organic solvent.

The ability
to separate compounds in a mixture using the
technique of LLE
depends upon how
differently the compounds of the sample
mixture partition themselves between the two
immiscible solvents.
Selective partitioning of the compound of interest
into one of two
immiscible or partially miscible phases occurs
by the proper choice of extraction of solvent.

In general liquid-liquid extractions can separate four
different classes of compounds:
a. Organic bases: Any organic amine can be
extracted from an organic solvent with a
strong acid such as 1M hydrochloric acid
b. Strong acids: Carboxylic acids can be extracted
from an organic solvent with a weak
base such as 1M sodium bicarbonate
c. Weak acids: Phenols can be extracted from an
organic solvent with a strong base such
as 1M sodium hydroxide
d. Non-polar compounds stay in the organic layer

Disadvantages
• Large solvent consumption is needed
for extraction of drug.
• LLE is time consuming process when
compare to other methods.
• LLE require an evaporation step prior
to analysis to remove excess of
organic solvent.

Disadvantages
• LLE technique is not a suitable one for
the estimation of several analytes.
• Emulsion formation may be possible
when two immiscible phases were
used in the extraction procedure.

Mixed solvents
In some cases pure solvents will not be satisfactory
for the extraction of the compound of interest.
Alcohols are excellent solvent but
those with lower boiling points are too soluble in
water whereas
less miscible one are having high boiling points,
but the use of mixed
solvents containing alcohols can solve the problem.
1:1 mixture of tetrahydrofuran and
dichloromethane is a powerful solvent for the extraction of polar
compounds from aqueous solutions.

Several useful equations can help illustrate the
extraction process.
The Nernst distribution
law states that any neutral species will
distribute between two immiscible solvents so
that
the ratio of the concentration remains constant.
KD = Co/Caq

KD = Co/Caq
Where KD is the distribution constant,
Co is the concentration of the analyte in
the organic
phase, and
Caq is the concentration of the analyte in
the aqueous phase.

To increase the value of KD,
several approaches may be used:
• The organic solvent can be changed to
increase solubility of the analyte
• If the analyte is ionic or ionizable, its KD may
be increased by suppressing its ionization
to make it more soluble in the organic phase.

Liquid-liquid extraction
- Accomplished with a separatory funnel
- Shaken to increase surface area between phases
- When stop, denser phase settling to the bottom
Solid Phase Extraction
- Solid-phase refer to solid adsorbent in the cartridge
- many choice of adsorbent determined by the
properties of the species being retain in matrix in
which it is found.

•Replace liquid-liquid extraction due to
ease of use
faster extraction time
decreased volume of solvent and
able to concentrate the analytes.
•Solid-phase microextraction developed for
less sample.

A separatory funnel
Solid-phase extraction cartridge

Continuous Extraction
•For component of interest that
has unfavorable partition
coefficient.
•Extraction is accomplished by
continuously passing the
extracting phase through the
sample until a quantitative
extraction is achieved
•Involving solid samples are
carried out with a Soxhlet
extractor

Release of protein from biological
host
•To gain access to the product
•Access to the product is simple and
inexpensive when the
protein is produced extracellularly
•Microbial sources are preferred

• Mammlian cell hosts are preferred when
posttranslational modification is essential for
the function of eukaryotic
proteins
• Bulk enzymes are invariably produced
extracellularly by Bacillus species & fungi, as
are the proteins produced
by mammalian cell culture

Cell envelops of bacteria and yeast
Gram-positive Gram-negative Yeasts
bacteria bacteria
Cytoplasmic
membrane
Peptidoglycon
layer
Lipopolysaccharide
membrane
Mannan partially cross-
linked by phosphodiester
bridges
Glucan layer
with proteins
Animal cells: no cell wall, thus fragile in breaking
Plant cells: composed of cellulose and other polysaccharides

Cell Disintegration Techniques
Technique Example Principle
Gentle
Cell lysis Erythrocytes Osmotic disruption of cell
membrane
Enzyme digestion Lysozyme treatment
of bacteria
Cell wall digested, leading
to osmotic disruption
Chemical solubilization Toluene extraction of
yeast
Cell wall (membrane)
partially solubilized
chemically
lytic enzymes released
Hand homogenizer Liver tissue Cells forced through narrow
gap, disrupts cell membrane
Minicing (grinding) Muscle etc. Cells disrupted during
minicing process by shear
force

