MEDICAL Biochemistry

MEDICAL BIOCHEMISTRY


http://osp.mans.edu.eg/medbiochem_mi/cources/biochemistry/


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Physical Chemistry

Physical Chemistry

Objectives:

In this chapter you will learn some basics of chemistry which are necessary for understanding
biochemistry.
Also, some important physico-chemical phenomena as diffusion, dialysis and osmosis
will be studied.

Lec. Title PDF Swf

1 Introduction, Chemical bonds
2 Law of mass action, Water, pH
3 Acids and alkalis, Buffers
4 Solution, Solution phenomena
5 Osmosis, Expression of Concentration

Chemical bonds >> Lecture 1

Biochemistry is the study of living systems using the methods of chemistry and physics.
Physical chemistry deals with the physico-chemical phenomena, which are needed to understand biochemistry.
Organisms are complicated and highly organized. Each organism consists of many organ systems
( e.g. respiratory and reproductive systems). Each system is formed of many organs which are formed of tissues. The tissues are formed of cells that contain cell organelles. The cell and its organelles are made up from molecules.

Cells are composed of different types of molecules:

Water 70%
Proteins 15%
Nucleic acids 7%
Polysaccharides 3%
Lipids 1%
Minerals and others 4%
Molecules are formed of elements :
Carbon (C), hydrogen (H), oxygen (O), nitrogen (N), phosphorous (P) and sulfur (S) are the main
elements that constitute more than 99% of the human body.
Elements are substances that cannot be broken down further by ordinary chemical means.
Atoms are the smallest unit of matter retaining the properties of an element.
Each atom is made up of positively charged protons, neutral neutrons, and negatively
charged electrons.

Electrons are generally affected by the following three chemical phenomena:

Electrons tend to pair
[/color]They are negatively charged and so are subject to the electrostatic
attraction-repulsion rules.
The relative ability of an atom to draw electrons in a bond toward itself is called the electronegativity
of the atom.
Atoms with high electronegativitiy attract the electrons more than those that have small electronegativity.
Atoms tend to fill their outermost electron energy level (orbit) either by transfer or sharing of electrons.
These three factors are the basis for the different types of chemical bonds and chemical reactions
that occur in nature.

Chemical Bonds

The components of the system become more stable through the formation of bonds.

There are several types of chemical bonds:

1- Covalent Bonds

Covalent Bonds are formed by sharing of a pair of electrons. Electrons are shared in pairs.
Two atoms sharing a single pair of electrons have a single bond, while two atoms sharing two pairs have
a double bond and two atoms sharing three pairs have a triple bond.
Covalent Bonds are the strongest chemical bonds, the energy of a single covalent bond can
vary from 50 kcal/mol to 110 kcal/mol depending on the elements involved. Once formed, covalent bonds
rarely break spontaneously at room temperature because of the high amount of energy required.

Carbon-carbon bonds (C-C), carbon-oxygen bonds (C-O) and carbon-hydrogen
bonds (C-H) are all examples
of covalent bonds. Methane is formed of one carbon atom covalently attached to
4 hydrogen atoms by 4 covalent bonds.

2- Ionic Bonds
Ionic bonds are formed when there is a complete transfer of electrons from

one atom to another to
fill their outermost energy levels. This electron transfer results in two ions,
one positively charged
and the other negatively charged. These ions become attracted to each other by
the resulting electrostatic
charge differences. Ionic bonds are generally weak. They are often 4-7 kcal/mol in strength.
So they can be broken easily when subjected to heat or submerged in water.
An example of this process is the formation of a sodium chloride molecule.

Ionic bonds are also called electrostatic bonds as they result from the electrostatic attraction
between two ionized atoms or groups of opposite charge.

3- Hydrogen Bonds

Hydrogen bonds result from electrostatic attraction between an electronegative atom
e.g. oxygen or nitrogen (O or N) and a hydrogen atom that shares its electron with a
second electronegative atom.
Hydrogen bonds occur between two or more polar molecules.
A polar molecule is a molecule that has a slight positive charge at one end
and a slight negative charge
on the other (giving it poles). The bond is quite weak (5 kcal/mol) and easily broken,
unlike covalent bonds.
Accumulation of many hydrogen bonds provides specificity and significant stability
to macromolecular structures.

