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Property and composition of solutions

June 21, 2024

Property and composition of solutions. Quantitative composition of solutions. To prepare solutions. Colligative properties of solutions.

Physical and chemical characteristics of water.

Water (H₂O) is composed of two atoms of hydrogen and one of oxygen. Each hydrogen atom is linked to the oxygen atom by a single covalent bond. Because oxygen is more electronegative than hydrogen, there is a separation of charge within the molecule. The electron distribution in oxygen-hydrogen bonds may therefore be described as polar or asymmetrical. If water molecules were linear, then the bond polarities would cancel each other out, and water would be nonpolar. However, water molecules have a bent geometry with a bond angle of 104.5°

Molecules such as water, which have an unbalanced distribution of charge, are called dipoles. Such molecules have opposite charges on two points. When molecular dipoles are subjected to an electric field, they orient themselves in the direction opposite to that of the field.

Water’s properties are directly related to its molecular structure.

One consequence of the large difference in electronegativity of hydrogen and oxygen is that the hydrogens of one water molecule are attracted to the unshared pairs of electrons of another water molecule. This noncovalent relationship is called a hydrogen bond. In addition to hydrogen bonds, three other types of noncovalent interactions play important roles in determining the capacity of water to interact with other types of molecules. These are electrostatic interactions, van der Waal’s forces, and hydrophobic interactions. Because biological reactions take place in a water medium, an understanding of noncovalent bonding is important.

Covalent bonds between hydrogen (electropositive atoms) and oxygen are polar. For example, each of the two hydrogens in water molecules will be weakly attracted to oxygen atoms in other nearby water molecules. The resulting intermolecular “bonds” act as a bridge between adjacent molecules. Although considerably weaker than ionic and covalent bonds, hydrogen bonds are stronger than most other types of noncovalent bonds.

Electrostatic interactions occur between oppositely charged atoms or groups. An important aspect of all electrostatic interactions in aqueous solution is the hydration of ions that occurs. Because water molecules are polar, they are attracted to charged ions. Shells of water molecules, referred to as solvation spheres, cluster around both positive and negative ions. As ions become hydrated, the attractive force between them is reduced, and the charged species dissolves in the water. Water, sometimes called the universal solvent.

Melting point of water – 0 ^oC; boiling point – 100 ^oC.

Water plays an important role in the thermal regulation of living organisms. Water’s high heat capacity coupled with the high water content found in most organisms (between 50% and 95%, depending on species) contributes to the maintenance of an organism’s internal temperature. The evaporation of water is used as a cooling mechanism, since it permits large losses of heat. For example, an adult human may eliminate daily as much as 1200 g of water in expired air, sweat, and urine. The associated heat loss may amount to approximately 20% of the total heat generated by metabolic processes.

Water is a remarkable solvent. Water’s ability to dissolve a large variety of ionic and polar substances is determined by its dipolar structure and its capacity to form hydrogen bonds. Salts such as sodium chloride (NaCI) are held together by ionic (or electrostatic) forces. They dissolve easily in water because dipolar water molecules are attracted to the Na⁺ and Cl^– ions.

Organic molecules with ionizable groups and many neutral organic molecules with polar functional groups also dissolve in water. Their solubility is due primarily to the hydrogen bonding capacity of water. Nonpolar compounds are not soluble in water. Because they lack polar functional groups, such molecules cannot form hydrogen bonds.

Liquid water molecules have a limited capacity to ionize to form a hydrogen ion (H⁺) and a hydroxide ion (OH^–). (H⁺ does not actually exist in aqueous solution. In water a proton combines with a water molecule to form the hydrated hydrogen ion, H₃O⁺, commonly called a hydronium ion. For convenience, however, the hydrated proton is usually represented as H⁺.)

The state and distribution of water in the organism.

There are two water compartments in the body:

1. Intracellular water

2. Extracellular water

Extracellular fluid is divided into:

1. interstitial fluid

2. plasma

Distribution of water in an adult man, weighing 70 kg

Compartment

Body weight (%)

Volume (l)

Total

ICF

ECF

Interstitial fluid

10,5

Plasma

3,5

Biological role of water:

1. Water is an essential constituent of cell structures and provides the media in which the chemical reactions of the body take place and substances are transported.

2. It has a high specific heat for which, it can absorb or gives off heat without any appreciable change in temperature.

3. It has a very high latent heat. Thus, it provides a mechanism for the regulation of heat loss by sensible or insensible perspiration on the skin surface.

4. The fluidity of blood is because of water

5. Water is the most suitable solvent in human body

6. Dielectric constant : Oppositely charged particles can coexist in water. Therefore, it is a good ionizing medium. This increases the chemical reactions.

7. Lubricating action: Water acts as lubricant in the body to prevent friction in joints, pleura, conjunctiva and peritoneum

A solution is a homogeneous mixture of two or more chemically non-reacting substances whose composition can be varied within certain limits.

Solutions can be classified on the basis of their state: solid, liquid, or gas.

THE MECHANISM OF DISSOLVING. Consider what happens when we add a solid solute to a liquid solvent. Immediately after the addition, the solid-state structure begins to disintegrate, as, little by little, solvent molecules chip away at the surface of the crystal lattice, prying out solute particles, surrounding them, and finally dispersing them throughout the body of the solution. The ease with which all this takes place depends primarily on the relative strengths of three attractive forces: (1) the forces between the particles of the solute before it has dissolved {solute-solute forces), (2) the forces between solvent particles before dissolution has taken place {solvent-solvent forces), and (3) the forces that are formed between solute and solvent particles during the dissolving process {solute-solvent forces). As dissolving takes place, solute-solute and solvent-solvent forces are replaced by solute-solvent forces. In general, high solubility of a solid in a liquid is favored by weak solute-solute forces (measured in solids by lattice energy) and by strong solute-solvent forces (measured by solvation energy; see below). A generalization that is old but still very useful is “like dissolves like,” which means that a solvent will dissolve a solute if they have similar properties. More specifically, polar solvents tend to dissolve polar and ionic solutes, and nonpolar solvents tend to dissolve nonpolar solutes. Although this rule is not perfect, it is very useful as a rough guideline for predicting solubilities. Let us see how it can be justified in terms of the structural changes that take place during dissolution.

NONPOLAR SOLVENTS. First, consider carbon tetrachloride, a nonpolar solvent. (CCl₄ has no dipole moment because it is tetrahedral.) Those solutes that have high solubilities in CCl₄ are also nonpolar, such as iodine (I₂) and sulfur (S₈). In nonpolar substances, the only intermolecular attractions are comparatively weak London forces. Thus, when a solution is formed from two nonpolar substances, solute-solute and solvent-solvent forces are easily replaced by solute-solvent forces. (Here, “easily” means that there is no great energy deficit.) Wheonpolar solute molecules are inserted among nonpolar solvent molecules, neither experiences a great change in environment, and so the solution is readily formed. On the other hand, nonpolar liquids tend to be poor solvents for polar and ionic compounds. Lattice energies in these compounds tend to be higher than ionpolar solutes, and because of the lack of polarity of the solvent molecules, solute-solvent forces are too weak to overcome solute-solute forces. All other factors being equal, polar compounds tend to be less soluble ionpolar liquids than are nonpolar compounds, and ionic compounds tend to be still less soluble.

POLAR SOLVENTS. Now, consider the solvent properties of a polar solvent, water. Water is a good solvent for polar solutes, such as hydrogen chloride (HC1), and for ionic solutes, such as sodium chloride (NaCI). In these cases, the establishment of strong solute-solvent attractions provides the energy needed to overcome strong attractive forces in the solute. Water is an especially good solvent for many ionic compounds because the ion-dipole forces established in the solution are strong (recall that water is a highly polar molecule), and many H₂O molecules cluster around each ion. However, water is usually a poor solvent for nonpolar substances. This is true because the solvent-solvent (dipole-dipole) attractions in water are so strong that they cannot be overcome by weaker solute-solvent interactions, such as those provided by a nonpolar solute. (The formation of weak solute-solvent attractions does not release enough energy to compensate for that required to separate the solvent dipoles.)

Although the like-dissolves-like rule is a handy generalization, it tends toward oversimplification. For example, not all ionic compounds are highly soluble in water. Although barium sulfate, BaSO₄, is an ionic compound, its solubility in water is very low. Apparently the 2 + and 2 – ions in BaSO₄ cause the lattice energy in BaSO₄ to be so high that the solubility of this compound in water is low. Evidently we must look more closely at the specific interactions in solute, solvent, and solution. Another case in point: The ethanol molecule (C₂H₅OH; dipole moment, 1.70 D) is less polar than ethyl chloride (C₂H₅Cl; 2.05 D), yet ethanol is miscible with water in all proportions (the term infinitely soluble is often used), but the solubility of ethyl chloride in water is extremely low. Here, again, considering only polarity of solvent and solute leads to an error; specific strong interactions between ethanol and water (described below) account for their mutual solubility.

Gas solution is not possible to prepare a heterogeneous mixture of two gases because all gases mix uniformly with each other in all proportions. Gaseous solutions have the structure that is typical of all gases. (The molecules are spaced far apart and are in rapid, random motion, colliding frequently with each other and with the walls of the container.) The only difference between a gaseous solution and a pure gas is that in a solutioot all the molecules are alike. Air, the gaseous solution with which we come in closest contact, is composed primarily of N₂ (78 % by volume), O₂ (21 %), and Ar (1 %), with smaller concentrations of CO₂, H₂O, Ne, He, and dozens of other substances at very low levels.

Liquid solutions have the internal structure that is typical of pure liquids: closely spaced particles arranged with little order. Unlike a pure liquid, however, a liquid solution is composed of different particles. Much of this chapter is devoted to the properties of liquid solutions, and special emphasis is given to aqueous solutions, in which the major component is water.