Cell Disintegration Techniques
Technique Example Principle
Moderate
Blade homogenizer Muscle tissue, most Chopping action breaks up
(waring type) animal tissues, plant large cells, shears apart
tissues smaller ones
Grinding with abrasive Plant tissues, bacteria Microroughness rips off
(sand, alumina) cell walls
Vigorous
French press Bacteria, plant cells Cells forced through small
orfice at very high press-
ure; shear forces disrupt
cells
Ultrasonication Cell suspensions Micro-scale high-pressure
sound waves cause dis-
ruption by shear forces
and cavitation

Mechanical Methods
Limitations
• High risk of damage to the product
• Heat denaturation a major problem
• The release of proteases from cellular
compartments can lead to enzymatic degradation
of the product
• Products released encounter an oxidizing
environment, that can cause denaturation and
aggregation

Non-Mechanical Methods
Physical Rupture of Microbial Cells
Desiccation: by slow drying in air, drum drying, etc
followed by extraction
Osmotic shock: Changes in the osmotic pressure of the
medium may result in the release of certain enzymes,
particularly periplasmic proteins in gram negative cells.
Suspending a cell suspension in a solution with high salt
concentration

High temperature: Exposure to high
temperature can be an effective approach to
cell disruption but is limited to heat-stable
products.
Heating to 50 – 550C disrupts outer
membrane, releases periplasmic proteins.
Heating at 900C for 10 min can be used for
releasing cytoplasmic proteins

Physical Rupture of Microbial Cells
Freeze-thawing: Rupture with ice crystals is
commonly used method. By slowly freezing and
then thawing a cell paste, the cell wall and
membrane may be broken, releasing enzymes
into the media

Nebulization: In nebulization gas is blown over a
surface of liquid through a neck. Because of the
differential flow within the neck, the cells are
sheared
Decompression: When pressurized, the microbial
cells are gradually penetrated and filled with gas.
After saturation by the gas, the applied pressure is
suddenly released when the absorbed gas rapidly
expands within the cells leading to rupture
Note: Methods produce low protein yields and require
long process time

• Dialysis is based on diffusion during which the
mobility of solute particles between two liquid
spaces is restricted, mostly according to their
size.
• Size restriction is achieved by using a porous
material, usually a semi-permeable membrane
called dialysis membrane.
• This membrane is permeable only for particles
below a certain size.

A dialysis membrane is a semi-permeable film
(usually a sheet of regenerated cellulose)

Type 1: Cross Filtration
 Flow parallel to membrane
surface
 Does not cause buildup, therefore
does not suffer from reduced flow
overtime

Type 2: Dead End Flow
 Flow perpendicular to
membrane surface
 Causes build up of filter
cake on membrane

What properties of proteins
can be used to separate and
purify them from each other?

Solubility of proteins
• Multiple acid-base groups on proteins affect their
solubility properties.
• Solubility of a protein is therefore dependent on
concentrations of dissolved salts, the polarity of
solvent, the pH and the temperature.
• Certain proteins will precipitate from solutions under
conditions which others remain soluble-so we can
use this as an initial purification step of proteins.
• Salting out or salting in procedures take advantage
of ionic strength

•Ammonium sulfate is the most commonly
used reagent
•High solubility (3.9 M in water at 0 ºC)
•High ionic strength solution can be made (up to
23.5 in water at 0 ºC)
Note-certain ions (I-, ClO4
-, SCN-, Li+, Mg2
+, Ca2
+ and Ba+) increase the
solubility of proteins rather than salting out. (also denature proteins).

•Water-miscible organic solvents also
precipitate proteins.
•Acetone, ethanol
•This technique is done at low temperatures
(0 ºC) because at higher temperatures, the
solvent evaporates.
•Some water-miscible organic solvents
(DMF, DMSO) are good at solubilizing
proteins (high dielectric constants).

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