Hydrogen bonds are frequently found in proteins and nucleic acids in large numbers
and serve to keep the protein or nucleic acid structure secure.

Perhaps the most famous example of hydrogen bonds is the bond which is formed between oxygen
of water molecule and the hydrogen of another water molecules.
Each water molecule can form up to 4 hydrogen bonds.

4- Van der Waals Interactions

Van der Waals interactions are intermolecular forces of attraction that occur when there is
a transient asymmetry in the distribution of charge around atoms in a molecule.
The consequent charge imbalance affects and attracts adjacent atoms.
Van der Waals interactions are
very weak bonds (1 kcal/mol) formed between non-polar molecules or non-polar parts
of a molecule that have slight transient charge displacements.

5- Hydrophobic Interactions

Hydrophobic interactions occur between clusters of nonpolar molecules that
tend to aggregate
so as to minimize the surface area that is exposed to water. Hydrophobic molecules
tend to aggregate together in avoidance of H2O molecules.

6- Steric Hindrance

Atoms occupy a fixed volume of space that is very difficult to compress, except by covalent
bond formation.
Thus, atoms cannot overlap in their position. The effect of this on protein structure
is called steric hindrance.
Bulky side-chains such ( as that found in isoleucine amino acid) restrict the possible side-chain
angles in protein structure.

Biological Functions where non-covalent interactions play roles :

1- Binding Specificity

Specificity of enzyme substrate binding is due to formation of enough
noncovalent bonds to
hold the enzyme and substrate together. Also, specificity of antigen antibody
reactions is
due to formation of enough noncovalent bonds.

2- Protein Structure

Secondary, tertiary and quaternary protein structures are stabilized by noncovalent bonds.
Collagen ,a protein whose function depends on its ability to maintain a long and fibrous structure, is stabilized by noncovalent bonds.

3- DNA Base Pairing

Hydrogen bonding between adjacent base pairs underlies the ability of one strand of DNA
to pair with another and serves to hold the two strands of the DNA double helix together.

Physical Chemistry
Lecture 2
Law of mass action,

Water, pH

Chemical Reactions

Chemical reactions are classified into:

Irreversible reaction, which proceeds in one direction e.g. reaction of acids with alkalies to form salts and water

Reversible reaction, which proceeds in both directions. Most of the reactions in biological systems are of this type.
In reversible reaction, the reactants (A and B) react together to give the products (C and D) and vise versa i.e. the reversible reaction does not proceed to completion in either direction.

Law of Mass Action

The law of mass action states that "In reversible reactions the velocity of the reaction at a certain temperature is directly proportional to the molecular concentration of the reacting substances".
If we have a reversible reaction in which A and B are the reacting substances, C and D are the products, V1 is the velocity of the reaction in the forward direction and V2 is the velocity of the reaction in the reverse direction.

According to the law of mass action, the velocity of the reaction in the forward direction (V1) is directly proportional to the molecular concentration of A and B.

Also, the velocity of the reaction in the reverse direction (V2) is directly proportional to the molecular concentration of C and D.

At equilibrium, the velocity of the reaction in the forward direction (V1) is equal to the velocity of the reaction in the reverse direction (V2).

Where Keq is the equilibrium constant.
So, in reversible reactions at equilibrium the product of molecular concentrations of the resulting substances divided by the product of molecular concentrations of the reacting substances is a constant called equilibrium constant (Keq).

Water

Water is the liquid in which all known life forms exist. By far the most common molecule in the living organisms is water.

Since most of the molecular interactions in the body take place in an aqueous environment, we must always consider the effects of water on biomolecules.

Water molecule is formed of one oxygen atom covalently linked to 2 hydrogen atoms. Since O has a much higher electron affinity (electronegativity) than H, it attracts the electrons away from the hydrogen atoms, resulting in a partial negative charge on the oxygen and a partial positive charge on each of the hydrogen atoms. This creates a positive pole near the hydrogen atoms and a negative one on the opposite side. So, water is best described as a polar compound. One end, or pole, of the molecule has a partial positive charge delta+ () on the hydrogen atoms, and the other end has a partial negative charge (2 ) on the oxygen atom.