Two kinds of solid solutions are common. The first, the substitutional solid solution, exhibits a crystal lattice that has structural regularity but in which there is a random occupancy of the lattice points by different species. For example, mixtures of potassium chloride and potassium bromide ranging in composition from pure KC1 to pure KBr can be crystallized out of aqueous solution. For each intermediate composition, a single crystalline phase with the rock salt structure is formed, but it has Cl^– and Br^– ions randomly distributed at the anion sites. The radii of Cl^– and Br^– are sufficiently similar to allow one ion to substitute for the other. Similarly, I^– ions can be incorporated into the rock salt structure of AgBr by substitution for Br^– ions, but in this case a maximum of only about 70 % of the Br^– ions can be replaced.

In the second type of solid solution, the interstitial solid solution, foreign atoms, ions, or molecules occupy the nooks and crannies, the interstices, in the host lattice. One example of this is austenite, in which carbon atoms occupy some of the interstices in a face-centered cubic array of iron atoms. Most of the transition elements can form such solid solutions with small atoms such as H, B, C, and N.

In many solutions, one component is present in considerable excess over the others. This component is called the solvent, and other components are called solutes. For example, in the case of a solution prepared by dissolving 1 g each of table salt and table sugar in 100 g of water, we refer to the water as the solvent and the salt and sugar as solutes. As a practical matter, it is often useful to consider the solvent as the component throughout which the particles of the solute(s) are randomly dispersed.

The terms concentrated and dilute are commonly used to give a qualitative indication of the concentration of the solute in a solution; “concentrated” implies a relatively high concentration of solute, “dilute” a relatively low one.

The composition of a solution is described quantitatively by specifying the concentrations of its components. Commonly used concentration units include mole fraction, mole percent, molarity, molality, percent by mass, and normality. Definitions and examples of each of these units follow, except for normality.

Concentration units of solution

Mass fraction (w_i) of solute in solution is the ratio of the mass solute (m_i) to the mass of solution m_i +m_s; m_s– mass of a solvent:

Percentage by weight (mass) or mass percent, is the quantity of one component of a solution expressed as a percentage of the total mass:

where w_m – percent by mass,

m_A, m_B, m_C – mass of components in the solution.

Eexample: What is the percent by mass of the sucrose in the solution of Example1?

Solution:

Mole (n) is quantity of matter, contains such quantity atoms, molecules, ions, as is atoms in 0,012 kgs of an isotope Carbon ₁₂C.

Molar mass of compound is the ratio of mass compound to numbers mole compound. It is equal to a ratio of matter mass (m) to quantity moles its (n), gram/ mol:

The mole fraction (X) of a component in solution is the ratio of the number of moles of that component to the total number of moles of all components. Letting n represent number of moles and designating different components as A, B, C, . . ., we can write

and

and so on. Note that X _A + X_B + X_C + X_D + X_F = 1

Example: If 28.6 g of sucrose (cane sugar, C₁₂H₂₂O₁₁) is dissolved in 101.4 g of water, what is the mole fraction of sucrose in the solution? What is the mole percent sucrose in the solution?

Solution: The molecular mass of C₁₂H₂₂O₁₁ is 342.3, and so 28.6 g of sucrose is

The molecular mass of H₂O is 18.02, and so 101.4 g of water is

Therefore,

mol % _sucrose = X_sucrose ×100 = 1.46×10^-2×100 = 1.46% (The solution is 1.46 mol % sucrose and 98.54 mol % water.)

A concentration unit closely related to mole fraction is mole percent, which is the number of moles of one component expressed as a percentage of the total number of moles present:

(Note that this is simply mole fraction times 100).

Volume fraction of a component in solution is the ratio of the volume solute (V_i) to the total volume of solution (V_i + V_s), V_s – volume solvent:

Mass concentration, titer (T) is number grams of solute (m_i) per one milliliter of solution (V_s). Or it is the ratio of the quantity grams of solute and volume solution:

Molarity (C), or molar concentration, is the number of moles of solute dissolved per liter of solution. If n_A represents the number of moles of solute A, and V the volume of the solution in liters, then

Where:- number mole solute;

m – mass solute, grams;

M_B – molar mass solute, in grams/mole.

A solution that contains 1.67 mol of solute dissolved in a volume of 1.00 l is referred to as a 1.67-molar (written 1.67 M) solution.

The molarity expresses by mol per liter, or mol per m³, mol per sm³.

Example: 10.0 g of ascorbic acid (vitamin C, H₂C₂H₆O₆) is dissolved in enough water to prepare 125 ml of solution. What is the molar concentration (molarity) of this compound in the solution?

Solution: The molecular mass of H₂C₂H₆O₆ is 176.1. Therefore, 10.0 g of ascorbic acid is

The volume of the solution is 125 ml, or 0.125 l. The molarity is therefore

Molality is defined as the number of moles of solute dissolved per kilogram of solvent. Thus, the molality of solute A in a solution is

where n_A is the number of moles of A, and m_solvent is the mass of solvent in kg

A solution containing 3 mol of solute per kilogram of solvent is said to be 3-molal, written 3 m.

Example: If 28.6 g of sucrose (cane sugar, C₁₂H₂₂O₁₁) is dissolved in 101.4 g of water, what is the molality?

Solution: The molecular mass of C₁₂H₂₂O₁₁ is 342.3, and so 28.6 g of sucrose is

We know the number of moles of sucrose, and the mass of the water is 101.4 g = 0.1014 kg. Therefore, the molality of the sucrose is

In measure analysis for the characteristic the composition of solution will use molar mass of an equivalent (equivalent mass)

Molar mass of an equivalent of element is the mass of the element which combines with or displaces 1.008parts by mass of hydrogen or 8 part by mass of oxygen or 35.5 parts by mass of chlorine:

E = f_equiv · M_B

The factor of equivalence (f_equiv) – number, which is demonstrated which part of matter (equivalent) can react with one atom of Hydrogen, or one electron in reduction reactions.

Molar concentration of an equivalent (normal concentration), normality is quantity gram-equivalent of solute per one liter of solution (V_s):

Example: What volume should be taken solution HCl with mass fraction 4 % and density 1,018 g / m³ for making 250 mls 0,12 mol/l of solution?

The first of all we calculate normal concentration 4 % solution of HCІ by the formula:

After that we calculate a volume its solution:

Coligative properties of biological liquids.

The water – main component of human organisms and also is part of medium, in which lives the people. The main water property is solved a lot of matters with formatted solutions.

The water in organisms of the person, animal, plant is by its constituent (in a yumrn’s organism about 70 -80 % of water), solvent, and also participates in exchange reactions of matters (hydrolysis, hydration, swelling (turgescence), digestion). It executes a role of a transport system in processes of a feeding, carry of enzymes, products of a metabolism, gases, antibodies. The water is supported a condition to a homeostasis in an organism of the person (acid – alkaline, osmotic, hemodinamicil, thermal equilibrium). The water is indispensable for secrets iones, maintenance of a turgor of cages.

SOLYTIONS. TYPES OF SOLYTIONS.

The solution is homogeneous thermodynamic nonperishable systems, which consist with two or stable more components.

Distinguish gaseous, fluid and solid solutions. The gas solutions are mixture of gases. Air is solution of gases of azote, oxygen, carbonic oxide (IV), water vapour and inert gases. The slurries are mixture of liquids, or solutions of solid matters and gases in liquids. Solid solutions are solid phases, which one were derivated at cooling of infrequent melts.

A true solution it is a homogeneous system of a changeable structure derivated two and more components an elemental Composition and physical characteristicss of solution is identical in all volume of solution. The characteristics of solution are its structure and concentration.

Mass fraction (w_i) of solute in solution is the ratio of the mass solute (m_i) to the mass of solution m_i +m_s; m_s– mass of a solvent:

Percentage by weight (mass) or mass percent, is the quantity of one component of a solution expressed as a percentage of the total mass:

where w_m – percent by mass,

m_A, m_B, m_C – mass of components in the solution.

Mole (n) is quantity of matter, contains such quantity atoms, molecules, ions, as is atoms in 0,012 kgs of an isotope Carbon ₁₂C.

Molar mass of compound is the ratio of mass compound to numbers mole compound. It is equal to a ratio of matter mass (m) to quantity moles its (n), gram/ mol:

and

and so on. Note that X _A + X_B + X_C + X_D + X_F = 1

A concentration unit closely related to mole fraction is mole percent, which is the number of moles of one component expressed as a percentage of the total number of moles present:

(Note that this is simply mole fraction times 100).

Volume fraction of a component in solution is the ratio of the volume solute (V_i) to the total volume of solution (V_i + V_s), V_s – volume solvent:

Mass concentration, titer (T) is number grams of solute (m_i) per one milliliter of solution (V_s). Or it is the ratio of the quantity grams of solute and volume solution:

Where:- number mole solute;

m – mass solute, grams;

M_B – molar mass solute, in grams/mole.

A solution that contains 1.67 mol of solute dissolved in a volume of 1.00 l is referred to as a 1.67-molar (written 1.67 M) solution.

The molarity expresses by mol per liter, or mol per m³, mol per sm³.

Molality is defined as the number of moles of solute dissolved per kilogram of solvent. Thus, the molality of solute A in a solution is

where n_A is the number of moles of A, and m_solvent is the mass of solvent in kg

A solution containing 3 mol of solute per kilogram of solvent is said to be 3-molal, written 3 m.

In measure analysis for the characteristic the composition of solution will use molar mass of an equivalent (equivalent mass),

Molar mass of an equivalent of element is the mass of the element which combines with or displaces 1.008parts by mass of hydrogen or 8 part by mass of oxygen or 35.5 parts by mass of chlorine:

E = f_equiv · M_B

The factor of equivalence (f_equiv) – number, which is demonstrated which part of matter (equivalent) can react with one atom of Hydrogen, or one electron in reduction reactions.

Molar concentration of an equivalent (normal concentration), normality is quantity gram-equivalent of solute per one liter of solution (V_s):

MODERN THEORY ABOUT THE NATURE OF SOLUTIONS

The medical men are especially interested for liquors. The biological liquids (blood, the lymph, urine), wich is by complex mixtures of proteins, lipids, carbohydrates, salts. Physic-chemical regularity of interplay these miscellaneous behind properties and sizes of fragments both between itself, and with water moleculas ambient them, is extremely relevant for habitability of an organism.

During development of the doctrine about solutions two theories are designed: chemical and physical.

According to the physical theory (S. Arrenyus, V. Osvald, Ye. Vant-Hoff), the process of dissolution is esteemed as an even distribution particles of solvend in all volume of solution. The solvent is by inert medium, the moleculas of solvend and solvent do not interact between themselves.

The chemical theory (D. I. Mendelaev, I.A. Kablucov, M. S. Kurnacov) regarded solution as systems, which one were derivated from parts of solvend, solvent and non-persistent chemical combinations, which one will be derivated in solution with the help of hydrogen bindings, or electrostatic attractive forces at interplay particles of a solvent and solvend.

The modern theory of solutions integrates the physical and chemical theories regarded process formation of solutions as interplay between particles of different polarity.

PROCESSES OF DISSOLUTION AND DISSOLVING CAPACITY OF WATER

The water well solves many matters, its dissolving capacity is determined by high inductivity (e= 78). The moleculas of water enter with ions of salts ion – dipole interplay, derivating hydrated ions.

The hydration of ions results in their stabilization in solution. Each negatively ionized atom attracts positive poles of dipoles of water and aims to hold them about itself. The positively ionized atoms hydrated even more strongly, as the charge density for them is higher.

The dissolubility of organic matters in water depends on availability of polar groups (having large affinity to water) in moleculas; such groups are called hydrophylic, are capable to interact with water (formation of hydrogen bindings).

In common cases dissolubility of matters are subject to the rule of thumb “similar solves in similar”. For example, the hydrocarbon a naphthalene well solves in gasoline (mixture of hydrocarbon), it is worse – in alcohol (names hydrocarbon radicals), almost water soluble.

The dissolubility of solid matters depends on temperature. For the majority of salts she increases with temperature rise.

The process of dissolution is conditioned by interplay of fragments of solved matter and solvent. The dissolution of solid matters in water and molecular dissociation on ions can be presented as follows: the dipoles of water, falling in an electrical field of polar moleculas, will orient around of polar groups or ions, which one are on a surface of a chip. Attracting to themselves of a molecula or ions, the dipoles of water relax, and then lacerate intermolecular or ionic bonds. For example, the water reduces strength of an ionic bond in a chip of sodium salt, between ions Na⁺and Cl^–. The separate moleculas or ions under influencing of heat motion of all fragments of solution are displaced (diffuse), uniformly being arranged between moleculas of a solvent. At dissolution often descends not only gap of bonds in solute, but also the associates of moleculas of a solvent are blasted. In solution there are new associates as well from moleculas of a solvent, as from fragments of solute.

The processes of dissolution are accompanied by heat effects, which depend on energy consumption on transfer of solid matter in a liquid state (that is its dissolution), and from a secured heat at interplay of fragments of solved matter and solvent (solvation). If quantity of energy expended for a gap of bonds in a chip, is more than a heat generated of a solvation, dissolution – endothermal process. At minor heat of dissociation on ions and large heat of a solvation – exothermal process.

At formation of solutions the polarity of moleculas of a solvent matters, than their electrical dipole moment more, the above their capacity to a solvation, the is more excreted of energy and better matter solves. Capacity of matters will be dissolved in solvent liquids, as depends on polarity of their moleculas.

The polar matters well solve in polar solvents, (for example, sodium salt in water). To the contrary, the matter with low-polarity by moleculas, for example, fatty acids, solve in solvents, molecula which one poorly polar (ethers) or not polar (benzole) better.

An unsaturated solution is one in which the concentration of solute is less than its concentration in a saturated solution. (Additional solute can be dissolved in an unsaturated solution, until the solution becomes saturated.)

A supersaturated solution is one in which the concentration of solute is greater than its concentration in a saturated solution. A supersaturated solution is unstable and its solute tends eventually to crystallize out of solution, much as a supercooled liquid tends eventually to crystallize. Solubility equilibrium is not possible in a supersaturated solution. If more solute is added, crystallization occurs, usually rapidly, as solute leaves the solution to crystallize on the surfaces of the crystals of added solute. (The situation is very much like the rapid freezing of a supercooled liquid brought about by adding a seed crystal.)

The solubility of a solute in a given solvent is defined as the concentration of the saturated solution. At 25°C, urea (see above) can be dissolved in water until a total of 19 mol has been added for each liter of solution formed. Some solutes are infinitely soluble in a given solute (this means that solute and solvent will mix in all proportions), while others have solubility’s so low that they are not measurable by direct methods. Although there is probably no such thing as complete insolubility, the term insoluble is commonly applied to a substance whose solubility is extremely low.

SOLUTIONS OF GASES IN LIQUIDS

Henry’s Law: The solubility of a gas dissolved in a liquid is proportional to the partial pressure of the gas above the liquid.

This is a statement of Henry’s law, which can be written

X = KP

where X is the equilibrium mole fraction of the gas in solution (its solubility), P is its partial pressure in the gas phase, and K is a constant of proportionality, usually called the Henry’s-law constant.

The partial pressure is a part of common pressure, which one is a share of each gas in gas mixture.

Henry’s law applies only when the concentration of the solute and its partial pressure above the solution are both low, that is, when the gas and its solution are both essentially ideal, and when the solute does not interact

Oo-bottoms, where the external pressure increases, the dissolubility of gases in a blood is augmented. At fast ascent from depth the dissolubility sharply decreases, they are excreted by the way is bubble and seal vessels – aeroembolism.

Properties of a solution which depend only on the concentration of the solute and not upon its identity are known as colligative properties. These include vapor-pressure lowering, boiling-point elevation, freezing-point depression, and osmotic pressure. Each of these properties is a consequence of a decrease in the escaping tendency of solvent molecules brought about by the presence of solute particles. Escaping tendency is the tendency shown by molecules to escape from the phase in which they exist.

Osmosis.

Suppose а concentrated solution of copper sulphate (deep blue in colour) is placed in а beaker and water (or а dilute solution of copper sulphate) is added slowly along the walls of the beaker without much disturbing the concentrated copper sulphate solution. The two layers are more or less well defined. Now if the beaker is allowed to stand, it is observed that after а few days, the solution in the beaker becomes uniformly blue throughout. This must be obviously due to the fact that the particles of the solute (Cu⁺² and SO₄^-2 ions) move slowly into the solvent and the molecules of the solvent (water) move into the copper sulphate solution. In other words, the particles of the solute and solvent mix spontaneously into each other.

Now suppose the experiment is performed in а slightly different manner. Suppose the beaker is divided into two compartments, by а semi-permeable membrane i.е. а membrane which allows the solvent molecules to pass through but not the solute particles. Suppose again that copper sulphate solution is placed in one compartment and water in the other. It is observed that the level on the solution side begins to rise. This must be obviously due to the fact that greater number of solvent (water) molecules from the solvent side pass into the solution side through the semi-permeable membrane than the number of solvent molecules going into the solvent from the solution through the semi-permeable membrane. Similarly, if а concentrated solution is separated from а dilute solution by а semi-permeable membrane, there is а net flow of solvent from the dilute solution to the concentrated solution through the semi-permeable membrane.

The spontaneous mixing of the particles of the solute (present in the solution) and the solvent (present above the solution) to form а homogeneous mixture is called diffusion, just as the term is used for the spontaneous mixing of gases to form homogeneous mixtures.

The net spontaneous flow of the solvent molecules from the solvent to the solution or from a less concentrated solution to а more concentrated solution through а semi-permeable membrane is called osmosis (Greek: push).

Difference between Diffusion and Osmosis. The main points of difference between diffusion and osmosis nау be summed up as given below:

Osmosis:

1. In osmosis, а semi-permeable membrane is used.

2. In this process, there is only flow of solvent molecules and that too through the semi-permeable membrane.

3. It takes place from lower concentration to higher concentration.

4. It applies to solutions only.

5. It can be stopped or reversed by applying pressure on the solution with higher concentration.

Difference:

1. In diffusion, no semi-permeable membrane is used.

2. In this process, the solvent as well as the solute molecules move directly into each other.

3. It takes place from higher concentration со lower concentration.

4. It takes place in gases as well as solutions.

5. It cannot be stopped or reversed.

Semi-permeable membranes. The semi-permeable membranes (as defined above) are of two types:

(i) Natural semi-permeable membranes е.g vegetable membranes and animal membranes which are found just under the outer skin of the animals and plants. The pig’ s bladder is the most common animal membrane used.

(ii) Artificial semi-permeable membranes. The well known examples of the artificial semi-permeable membranes are parchment paper, cellophane and certain freshly precipitated inorganic substances е.у. copper ferrocyanide, silicates, of iron, cobalt, nickel etc. The precipitated substances have to be supported on some material and this is achieved by preparing the precipitate in the walls of а porous pot.