Biomolecules Interacting with Water Fall into Three Categories:

Hydrophilic (water loving) molecules which readily form hydrogen bonds with water and therefore dissolve in water. These are substances that are ionic or can engage in hydrogen bonding with water molecules e.g. sodium chloride and glucose.
Hydrophobic (water hating)molecules that can not form hydrogen bond with water and can not form ionic bonds. These are nonionic and nonpolar substances that are unable to engage in attractive interactions with water molecules e.g. neutral fat and oils.
Amphipathic molecules which have both a hydrophobic and a hydrophilic portion. In aqueous solutions these compounds spontaneously aggregate with the hydrophilic portion exposed to the water and the hydrophobic portion in the interior where it is shielded from the water molecules e.g. lipoproteins and phospholipids.

Dissociation (Ionization) of water

can act as both a proton donor (acid) and proton acceptor (base) for itself. A proton can be transferred from one water molecule to another, resulting in the formation of one hydroxyl ion and one hydronium ion
.
This is called the autoionization or dissociation of water. This equilibrium can also be expressed as:

Although this is misleading because free protons are rarely present in pure water.
At equilibrium :

Where:

Keq is the equilibrium constant. It is a constant that equals 1.8 x 10-16
is hydrogen ion concentration in mole/liter (mol/L)
is the hydroxyl ion concentration in mol/L.
is the concentration of water in mol/L. Each liter of pure water weighs 1000 gram and the molecular weight of water is 18. So, in pure water = 1000/18 =55.5 M

By substitution:

Rearranging gives:

Keq X 55.5 is another constant called ion product of water (Kw)

Pure water which is neutral has the same concentration of as ions

As the ion-product constant of water (Kw) always remains constant at equilibrium, consequently:

If the concentration of either or increases in solution, then the other must decrease to compensate.
An acid causes an increase in the concentration of ions and an alkali causes an increase in the concentration of ions.
So, in acidic solutions is more than , while in alkaline solutions is more than .
Hydrogen ion concentration is used to express acidity or alkalinity in aqueous solution as follows:

In neutral solution = =
In acidic solution > and > while <
In alkaline solution > and < while >
pH

pH is a measure of the concentration of hydrogen ions. It is defined as the negative logarithm of hydrogen ions concentration in mol/L.

pH = - log

As mentioned before, the in pure water at 25 C is ; therefore the pH of pure water is:

pH = -log()
pH = - (-7)
pH = 7

pH ranges from 0 to 14:
pH = 0, the most acidic (1 molar HCl)
pH < 7, solution is acidic
pH = 7, solution is neutral
pH > 7, solution is alkaline
pH = 14, the most alkaline (1 molar NaOH)

A lower pH always means a higher concentration of [H+] . The most acidic fluid in the body is gastric juice, which has a pH range 1-2. While the most alkaline juice in the body is the pancreatic juice with pH about 8. Milk (pH = 6.6 - 6.9) and urine (pH=6) are slightly acidic while blood, saliva and aqueous humor of the eye are slightly alkaline (pH = 7.4).

Blood pH is kept within very narrow range (7.37 - 7.43) by blood buffers.
Also, the negative logarithmic scale is useful for measuring hydroxyl ion concentration :

pOH = - log

As discussed in dissociation of water:

By logarithmic transformation, we obtain the following useful expressions:
pH + pOH = - log = 14

This means that the sum of pH and pOH is always 14 and if either of them increases then the other must decrease to compensate.

The Relationship Between pH and pOH
pH pOH mol/L mol/L
0 14 1.0 10-14
2 12 0.01 10-12
4 10 0.0001 10-10
6 8 10-6 10-8
8 6 10-8 10-6
10 4 10-10 0.0001
12 2 10-12 0.01
14 0 10-14 1.0
N.B.: The pH scale is logarithmic scale representing the concentration of ions in a solution. The pH scale is logarithmic, so pH difference of 1 means a tenfold difference in the relative concentration of hydrogen ions.

Measurement of pH :

Colorimetric method
It depends on indicators, which are weak organic acids that change their colours on ionization i.e. they have certain colour in the non-ionized state and another colour in the ionized state. For example the use of phenolphthalein solution or litmus paper to measure pH.
Electrometric method
pH is measured by pH meter, which uses hydrogen sensitive electrode that measures hydrogen ion concentration in the solution.
Buffers

They are solutions that resist appreciable changes in pH when an acid or alkali is added to it.
Biological systems use buffers to control pH.