Fig: Osmosis and measurement of osmotic pressure of а solution. (А) 1. Jar containing solution. 2. Water. Arrows show direction of movement of яа1ег. (В) Pfeffer0s apparatus for measurement of osmotic pressure. 1. Pot containing solution. 2. Water. 3. Mercury manometer (see text).

Osmotic Pressure – The upward movement of water taking place in Fig.1 (А) can be prevented if we apply mechanical force on top of the solution in the jar. The pressure just sufficient to stop osmotic pressure exerted by the solution in the jar will be the osmotic pressure exerted by the solution present in the jar. The osmotic pressure of а solution may thus be defined as the equivalent of excess pressure which must be applied, to the solution in order to prevent the passage of the solvent into it through а semi-permeable membrane separating the two, i.e. the solution and the pure solvent. As mentioned above, due to osmosis, there is а flow of solvent from the solvent to the solution or from the less concentrated solution to the more concentrated solution through the semi-permeable membrane. This flow of the solvent does not continue indefinitely. For example, consider an inverted thistle funnel at the mouth of which is tied а semi-permeable membrane (pig’ s bladder or cellophane). Suppose the thistle funnel is filled with sugar solution and then lowered into distilled water contained in а beaker. It is observed that the level of the solution inside the stem of the thistle funnel starts rising and then after some time, it becomes constant. The rise of level in the stem of the thistle funnel is obviously due to the net flow of solvent into the solution through the semi-permeable membrane. The constancy in the level shows а state of equilibrium i.е. as many molecules of the solvent eater into the solution through the semi-permeable membrane, the same number of solvent molecules from the solution go into the solvent through the semi-permeable membrane in the same time. The pressure exerted by the column h of the solution is called osmotic pressure. Thus: Osmotic pressure may be defined as the equilibrium hydrostatic pressure of the column set up as а result of osmosis.

Expression for the osmotic pressure. Osmotic pressure (Р) of а solution is found to be directly proportional to the concentration (С) of the solution and its temperature (Т). Mathematically,

Р µС; µТ; Р µС·Т or P=R·C·T

where R is а constant (called solution constant) and its value is found to be same as that of the “Gas constant”. The above equation is usually written as

Р = CRT ; p_osmotic = C_osmotic RT;

Since molarity equals the number of moles of solute (n₂) per liter of solution. V, that is. Since: C = n₂/V;

PV= nRT – van’t Hoff equation for dilute solutions.

Measurement of osmotic pressure

The osmotic pressure of а solution can be measured by many methods, but only two methods will be described.

1. Pfeffer’s method – А very simple apparatus was used by Pfeffer for this purpose. А battery pot with, а semipermeable membrane deposited in its wall is cemented to, а wide glass tube which ends in а thin tube at the top and carries а manometer in the side. The manometer is closed at its upper end and is filled with Hg and N₂. The solution under investigation is introduced into the pot through this tube, The apparatus is then made airtight by sealing off the tube at the top. А portion of the pot is immersed in distilled water kept at. а constant temperature. In the course of а few days, the manometer registers the maximum pressure, which is the osmotic pressure of the solution.

2. Freezing point determination method – It ha been found that there is а decrease of 1.8б⁰С in the freezing point of а solution when its osmotic pressure is0equal to one osmole. This method is much more rapid and accurate than Pfeffer’s method. A special apparatus is used to determine the freezing point о f the solution under investigation which is then compared with freezing point of the pure. solvent.

The decrease in the freezing point of the solution is one of the colligative properties of colloidal solutions. The other colligative properties e.g. elevation of boiling point and depression of the vapor density can also be used in the determination of the osmotic pressure of а solution.

Laws of osmotic pressure – These are the same as gas laws and apply to dilute solutions which occur in the living body.

1. The osmotic pressure is directly proportional to the concentration of the solute. For example, 1 % NaCI solution will have double the osmotic pressure of 0.5 % NaCl solution. The osmotic pressure of а solution is dependent upon only the number of dissolved or dissociated particles per unit volume and is independent of chemcial nature of particles. Thus а sodium ion (at. wt. 23), а molecule of glucose (mol. wt. 180) and а molecule of serum albumin (mol. wt. about 70,000) will exert equal osmotic pressure.

2. The osmotic pressure of а solution is directly proportional to the absolute temperature. For dilute solutions, the osmotic pressure is equal to the value CRT, where С = molar concentration, R = gas constant (0.082 liter atmosphere per degree per mole) and T = absolute temperature. А molar solution of a non-electrolyte (е.g. glucose or urea) at 0⁰С (or 273⁰ absolute) will exert an osmotic pressure equal to CRT = 1 х 0.082 x 273 = 22.4 atmospheres = 22.4 х 760 mm Hg = 17024 mm Hg.

The situation however is different for electrolytes; their molecules in solution are split into more than one particles or ions. For example, NaCl ionizes into Na⁺ and Cl^–. Thus а molar solution of NaCl will have double the number of particles as compared to а molar solution of а non-electrolyte in а given volume. As already mentioned, the osmotic pressure of а solution is dependent only on the total number of particles present in it. Therefore, а molar solution of NaCl will exert double the osmotic pressure exerted by а molar solution of а non-electrolyte. In the same way, а molar solution of Na₂SO₄ will exert an osmotic pressure equal to three times the osmotic pressure of а molar solution of а non-electrolyte. In summary, in case of еlесtrоlуtеs, osmolarity = molarity х number of particles resulting from ionization of each molecule. This, however, is only а generalization; in actual practice, the osmolarity is somewhat less than that calculated as above.

The unit of osmotic pressure described above, i.e. the osmole, (expressed as osmole per liter) which is equal to 22.4 atmospheres or 17024 mm Hg is too big for use in biology. Milliosmole which is equal to 1/1,000th of an osmole, i.e. about 17 mm Hg is usually employed; Blood plasma, gastric juice, pancreatic juice, liver bile and cerebrospinal fluid exert nearly equal osmotic pressure which is about 300 (28(Ito 295) milliosmoies/liter. The serum osmolarity can be easily found out by using the formula;

Serum osmolarity in mosm/l = 2 (Serum Na⁺ + Serum К⁺, both in mmol/l) + (Serum Glucose mg/dl)/18 + (Serum Urea mg/dl)/6. As serum К⁺ level is quite 1ow in both health hand disease as compared to serum Na⁺, it can be neglected, Another term osmolality is replacing osmolarity. Whereas osmolarity is expressed as osm/l, the osmolality is expressed as osm/kg Н₂0. In the body which has dilute fluids, both terms are virtually the same and are used interchangeably.

VALUE OSMOS IN BIOLOGICAL PROCESSES

The blood, lymph and also all intercellular lymphs alive organisms is by aqueous solutions of moleculas and ions of many matters – organic and mineral. These solutions have definite osmotic pressure. So, the osmotic pressure of a blood of the person is value a constant and equally 7,4 105 – 7,8 105 Pa(pascal). Such high value osmotic pressure in a blood is conditioned by availability in her of a plenty of ions. High-molecular connection, mainly, proteins (albumines, the globulins), introduce 0,5 % common osmotic pressure of a blood. This part of osmotic pressure of a blood call oncotical as pressure, the value which one is equal 3,5-3,9 kPa. Oncotical pressure has large value for alive organisms. At a decrease oncotical of pressure the water goes in the party by high pressure – in a tissue, producing so called oncotical edemas of a hypodermic fat.

The osmotic pressure of a blood of the person is responded osmomolar concentration Dissoluble in plasma of matters, which one equal 0,287 – 0,0303 mol / liter.

Solutions with osmotic pressure, which is equal osmotic pressure of standard solution, is called isotonic. The solutions with osmotic by pressure are called as maximum for standard, hypertonic, and solutions with the lowest osmotic pressure hypotonic. In medical practice isotonic call solutions with osmotic pressure equal to osmotic pressure of a blood plasma. Such solution is 0, 89 % solution of sodium salt, and also 4,5 -5 % solution of a glucose. The isoosmotic solutions can be entered into an organism of the person in plenties. The hypertonic salt solutions enter in an person’s organism only in small amounts. At the introducing of a plenty hypertonic of solution the erythrocytes owing to loss of water decrease in volument and shrivel. Such phenomenon is called as a plasmolysis.

Importance of osmotic pressure of plasma proteins – The plasma proteins form а colloidal solution and are the chief colloid of the plasma. The osmotic pressure of plasma proteins which is called oncotic pressure though negligible (25 to 30 mm Hg) as compared to that of plasma crystalloids (about 5,000 mm Hg) is the main force which tends to keep the plasma water within the blood vessels. This is so because the concentration of non-colloids is almost the same in plasma and in the extracellular fluid and the osmotic pressure exerted by the non-colloids on the outside and the inside of capillaries are therefore balanced by each other. If the concentration of plasma proteins decreases markedly, water leaks into tissue spaces and the pathological condition called edema results.

Isosmotic, isotonic, hyposmotic, hypotonic, hyperosmot1c and hypertonic solutions – Isosmotic solutions are those which have the same osmotic pressure; 0.15 molar or about 0.90 % NaC1 solution in water is isosmotic with the human blood plasma. If human red blood cells are placed in this solution, they remain intact and retain their original shape and volume. 0.90% NaC1 solution is therefore isosmotic as well as isotonic with red blood cells because in this case0the amount of water entering the cells is equal to that leaving them; thus there is no net gain or loss of water by the cells.

If а solution of NaCl more concentrated than 0.90 % is used to suspend the human or other mammalian red blood cells, water will leave the cells and the cells shrink and become crenated. Such solutions are called hypertonic. On the other hand, if red blood cells are suspended in hypotonic solutions which have less than 0.90 % NaC1, then water will enter the cells making them swollen and if the solution is very dilute the cells will rupture releasing their hemoglobin in the solution. The rupture of red blood cells is called hemolysis. Frog’s blood plasma is isotonic with about 0.6% NaCl solution.

It should however be noted that osmosis and tonicity are different. А hyperosmotic solution is not necessarily а hypertonic solution. Osmolarity is а function of the number of solute particles in solution while tonicity is а function of how well а given solute causes osmosis across а cell membrane. Suppose we suspend the RBCs in two solutions of different solutes equally hyperosmolar as compared to the RBC interior, to one of which the RBC membrane is impermeable but is permeable to the other. Result will be that in the first case the outside medium of the BBCs will remain hyperosmotic and water will flow out of the RBCs. Thus this solution is both hyperosmotic as well as hypertonic. In the second case some solute particles will enter the cell interior from outside, raising its osmotic pressure and the cells will not lose water. This second type of solution is therefore hyperosmotic but not hypertonic. In the same way а solution may be hyposmotic but not hypotonic.

Cell contains а fluid (cell sap) and its wall is composed of а living cytoplasmic membrane which is semi-permeable and is responsible for the phenomenon of osmosis in living organisms. If such а cell comes in contact with water or some dilute solution, the osmotic pressure of which is less than that of cell sap present in the cell, there will be а tendency of water to enter into the cell through the cell wall. The pressure developed inside the cell due to the inflow of water into it is called turgor. On the other hand, if the cell comes in contact with а solution of higher osmotic pressure, the cell would shrink due to going out of water from the cell through the cell wall. This shrinking of the cell is called plasmolysis.

Fig. The effect of hypertonic and hypotonic solutions on animal cells.(а) Hypertonic solutions cause cells to shrink (crenation); (b) hypotonic solutions cause cell rupture; (c) isotonic solutions cause no changes in cell volume.

(а) (b) (с)

Osmotic pressure creates some critical problems for living organisms. Cells typically contain fairly high concentrations of solutes, that is, small organic molecules and ionic salts, as well as lower concentrations of macromolecules. If cells are placed in а solution that has an equal concentration of solute, there will be no net movement of water in either direction. Such solutions are called isotonic. For example, red blood cells are isotonic to а 0.9% NaCI solution. When cells are placed in а solution with а lower solute concentration (i.е., а hypotonic solution), water will move into the cells. Red blood cells, for example, will swell and rupture in а process called hemolysis when they are immersed in pure water. In hypertonic solutions, those with higher solute concentrations, cells shrivel because there is а net movement of water out of the cell. The shrinkage of red blood cells in hypertonic solution (е.g., а 3% NaCI solution) is referred to as crenation.

Because of their relatively low cellular concentration macromolecules have little direct effect on cellular osmolarity. However, macromolecules such as the proteins contain а large number of ionizable groups. The large number of ions of opposite charge that are attracted to these groups have а substantial effect on intracellular osmolarity. Unlike most ions proteins are unable to penetrate cell membranes. (Cell membranes are not, strictly speaking, osmotic membranes, since they allow the passage of various ions, nutrients, and waste products. The term dialyzing membrane gives а more accurate description of their function.) As а result, at equilibrium the concentrations for each ionic species will not be the same on both sides of а cell’s plasma membrane. Instead, the intracellular concentrations of inorganic ions will be higher than that found outside the cell. There are several consequences of this phenomenon, called the Donnan effect:

1. а constant tendency toward cellular swelling because of water entry due to osmotic pressure,

2. the establishment of an electrical gradient called а membrane potential.

Because of the Donnan effect, cells must constantly regulate their osmolarity. Living organisms use several strategies to accomplish this goal. Many cells, for example, animal and bacterial cells, pump out certain inorganic ions such as Na+. This process, which requires а substantial proportion of cellular energy, maintains cell volume within acceptable limits. Several species, such as some protozoa and algae, periodically expel water from special contractile vacuoles. Since plant cells have rigid cell walls, plants use the Donnan effect to create an internal hydrostatic pressure called turgor pressure. This process is the driving force in cellular growth and expansion. It is also responsible for the rigidity of many plant structures.

Fig. Osmotic Pressure and Plant Cells. (а) Isotonic solutions cause no changes in cell volume. (b) Plant cells typically exist in а hypotonic environment. Water enters these cells, and they become swollen. Cell bursting is prevented by the restraining force of rigid cell walls. (с) In а hypertonic environment the cell membrane pulls away from the cell wall because of water loss. This is the reason plants wilt when they receive insufficient water.

a) b) c)

Vapor-pressure lowering.

The escaping tendency of a liquid is measured by its vapor pressure, which is decreased by the presence of any solute. We saw that at any given temperature the vapor pressure of a pure liquid depends on the fraction of its molecules that have sufficient kinetic energy to escape from the attractions of their neighbors. Shows a pure liquid solvent in equilibrium with its vapor. Shows a solution at the same temperature, also in equilibrium with its vapor. (The solute is assumed to be nonvolatile, and so the only molecules in the gas phase are those of the solvent.) In this illustration, the concentration of molecules in the gas phase can be seen to be less. In the solution, not all of the molecules at the surface are solvent molecules, and so not all are potentially capable of escaping from the liquid. Thus, the rate of evaporation from the solution is less than from the pure solvent. As a result, the equilibrium concentration of molecules in the gas phase above the solution is less than that above the pure solvent. Thus, the vapor pressure of the solution is less than that of the pure solvent.

Shows two vapor-pressure curves, one for a pure solvent and one for its solution. The vertical distance between the two curves shows the magnitude of the vapor-pressure lowering at each temperature.

Raoult’s law. The relationship between vapor-pressure lowering and concentration in an ideal solution is stated in Raoult’s law.

Raoult’s Law: The partial vapor pressure of a component in liquid solution is proportional to the mole fraction of that component, the constant of proportionality being the vapor pressure of the pure component.

This means that a component’s vapor pressure is equal to the product of its mole fraction times its vapor pressure when pure.

Representing the solvent by the subscript 1, Raoult’s law can be written as

P₁ = X₁P₁⁰

where P₁ and P₁⁰ are the vapor pressure of the solution (actually, the vapor pressure of the solvent in the solution) and that of the pure solvent, respectively, and X₁ is the mole fraction of the solvent in the solution. Note that as long as the solute is not volatile, P₁ is the total vapor pressure of the solution. Now, since X₁ = 1 – X₂, where the subscript 2 represents the solute, P₁ = (1- X₂)P₁⁰

or, solving for X₂,

On the right-hand side of this relationship, P₁⁰ – P₁is the vapor-pressure

lowering brought about by the presence of the solute. is called the fractional vapor-pressure lowering, which can be seen to be equal to the mole fraction of the solute.

Boiling-point elevation.

The boiling point of a liquid is the temperature at which its vapor pressure equals atmospheric pressure. At any temperature, the presence of a solute lowers the vapor pressure of a liquid, and so in order to cause a solution to boil it is necessary to raise its temperature above the boiling point of the pure solvent.

Because it is a colligative property, vapor-pressure lowering in dilute solutions depends on the concentration of solute particles but not on their identity. Therefore, we anticipate a similar relationship between boiling-point elevation and solute concentration. It can be shown that in dilute solutions the boiling-point elevation is proportional to the molality of the solute particles. (As before, we assume that the solute is not volatile.) In other words, if DT_b, represents the boiling-point elevation, , then: DT_b = K_bm:

Where m is the molality of the solute and K_b, is a proportionality constant known as the molal boiling-point elevation constant.

Freezing-point depression.

The phenomenon of boiling-point elevation occurs because the presence of a solute lowers the escaping tendency of the solvent. Therefore, in order to cause a solution to boil, it is necessary to raise its temperature above the boiling point of the pure solvent. Escaping tendency means tendency for molecules to escape to any other phase, however. Consequently, in order to cause solvent molecules to freeze out of a solution, the solution must be cooled to a temperature lower than the freezing point of the pure solvent so as to compensate for the decreased escaping tendency of its molecules. The presence of a solute always lowers the freezing point of a solvent, as long as the solute is not also soluble in the solidified solvent. The depression of the freezing point causes the solid-liquid equilibrium line to be moved to the left in the phase diagram. Shows a composite representation of two phase diagrams: the first is for pure H₂O, and the second is for an aqueous solution of a solute that is not soluble in ice. Note that the solid-gas (sublimation) equilibrium line is unaffected. The diagram shows that the presence of the solute decreases the temperature at which ice and liquid can coexist at any pressure.

The relationship between freezing-point depression and molality in dilute solutions is a direct proportionality and is similar to that between boiling-point elevation and molality: DT_f= K_fm

Where: m – molality of solute; K_f – molal freezing-point depression constant

– freezing-point depression.

Biogenic elements. Qualitative reaction for ions of some macro- microelement. Complex compound in biological systems. Determination of water hardness.

Physical and chemical characteristics of water.

Water’s properties are directly related to its molecular structure.

Melting point of water – 0 ^oC; boiling point – 100 ^oC.

The state and distribution of water in the organism.

There are two water compartments in the body:

1. Intracellular water

2. Extracellular water

Extracellular fluid is divided into:

1. interstitial fluid

2. plasma

Distribution of water in an adult man, weighing 70 kg

Compartment

Body weight (%)

Volume (l)

Total

ICF

ECF

Interstitial fluid

10,5

Plasma

3,5

Biological role of water:

1. Water is an essential constituent of cell structures and provides the media in which the chemical reactions of the body take place and substances are transported.

2. It has a high specific heat for which, it can absorb or gives off heat without any appreciable change in temperature.

3. It has a very high latent heat. Thus, it provides a mechanism for the regulation of heat loss by sensible or insensible perspiration on the skin surface.