Buffers may be formed of one of the following mixtures:

Weak acid and its salt with strong base e.g. carbonic acid and sodium bicarbonate.
Weak base and its salt with strong acid e.g. ammonium hydroxide and ammonium chloride.

Mechanism of Buffer Action

In the case of bicarbonate buffer (NaHCO3 / H2CO3)
If strong alkali as sodium hydroxide (NaOH) is added it will react with the carbonic acid forming sodium bicarbonate, which is a weak alkali. So there is no appreciable change in the pH.

If strong acid as hydrochloric acid (HCl) is added it will react with the sodium bicarbonate forming carbonic acid, which is a weak acid. So there is no appreciable change in the pH.

Buffer Systems of The Blood :

Plasma buffers
They include:
a) Bicarbonate buffer (NaHCO3 / H2CO3)
b) Phosphate buffer (Na2HPO4 / NaH2PO4)
c) Plasma proteins (Na proteinate / H protein)
RBCs buffers
They include:
a) Bicarbonate buffer (KHCO3 / H2CO3)
b) Haemoglobin buffer (KHb /HHb)
c) Oxyhaemoglobin buffer (KHbO2 /HHbO2)

Clinical significance of Blood Buffers :

The pH of blood is maintained in a narrow range around 7.4 (7.37 - 7.43) by the buffer systems in blood plasma and red blood cells.

A decrease in blood pH, which can be compensated, is called acidosis while uncompensated decrease in blood pH is called acidaemia.

An increase in blood pH, which can be compensated, is called alkalosis while uncompensated increase in blood pH is called alkalaemia .

Even relatively small changes in this value of blood pH can lead to severe metabolic consequences. Therefore, blood buffers are extremely important in order to maintain homeostasis.

Physical Chemistry >> Acids and alkalis, Buffers >> Lecture 3
Acids and Alkalis

Acids :

They are substances that give hydrogen ions in solution i.e. they are potential proton [H+] donors.

The strength of an acid depends on the degree of dissociation (ionization) in solution. Hydrochloric acid (HCl) is almost completely dissociated so, it is a strong acid while acetic acid (CH3COOH) is a weak acid as it slightly dissociates.

Types of acidity :

There are 2 types of acidity; true and titratable acidity

True Acidity
It is the concentration of hydrogen ions in solution..
It depends on the degree of ionization.
It is measured by pH meter.
Titratable (Total) Acidity
It is the total concentration of hydrogen available for ionization whether ionized or not.
It does not depend on the degree of ionization.
It is measured by titration against an alkali of certain normality such as 0.1N sodium hydroxide.
Alkalis :

They are substances, which give hydroxyl ions in solution. Also, they are potential proton acceptors.
The strength of an alkali depends on the degree of dissociation (ionization) in solution. Sodium hydroxide (NaOH) and potassium hydroxide (KOH) are completely dissociated so, they are strong alkalis while ammonium hydroxide (NH4OH) is a weak alkali as it slightly

Bases :

They are substances, which can accept protons .
All alkalis are bases but not all bases are alkalis.
Thymine, uracil, adenine and guanine are examples of bases.

Amphoteric Substances :

They are substances that can act as acids in certain media and as alkalies in other media. Proteins and amino acids are examples for amphoteric substances.

Buffers

They are solutions that resist appreciable changes in pH when an acid or alkali is added to it. Biological systems use buffers to control pH.
Buffers may be formed of one of the following mixtures:

Weak acid and its salt with strong base e.g. carbonic acid and sodium bicarbonate.
Weak base and its salt with strong acid e.g. ammonium hydroxide and ammonium chloride.
Mechanism of Buffer Action :

In the case of bicarbonate buffer (NaHCO3 / H2CO3)
If strong alkali as sodium hydroxide (NaOH) is added it will react with the carbonic acid forming sodium bicarbonate, which is a weak alkali. So there is no appreciable change in the pH.



If strong acid as hydrochloric acid (HCl) is added it will react with the sodium bicarbonate forming carbonic acid, which is a weak acid. So there is no appreciable change in the pH.