4. The fluidity of blood is because of water

5. Water is the most suitable solvent in human body

6. Dielectric constant : Oppositely charged particles can coexist in water. Therefore, it is a good ionizing medium. This increases the chemical reactions.

7. Lubricating action: Water acts as lubricant in the body to prevent friction in joints, pleura, conjunctiva and peritoneum

Water balance. Endogenous water.

Water balance is an equilibrium persists between the intake and output of water in the body. In addition to other factors, certain hormones, such as ADH, vasopressin, oxytocin and aldosterone influence the regulatory mechanism of body water.

A. Water intake : Water is supplied to the body by the following processes:

1. Water taken orally (1200-1300ml).

2. Along with food (1000ml).

3. Oxidation of foodstuffs : Fats, proteins and carbohydrates yield water after combustion. Fats produce 107 ml./lOO gm., proteins 41 ml./lOO gm. and carbohydrates 56 ml./100 gm. (metabolic water 300-400 ml).

B. Water losses : Water is lost from the body by 4 routes:

1. Evaporation from skin and lungs.

2. Kidneys, as urine.

3. The intestines, in the the feces.

4. Perspiration.

C. Additional water losses in disease:

1. Water loss is more in diarrhea and vomiting and these losses can be fatal in infants.

2. In kidney disease, renal water loss is more.

3. In fever, insensible losses may rise much higher than normal.

4. Patients in high environmental temperatures also sustain extremely high extrarenal water losses.

Regulation of Water Metabolism

There are several factors which regulate the water metabolism in the body; they are as follows:

1. Antidiuretic hormone or Vasopressin: Posterior pituitary releases ADH which has got the property to enhance water reabsorption in the distal tubules and collecting ducts. Water permeability gets increased.

2. Hypothalamus: There is a centre in hypothalamus known as a thirst centre;

whenever there is dehydration in the body, osmoconcentration of plasma takes place which eventually stimulates the thirst centre producing thirst as a result of which animal gets provoked to drink the required amount of water. Besides this, osmoconcentration of plasma also stimulates supraoptic and paraventricular nuclei of hypothalamus; nerve impulses from them are responsible to increase the release of vasopressin from neurohypophysis into the blood. Lesions in the supraopticoparaven-tricular region or in the neurohypophysis produce diabetes insipidus in which large volume of dilute (hypotonic) urine is passed, may be 10 litres or so a day. Diabetes insipidus is due to the abnormalities in the ADH secretion. In primary diabetes insipidus, there is less secretion of the ADH hormone which is usually due to the destruction of the hypothalamic-hypophyseal tract either from a basal skull fracture, or tumor, or infection. Besides, diabetes insipidus may be hereditary also.

3. Adrenal Cortex and Water loss: Loss of electrolytes and the loss of water from the body are closely interlinked. Adrenal cortex plays a very important role in governing the reabsorption of water by the renal tubules. The excretion of sodium and potassium by the kidneys is controlled by a steroid hormone called aldosterone which is secreted by the zona glomerulosa of the adrenal cortex. In man, aldosterone first increases the elimination of potassium and hydrogen ion and then decreases the excretion of sodium without any change in GFR. Aldosterone acts mainly at the distal tubule but its effect on sodium reabsorption may be partly at the proximal tubule. Apart from its action on the renal tubuie, aldosterone increases the reabsorption of sodium from the secretions of the intestinal mucosa and of the salivary and sweat glands. Thus, the body content of Na⁺ rises and that of K⁺ decreases.

Aldosterone is not the only cortical hormone affecting water balance. The diuretic response to a water load gets impaired in patients whose adrenal glands are destroyed by diseases like Addison’s disease or are removed after operation. The ability to deal normally with water is, however, restored by the administration of cortisone or hydrocortisone.

4. Rennin-Angiotensin system: This system is also involved in the regulation of blood pressure and electrolyte metabolism. The primary hormone involved is angiotensin II, an octapeptide formed from angiotensiongen. A decrease in circulating blood volume stimulates rennin secretion from kidneys which in turn promotes angiotensin formation in plasma. Angiotensin II stimulates the synthesis and secretion of aldosterone and the release of vasopressin, and thereby increases renal absorption of Na⁺and H₂O.

5. Prostaglandins: They are also believed to help maintain glomerular filtration inspite of hypotension, by causing renal vasodilation. They may also increase urinary loss of water by inhibiting the antidiuretic effect of vasopressin and by increasing the urinary sodium.

6. Solutes: Osmotic effect of Na⁺ helps to retain water in extracellular fluids. Elevation in plasma Na⁺ raises the ECF volume in primary aldosteronism while an increase in urinary Na⁺ raises the urinary water output in Addison’s disease. K⁺ helps to retain water in the cells, whereas, plasma proteins do help to retain water in the body by their osmotic effects. Increase in urinary urea or excretion of glucose in urine increases osmotically the urinary loss of water (osmotic diuresis).

Dehydration

Dehydration may be defined as a state in which loss of water exceeds that of intake, as a result of which body’s water content gets reduced. In this state, the body is iegative water balance.

Causes

1. Primary dehydration: There is purely water depletion and no salt depletion. It occurs in following states:

(a) Due to deprivation of water as generally happens during desert travelling.

(b) In mental patients who refuse to drink water/fluids.

(d) It occurs more quickly during fever or in the high temperature of the environment.

(e) Excessive water loss due to vomiting, prolonged diarrhoea, gastroenteritis.

(f) Due to excretion of large quantities of urine or sweat.

This type of dehydration raises the concentration and osmotic pressure of extracellular fluid as a result of which there is consequent outflow of the intracellular water to the ECF; thus, ECF volume gets largely restored but there becomes deficiency of water inside the cells as a result of which they suffer from osmoconcentration; symptoms of which include dry tongue, poor salivation, dry shrunken skin, nausea, reduced sweating and intense thirst.

When the blood becomes hypertonic, it lowers the urinary output and also makes the urine concentrated as a result of which there is less elimination of NPN and other acids which leads to acidosis and eventually coma. Death occurs in man due to renal failure, acidosis, intracellular hyperosmolality, circulatory collapse or neural depression, when body water falls by 20%.

Drinking of concentrated saline like sea-water or failure of Na⁺ excretion (e.g., in Cushing’s syndrome and Primary aldosteronism) may cause hypertonicity of ECF which in turn is responsible for withdrawal of water from tissue cells, dehydration of tissues, but a rise in ECF volume. Mg²⁺ of sea-water may be responsible for an increased intestinal loss of water due to its osmotic effect in the intestinal lumen.

2.Secondary dehydration: The concentration of the electrolytes of the body fluids is maintained constant either through The elimination or retention of water. The reduction or elevation in the total electrolytes, which affects the basic radicals chiefly i.e. Na (extracellular) and K (intracellular) and the acid radicals HCO₃ and Cl are accompanied by a corresponding increase or decrease in the volume of body water which is eventually the cause of intracellular edema; as a result of which there is slowing of circulation and impairment of urinal functions. All this causes an individual to become weak bodily.

3. Dehydration due to injection of hypertonic solution: When a highly concentrated solution of sugar or salt is injected into the body of an individual, the osmotic pressure of blood will increase which results in the flow of fluid from the tissues into the blood unless an equilibrium is reached. Consequently, the blood volume increases. This increased blood volume soon returns to normal by the loss of excess material through excretion which finally causes a net loss of body water producing dehydration.

Effects of dehydration

There are various side effects of dehydration as follows, which may be overcome as soon as the body gets hydrated; otherwise the consequences are serious and may even lead to death.

1. Disturbance in acid-base balance.

2. Loss of body weight due to the reduction in tissue water.

3. Rise in nonproteiitrogen (NPN) of blood.

4. Dryness, wrinkling and looseness of the skin.

5. Elevation in the plasma protein concentration and chloride.

6. Rise in the temperature of body due to reduction of circulating fluid.

7. Increased pulse rate and reduced cardiac output.

8. Exhaustion and collapse i.e. death.

Correction of dehydration

1. Ordinarily, sodium chloride solution may be given parenterally to compensate the loss.

2. In several disorders like diarrhoea, gastroenteritis, pancreatic or biliary fistulas, etc., a mixture of two-thirds isotonic saline solution and one-third sodium lactate solution (M/6) should be administered intravenously.

3. Dehydration is a burning problem in several disorders like diabetes mellitus, Addison’s disease, uremia, shock and extensive burns which is difficult to correct by the above two ways.

Water Intoxication

This condition is generally caused due to the retention of excess water in the body and can occur due to renal failure, excessive administration of fluids parenterally and hypersecretion of ADH. Symptoms of water intoxication include nausea, headache, muscular weakness, etc.

Biological role of potassium and sodium

Potassium ions promotes the protein synthesis by ribosomes;

number of enzymes require K⁺ for maximum activity (for example in the glycolitic sequence K⁺ is required for maximum activity of pyruvate);

metabolically supported gradients of Na⁺ and K⁺ across the cell membrane are involved in the maintenance of the membrane potential of excitable tissues, which is the vehicle for transmission of impulses in the form of an action potential;

K⁺ ions enhance the function of parasympathetic nervous system and acetylcholine action on the nervous terminals in muscles;

K⁺ ions reduce the exciting influence of ions on muscles;

a proper plasma K⁺ level is essential for the normal heart functioning more precisely for relaxation of miocardium (diastole);

sodium ions play the main role in regulation of osmotic pressure and retention of water in an organism;

sodium chloridum of blood plasma is the main origin of hydrochloric acid formation;

Na⁺ ions take part in the formation of a short-term memory.

Regulation of the Na and K metabolism in organism.

Sodium content in blood plasma is 130-150 mmol/l.

Potassium content in the blood is 3.4-5.3 mmol/l, this is only 2 % of all potassium content in the human body.

Kidneys are the main regulator of body Na⁺ and normally 98 % of the body loss of Na⁺ occurs in the urine. If more Na⁺ is ingested, its excretion in the urine increases. If less Na⁺ is ingested or if plasma Na⁺ falls due to any reason, Na⁺ may totally dissappear from the urine. This is brought about through the aldosteron which increases the tubular reabsorption of Na⁺ in the distal part of the nephrons.

The various factors can affect the urinary escretion of Na⁺:

1. When glomerular filtration is broken, the amount of Na⁺ filtered is smaal, the renal tubules reabsorb all the filtered Na⁺ resulting in Na⁺ retention and hypernatriemia.

2. When tubular reabsorption is broken (in chronic renal failure), it resulting in excessive urinary loss of Na⁺ and hyponatriemia.

3. Severe acidosis aggravates Na⁺ loss in the urine because the ranel tubules may fail to produce NH₃ in sufficient amount to buffer H ions in the tubular lumen; in this way, there arises a defficiency of NH₄ ions which could be excreted in the urine in exchange for Na⁺ ions.

4. Diuresis. Most Na⁺ is lost in diuretic conditions as diabetes mellitus, diabetes insipidus or after administration of mannitol or urea (osmotic diuretics).

5. Hormones. The main hormones, regulating Na⁺ metabolism, are mineralocorticoids and atrial natriuretic peptide (ANP). The mineralocorticoids, aldosterone and deoxycorticosterone, increase Na⁺ reabsorption from the tubular fluid and therefore their excess causes Na⁺ retention. In addition, these hormones increase the elimination of more K⁺ and Na⁺ in the urine. A greater formetion of aldosterone (primary aldosteronism or Conn’s syndrome) is associated with an increased Na⁺ retention in the body (hypernatriemia) with hypokaliemia and metabolic alkalosis. Conditions like congestive heart failure, cirrhosis of the liver and nephrotic syndrome also lead to a greater formation of aldosterone (secondary aldosteronism).

The atrial natriuretic peptide is produced y the atrial muscle fibers. It increases the urinary loss of Na⁺.

The severe decrease of Na⁺ in the extracellular space may lead to hypovolemia, hypotension, circulatory collapse and syncope.

Patients with Na⁺ excess show a raised venous pressure, peripheral and pulmonary edema with eventual respiratory failure. Cerebral symptoms may be seen due to hyperosmolality of the plasma.

The main causes of hyperkaliemia are:

1. Release of cellular K⁺from muscle tissue (hard traumas), in intravascular hemolysis, after extensive surgical operations.

2. Renal failure – the K⁺ secretion by the distal tubules is decreased and retention of K⁺ takes place.

3. Chronic dehydration and shock (associated with decreased formation of urine and K⁺ retention).

4. Acidosis – H⁺ ions displace K⁺ ions from the cells.

5. Addison’s disease.

Symptoms of hyperkaliemia are exerted mostly on the heart and nervous systems. When the serum K⁺ level is above 7 mmol/L, ECG changes are observed, bradycardia and arrhytmias appear. The heart becomes more susceptible to vagal stanstil, and heart may stop in diastole.

Hypokaliemia may be observed in decreased K⁺ intake (in starvation, malnutrition states such as kwashiorkor), in excessive renal loss (in metabolic alkalosis, using of some diuretics, such as furosemid, in renal tubular disorders, in hyperaldosteronism Iincreased production of aldosterone), in severe vomitting or diarrhea. Symptoms are: anorexia, nausea, muscle weakness and mental depression. Irregular pulse and a fall of blood pressure are observed.

Function of Na⁺, K⁺-ATP-ase

Most animal cells maintain intracellular K⁺ and extracellular Na⁺ at a relatively high concentration due to the operation of the special transmembrane enzyme which is called Na⁺, K⁺-ATP-ase. Na⁺, K⁺-ATP-ase use the energy derived from ATP to drive the transport of Na⁺ and K⁺ ions against the concentration gradient. The Na⁺, K⁺-pump is a prominent example of a primary transporter. Na⁺, K⁺-ATP-ase has the molecular weight of about 250000 to 300000 and contains two different types of subunits. The large subunit is the portion of the molecule that is phosphorylated as ATP is hydrolyzed. It has binding sites for Na⁺, K⁺and appears to extend through the entire thickness of the cell membrane. The smaller subunit is a glycoprotein and contains sialic acid as well as glucose, galactose and other hexose residues.

Biological role of Calcium and phosphorus

Calcium forms about 1% of adult body weight. It is the most abundant electrolyte in the human body due to its structural function for the skeleton. Normal serum or plasma calcium level is 2.3-2.75 mmol/l. More than 99% of calcium in the body occurs in bones as its phosphate and carbonate; only 0.03% of the total body calcium occurs in blood. The bone calcium is constantly exchanged with the calcium of interstitial fluid and this process is regulated primarily by the parathyroid hormone, active vitamin D and also by calcitonin.

Milk and milk products are the best dietary sources of calcium. Other good sources are egg yolk, leafy vegetables and hard drinking water. In spite of their high calcium content, some vegetable foods such as spinach contain also oxalates and benzoates and are a poor source of calcium because calcium oxalate and benzoate thus formed are insoluble and are not absorbed.

Functions of calcium in the body:

1. Calcium salts take part in bone and tooth development. Deficient supply of calcium leads to rickets in children and osteomalacia in adults. Sufficient calcium intake must be ensured in early life to build up the skeletal reserves. If this is not done, then there occurs an increased incidence of osteoporosis in old age because at that time deficiency of sex hormones especially in females results in calcium mobilization from bones leading to osteoporosis.

2. The clotting of blood needs calcium ions.

3. By regulating the membrane permeability calcium ions control the excitability of nerves. If plasma ionized calcium level falls markedly, tetany results in which spasms of various muscle groups occur. Death may occur from convulsions or from laryngospasm. An excess of plasma calcium depresses nervous activity.

4. Calcium ions act as a cofactor or activator of certain enzymes. A proteiamely calmodulin is present within cells, which can bind calcium. The calmodulin-calcium complex becomes attached to certain enzymes which ire activated. Such enzymes include adenylate cyclase, Ca²⁺ ATPase, phosphorylase kinase, myosin light chain kinase, phosphodiesterasc and phospholipase A; this mechanism also is required for the release of acetylcholine at the neuromuscular junctions.

5. Calcium ions take part in the contraction of muscle including heart muscle and are involved in the excitation-conraction coupling mechanism. In increased plasma calcium, heart stops in systole. In addition, a high plasma calcium decreases conduction of cardiac impulses and thus can produce heart block.

6. Calcium ions are responsible for initiating contraction in vascular and other smooth muscles. Calcium ions enter through specific channels just as is the case with cardiac muscle. Drugs that block these channels [Ca²⁺ channel blockers] have profound effect on the contractility of cardiac and smooth muscle as well as on the conduction of impulses within the heart. These drugs find use in the treatment of angina pectoris, cardiac arrhythmias and hypertension.

7. Calcium is essential for maintaining the integrity of capillary wall. In its deficiency, capillary walls become fragile and there is increased permeability of capillaries.

8. Calcium ions are involved in exocytosis and thus have an importrole in stimulus-secretion coupling in most exocrine and endocrine glands, e.g. the release of catecholamines from the adrenal medulla, neurotransmitters at synapses and histamine from mast cells is dependent upon Ca²⁺.

9. Some hormones exert their influence through Ca²⁺. For example, the effect of adrenaline on the liver cells to increase glycogenolysis is partly due to an increased Ca²⁺ within these cells which is independent of cAMP.

Biological role of phosphorus

An adult body contains 1 kg phosphate and it is found in every cell of the body. Most of it (about 80%) occurs in combination with calcium in the bones and teeth. About 10% of body phosphorus is found in muscles and blood in association with proteins, carbohydrates and lipids.

Biochemical functions:

1. Phosphorus is essential for the development of bones and teeth.

2. It plays a central role for the formation and utilization of high-energy phosphate compounds (ATP, GTP, creatine phosphate etc.).

3. Phosphorus is required for the formation of phospholipids, phosphoproteins and nucleic acids (DNA and RNA).

4. It is essential component of several nucleotide coenzymes eg. NAD, NADP, pyridoxal phosphate, ADP, AMP.

5. Several proteins and enzymes are activated by phosphorylation.

6. Phosphate buffer system is important for the maintenance of pH in the blood as well as in the cells.

Role of vitamins and hormones in regulation of phosphorous metabolism

The hormones – calcitriol, parathyroid hormone and calcitonin are the major factors that regulate the plasma phosphorus within a narrow range (1.2-2.2mmol/l). Calcitriol is the biologically active form of vit.D. It acts at 3 different levels (intestine, kidneys and bone). Calcitriol increases the intestinal absorption of calcium and phosphate. Calcitriol along with parathyroid hormone increases the mobilization of calcium and phosphorus from bone.

Calcitriol is also involved in minimizing the excretion of Ca and P through the kidney, by decreasing their excretion and enhancing reabsorption. Calcitonin inhibits the reabsorption of phosphorus in kidneys. Thus, calcitonin decreases the phosphorus content in blood. Parathyroid hormone decreases serum phosphorus and increases urinary PO₄(increase phosphorus excretion in urine).

Calcium

Calcium is present in the body in the largest amount of all the minerals present in the body. Calcium comprises 2 percent of the body weight. RBC is devoid of calcium. The normal serum level is 9 – 11 mg percent.

Calcium is present in three forms:

1. Ionized form.

This form is physiologically active form.