Buffer Systems of The Blood :

Plasma Buffers
They include:
a) Bicarbonate buffer (NaHCO3 / H2CO3)
b) Phosphate buffer (Na2HPO4 / NaH2PO4)
c) Plasma proteins (Na proteinate / H protein)
RBCs Buffers
They include:
a) Bicarbonate buffer (KHCO3 / H2CO3)
b) Haemoglobin buffer (KHb /HHb)
c) Oxyhaemoglobin buffer (KHbO2 /HHbO2)
Clinical Significance of Blood Buffers :

The pH of blood is maintained in a narrow range around 7.4 (7.37 - 7.43) by the buffer systems in blood plasma and red blood cells.

A decrease in blood pH, which can be compensated, is called acidosis while uncompensated decrease in blood pH is called acidaemia.

An increase in blood pH, which can be compensated, is called alkalosis while uncompensated increase in blood pH is called alkalaemia.

Even relatively small changes in this value of blood pH can lead to severe metabolic consequences. Therefore, blood buffers are extremely important in order to maintain homeostasis.

Solutions

A solution is formed when a substance is dissolved in a solvent.

Any solution is formed of 2 phases; the solute which is also called dispersed phase and solvent which is also called dispersion phase

Types of Solutions:
Solutions can be classified according to the concentration or the size of solute particles

A- According to The Concentration of Solute Particles, solutions may be:

Normal Solution
It is a solution, which contains the equivalent weight of the solute in grams dissolved in one liter of the solution.
Molar Solution
It is a solution, which contains the molecular weight of the solute in grams dissolved in one liter of the solution.
Molal Solution
It is a solution, which contains the molecular weight of the solute in grams dissolved in one kilogram of the solvent.
B- According to The Size of Solute Particles, solutions may be:

Crystalloids (True solutions)
They are solutions in which the size of solute particles is less than 1 m (millimicron) e.g. sodium chloride solution.
Colloids
These are solutions in which the size of solute particles ranges from 1 to . e.g. plasma proteins
Suspensions
They are solutions in which the size of solute particles is more than e.g. suspension of sand in water.
Colloids

These are solutions in which the size of solute particles ranges from 1 to . e.g. plasma proteins.

Types of Colloids :

According to their ability to take up the solvent, colloids are classified into emulsoids and suspensoids

1- Emulsoids

- They are lyophilic (solvent loving) colloids. If the solvent is water they are called hydrophilic colloids.
- They are more viscid.
- They are more stable and not easily precipitated as the solute particles are surrounded by 2 stability factors:

i- Negative or positive charge.
ii- A shell (layer) of solvent.


They can be precipitated by dehydration followed by neutralization of the charge.

Examples for emulsoids include protein, starch and egg white solutions.

2- Suspensoids

- They are lyophobic (solvent hating) colloids. If the solvent is water they are called hydrophobic colloids.
- They are less viscid than emulsoids.
- They are less stable and easily precipitated as the solute particles are surrounded by one stability factor that is negative or positive charge.


- They can be precipitated by neutralization of the charge.
Examples for emulsoids include colloidal gold and colloidal iron solutions.

The Charge on The Colloidal Particles may be due to :

Ionization of some groups on the surface of colloidal particles.
Ions from the solvent may be adsorbed on the surface of the colloidal particles.
Stability Factors of Colloids :

Presence of negative or positive charge on the surface of the colloidal particles.
Presence of a shell (layer) of solvent surrounding the particles of emulsions.
The size of colloidal particles; the smaller the size, the more stable the colloid.
Separation of Colloids :

Ultracentrifugation, which means centrifugation of the solution at a very high speed.
Ultrafiltration, which means forcing the solution, under pressure, through membranes of varying porosity.
Electrophoresis, which means migration of colloidal particles in electric field.
Precipitation of colloids
Colloids can be Precipitated by:

Dehydrating agents.
Strong agitation, freezing and heating.
Colloids of opposite charge.

Physical Chemistry >> Solution, Solution phenomena >> Lecture 4
[/color]http://osp.mans.edu.eg/medbiochem_mi/cources/biochemistry/


Physical Chemistry >> Osmosis, Expression of Concentration>> Lecture 5
http://osp.mans.edu.eg/medbiochem_mi/cources/biochemistry/

ORGANIC CHEMISTRY

Natural Products
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Organic chemistry had its genesis in the study of naturally occurring substances, and this remains a constant source of information and intellectual challenge. In this section some of the most interesting and important classes of natural products will be discussed.

Table of Contents

Carbohydrates
Glucose
The Structure and Configuration of Glucose
Anomeric Forms of Monosaccharides
Glycosides
Disaccharides
Polysaccharides

Lipids
Fatty Acids
Soaps & Detergents
Fats & Oils
Waxes
Phospholipids
Prostaglandins
Terpenes
Steroids
Lipid Soluble Vitamins
Biosynthetic Mechanisms

Proteins and Amino Acids
α-Amino Acids
Reactions of Amino Acids
Synthesis of Amino Acids
Peptides & Proteins
The Primary Structure of Peptides
Secondary & Tertiary Structure of Large Peptides and Proteins
Peptide Synthesis

Nucleic Acids
The Primary Structure of DNA
The Secondary & Tertiary Structures of DNA
RNA and Protein Synthesis



Carbohydrates

Chemistry of Carbohydrates
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Carbohydrates or saccharides are the most abundant class of biomolecules. Carbohydrates are used to store energy, though they serve other important functions as well. This is an overview of carbohydrate chemistry, including a look at the types of carbohydrates, their functions, and carbohydrate classification.

What Is a Carbohydrate?

Carbohydrates are a common class of simple organic compouds. A carbohydrate is an aldehyde or a ketone that has additional hydroxyl groups. The simplest carbohydrates are called monosaccharides, which has the basic structure (C·H2O)n, where n is three or greater. Monosaccharides link together to form oligosaccharides and polysaccharides. Two monosaccharides link together to form a disaccharide.

Functions of Carbohydrates

Carbohydrates serve several biochemical functions:
Monosaccharides are a fuel for celular metabolism.
Monosaccharides are used in several biosynthesis reactions.
Monosaccharides may be converted into space-saving polysaccharides, such as glyocogen and starch. These molecules provide stored energy for plant and animal cells.
Carbohydrates are used to form structural elements, such as chitin in animals and cellulose in plants.
Carbohydrates and modified carbohydrates are important for an organism's fertilization, development, blood clotting and immune system function.

Examples of Carbohydrates

Monosaccharides: glucose, fructose, galactose
Disaccharides: sucrose, lactose

Polysaccharides: chitin, cellulose

Carbohydrate Classification

Three characteristics are used to classify monosaccharides:

number of carbon atoms in the molecule
location of the carbonyl group
the chirality of the carbohydrate
aldose - monosaccharide in which the carbonyl group is an aldehyde
ketone - monosaccharide in which the carbonyl group is a ketone

triose - monosaccharide with 3 carbon atoms

tetrose - monosaccharide with 4 carbon atoms

pentose - monosaccharide with 5 carbon atoms

hexose - monosaccharide with 6 carbon atoms

aldohexose - 6-carbon aldehyde (e.g., glucose)

aldopentose - 5-carbon aldehyde (e.g., ribose)

ketohexose - 6-carbon hexose (e.g., fructose)

A monosaccharide is D or L depending on the orientation of the asymmetric carbon located furthest from the carbonyl group. In a D sugar, the hydroxyl group is on the right the molecule when written as a Fischer projection. If the hydroxyl group is on the left of the molecule, then it is an L sugar.

Lipids
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The lipids are a large and diverse group of naturally occurring organic compounds that are related by their solubility in nonpolar organic solvents (e.g. ether, chloroform, acetone & benzene) and general insolubility in water. There is great structural variety among the lipids, as will be demonstrated in the following sections. You may click on a topic listed below, or proceed page by page.

Fatty Acids
Soaps and Detergents
Fats and Oils
Waxes
Phospholipids
Eicosonoids
Terpenes
Steroids
Lipid Soluble Vitamins
Biosynthetic Pathways


Fats, Oils, Waxes & Phospholipids

1. Fatty Acids

The common feature of these lipids is that they are all esters of moderate to long chain fatty acids. Acid or base-catalyzed hydrolysis yields the component fatty acid, some examples of which are given in the following table, together with the alcohol component of the lipid. These long-chain carboxylic acids are generally referred to by their common names, which in most cases reflect their sources. Natural fatty acids may be saturated or unsaturated, and as the following data indicate, the saturated acids have higher melting points than unsaturated acids of corresponding size. The double bonds in the unsaturated compounds listed on the right are all cis (or Z).

(IMMUNOGLOBULINS (ANTIBODIES
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Each antibody binds to a specific antigen; an interaction similar to a lock and key.

An antibody (Ab), also known as an immunoglobulin (Ig), is a large Y-shaped protein produced by B-cells that is used by the immune system to identify and neutralize foreign objects such as bacteria and viruses.
The antibody recognizes a unique part of the foreign target, called an antigen.[1][2] Each tip of the "Y" of an antibody contains a paratope (a structure analogous to a lock) that is specific for one particular epitope (similarly analogous to a key) on an antigen, allowing these two structures to bind together with precision.
Using this binding mechanism, an antibody can tag a microbe or an infected cell for attack by other parts of the immune system, or can neutralize its target directly (for example, by blocking a part of a microbe that is essential for its invasion and survival). The production of antibodies is the main function of the humoral immune system.[3]
Antibodies are secreted by a type of white blood cell called a plasma cell. Antibodies can occur in two physical forms, a soluble form that is secreted from the cell, and a membrane-bound form that is attached to the surface of a B cell and is referred to as the B cell receptor (BCR). The BCR is only found on the surface of B cells and facilitates the activation of these cells and their subsequent differentiation into either antibody factories called plasma cells, or memory B cells that will survive in the body and remember that same antigen so the B cells can respond faster upon future exposure.
[4] In most cases, interaction of the B cell with a T helper cell is necessary to produce full activation of the B cell and, therefore, antibody generation following antigen binding.[5] Soluble antibodies are released into the blood and tissue fluids, as well as many secretions to continue to survey for invading microorganisms.
Antibodies are glycoproteins belonging to the immunoglobulin superfamily; the terms antibody and immunoglobulin are often used interchangeably.[6]
Antibodies are typically made of basic structural units—each with two large heavy chains and two small light chains.
There are several different types of antibody heavy chains, and several different kinds of antibodies, which are grouped into different isotypes based on which heavy chain they possess.
Five different antibody isotypes are known in mammals, which perform different roles, and help direct the appropriate immune response for each different type of foreign object they encounter.[7]
Though the general structure of all antibodies is very similar, a small region at the tip of the protein is extremely variable, allowing millions of antibodies with slightly different tip structures, or antigen binding sites, to exist. This region is known as the hypervariable region. Each of these variants can bind to a different antigen.[1]
This enormous diversity of antibodies allows the immune system to recognize an equally wide variety of antigens.[6] The large and diverse population of antibodies is generated by random combinations of a set of gene segments that encode different antigen binding sites (or paratopes), followed by random mutations in this area of the antibody gene, which create further diversity.[7][8]
Antibody genes also re-organize in a process called class switching that changes the base of the heavy chain to another, creating a different isotype of the antibody that retains the antigen specific variable region. This allows a single antibody to be used by several
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TISSUE CHEMISTRY

http://www.chemistrydaily.com/chemistry/Connective_tissue

Connective tissue
Connective tissue is any type of biological tissue with an extensive extracellular matrix. There are several basic types:

Bone contains specialized cells called osteocytes embedded in a mineralized extracellular matrix, and functions for general support.
Loose connective tissue holds organs and epithelia in place, and has a variety of proteinaceous fibers, including collagen and elastin.
Fibrous connective tissue forms ligaments and tendons.
In the osteichthyes, cartilage makes up virtually the entire skeleton. In most other vertebrates, it is found primarily in joints, where it provides cushioning. The extracellular matrix of cartilage is composed primarily of chondroitin sulfate.
The extracellular matrix of blood is the blood plasma, which transports dissolved nutrients, hormones, and carbon dioxide in the form of bicarbonate. The main cellular component is red blood cells.
Adipose tissue contains adipocytes, used for cushioning, insulation, and energy storage.

Disorders of connective tissue

: Various connective tissue conditions have been described, these can be both inherited and environmental.
Marfan syndrome - a genetic disease causing abnormal fibrillin.
Scurvy - caused by a dietary deficiency in vitamin C, leading to abnormal collagen.
Ehlers-Danlos Syndrome - a genetic disease causing progressive deterioration of collagens, with different EDS types affecting different sites in the body, such as joints, heart valves, organ walls, arterial walls, etc.

ادخل
http://www.ncbi.nlm.nih.gov/pmc/articles/PMC1258307/