2. Protein bound fraction.

On the far right we see some lamellar bone, looking like strata of geologic layers of earth. Into such bone, cones of cutting blood vessels penetrate. They then organize bone around their path of penetration forming bone in long layered tubes. We see those in cross section on the left. They are called Haversian Tubules (take a wild guess who they are named after).

Cells are seen in layers around a central canal (Haversian Canal). Radiating micro tubules reach out intercommunicating.

Cells within bone (osteo) can make (“blast”) or degrade and remove (“clast”) bone substance. Osteoblasts lay down bone and osteoclasts digest bone. How active is all this? Well, depending on who you are, roughly one third of the bone you had yesterday isn’t the bone you have today.

The layering and bundles of layered tubes are of cells and sheets of tissue made of very tough fiber called bone collagen. A bone with all of its calcium leached out can be tied in a knot. On and through that structure, a very special crystalline form of calcium is formed called hydroxyappetite. That stiffens the bone to make it hard.

Bone gets strength from two things.

1) What it is made from (kind of bone), and

2) How it is shaped and organized.

Hydroxyappetite figures in both aspects of strength. The first is kind of obvious. Some stuff is tougher than other stuff. But shape is very important. Tubes are stronger than rods. Tension and compression struts vastly support structures (look at power towers). As with certain crystals – such as those used in phonographs (remember those) – when pressure is placed on them, they polarize and exhibit an electric charge. A needle jiggling in a plastic track while pressing against a crystal will reflect the jiggling as a fluctuating charge which magnified gives us music – and we listen.

The compression of forces of daily activity on hydroxyappetite gives us zones of charge to which the osteoblasts listen. They respond by putting more bone substance where forces generate such charges. Where such charges fail to form, bone – always being dissolved – wheedles away. The form of bone follows function. In other words, as was spoken by Hypocrites a few years ago, “That which is not used, wastes away.” The paraphrase is, “Use it or lose it.”

This had to be rediscovered when perfected devices which held fractured bone pieces absolutely rigidly, better than ever before … produced poorer healing. Without SOME movement, bone formation is not very good.

Calcium

Calcium is an important substance beyond mere bone structure. In ionic form (the molecular water dissolved form) it is a charged atom as Ca⁺⁺ which means it has charge that pairs it with negative charged things such as two OH^– (hydroxyl) molecules or phosphate as PO₄⁼. In this charged form it is used to regulate, control or initiate processes such as nerve conduction, muscle contraction, hormone release among other things. It is very tightly kept at optimum in the blood – even if that means taking calcium from bones. Calcium is the stuff that leaves water spots on dish ware or which makes bathtub rings when combined with certain substances. That is, it easily precipitates. X-rays of injured or inflamed tissue may show deposits of calcium. That does not mean calcium CAUSED the problem, but more likely that calcium is precipitating due to the problem.

Indeed, the combined levels of Ca⁺⁺and PO₄⁼found normally in the blood exceed levels which can be achieved in water without precipitating. Blood is hyper saturated with calcium by means of other substances which stabilize calcium in solution. An interesting complication of this fact to babies – especially preemies – is that IV fluids used to maintain babies (who cannot eat) cannot contain EITHER enough Ca⁺⁺ OR enough PO₄⁼ to sustaieeds. Children can get a baby form of rickets which may look like bone loss with fractures. To get around this, if an IV is needed long term, solutions with high calcium have to be alternated with solutions of high phosphate. If the correct amounts of BOTH were placed in a single solution, the calcium phosphate would solidify in the bottle.

When calcium levels in the blood are off by substantial amounts then something else must be at the core of the matter: low protein, high protein, kidney disorder etc. The people who manage calcium and those other difficult blood salts and kidney function are the ones with tall foreheads, out of control hair and who horde all the back issues of the Journal of Clinical Investigation

Diphosphonates are drugs which deliver dual phosphate molecules. Pairs of phosphates glom onto hydroxyappetite like egg dye on an egg. That tends to “stabilize” the otherwise very rapid turnover of the crystal. When there is a process that draws away calcium, the diphosphonates protect the crystals. But they can get in the way of build up when the process is in the direction of accumulation. Newer more clever chemistries are being used to see if the plusses can be made to outweigh the minuses. The very uneven results seen with these drugs as a class reflects the inconsistency of the application. There is tremendous variation and an experienced professional is needed to know if it is helping or hurting.

Body Distribution of Calcium and Phosphate

There are three major pools of calcium in the body:

Intracellular calcium: A large majority of calcium within cells is sequestered in mitochondria and endoplasmic reticulum. Intracellular free calcium concentrations fluctuate greatly, from roughly 100 nM to greater than 1 uM, due to release from cellular stores or influx from extracellular fluid. These fluctuations are integral to calcium’s role in intracellular signaling, enzyme activation and muscle contractions.
Calcium in blood and extracellular fluid: Roughly half of the calcium in blood is bound to proteins. The concentration of ionized calcium in this compartment is normally almost invariant at approximately 1 mM, or 10,000 times the basal concentration of free calcium within cells. Also, the concentration of phosphorus in blood is essentially identical to that of calcium.
Bone calcium: A vast majority of body calcium is in bone. Within bone, 99% of the calcium is tied up in the mineral phase, but the remaining 1% is in a pool that can rapidly exchange with extracellular calcium.

As with calcium, the majority of body phosphate (approximately 85%) is present in the mineral phase of bone. The remainder of body phosphate is present in a variety of inorganic and organic compounds distributed within both intracellular and extracellular compartments. Normal blood concentrations of phosphate are very similar to calcium.

Fluxes of Calcium and Phosphate

Maintaining constant concentrations of calcium in blood requires frequent adjustments, which can be described as fluxes of calcium between blood and other body compartments. Three organs participate in supplying calcium to blood and removing it from blood wheecessary:

The small intestine is the site where dietary calcium is absorbed. Importantly, efficient absorption of calcium in the small intestine is dependent on expression of a calcium-binding protein in epithelial cells.
Bone serves as a vast reservoir of calcium. Stimulating net resorption of bone mineral releases calcium and phosphate into blood, and suppressing this effect allows calcium to be deposited in bone.

This form is physiologically inert.

3. In combination with citrates.

Protein bound fraction is non-diffusible whereas other two fractions are diffusible.

Biological role of magnesium.

Half of magnesium occurs in the inorganic matter of bones and the rest occurs in soft tissues and body fluids. Blood plasma contains 0.8-1.2 mmol/l of Mg.

Nuts, legumes, chlorophyll and whole grains are very good sources of magnesium.

Functions of Mg in the body:

1. It takes part in the formation of complex salts of bones and teeth.

2. It acts as a cofactor for many enzymes.

3. It serves to decrease neuromuscular irritability.

Effects of a high serum Mg²level – Experimentally, a serum Mg²⁺ level of 8 mmol/L produces immediate and profound anesthesia and paralysis of voluntary muscles. These effects can be reversed by an intravenous injection of a corresponding amount of Ca²⁺. Serum Mg²⁺ tends to rise in renal failure.

Deficiency of Mg may occur in the malabsorptive syndrome, increased renal losses (diuretics, gentamycin intake and primary renal disease), chronic alcoholism, diabetic acidosis, cirrhosis of the liver, primary aldosteronism, hyperparathyroidism, prolonged and severe losses of body fluid and prolonged administration of Mg-free intravenous fluids. Plasma Mg may be lowered after parathyroidectomy (along with hypocalcemia) due to avidity of bones for divalent ions. In acute pancreatitis, Mg may become bound as soaps thus decreasing its plasma level.

Symptoms and signs of hypomagnesemia.

1. Neuromuscular disorders – weakness, tremors, muscle fasciculations and sometimes tetany.

2. Central nervous system disorders – personality changes, delirium, psychosis and coma.

Biological role of iron.

Iron is part of the structure of many important body constituents, e.g. hemoglobin, myoglobin, enzymes like cytochromes, catalase, xanthine oxidase, mitochondrial α-glycerophosphate oxidase, etc. The iron content of hemoglobin is 0.34%.

Dietary sources – Animal sources are the best and include liver, red meat and egg yolk. Of the vegetables, spinach and other leafy vegetables are good sources. Dried fruits also contain appreciable amounts of iron.

Plasma iron transport – Before iron can leave the intestinal mucosal cells, it is first converted to Fe²⁺ form. On entering the plasma, it is again oxidized to Fe³ form and is taken up by a pink colored protein called siderophilin or transferrin, having a mol. wt. about 80,000. The transfer of iron to the transferrin is catalyzed by a Cu-containing protein, namely ceruloplasmin. One molecule of siderophilin binds 2 atoms of iron. The absorbed iron is utilized to form products such as heme, etc. and the remaining portion is mostly stored in the body as ferritin in the reticuloendothelial cells and hepatocytes. Ferritin is a conjugated protein; its iron is in combination with the protein part of the molecule called apoferritin. Apoferritin has a mol. wt. of about 450,000 and is composed of 24 polypeptide subunits; these form an outer shell within which resides a storage cavity for polynuclear hydrous ferric oxide phosphate. Iron within ferritin molecule occurs as ferric hydroxide-ferric phosphate complex and its iron content may be upto 30%. Ferritin also is present in blood plasma where its level is a good index of iron stores of the body. If iron is in excess, then ferritin molecules aggregate forming hemosiderin (iron content upto 55%). Hemosiderin is stored as microscopically visible golden brown granules because it is insoluble due to its containing a characteristic arrangement of the micelles of Fe(OH)₃.

Recently the genes for the human transferrin receptors and ferritin have been discovered and the mechanism of regulation of expression of transfcrrin receptors and intraccllular ferritin in response to the iron supply have been established. When iron is in excess, the synthesis of transferrin receptors is decreased and ferritin production is increased; this favors iron storage. When iron is deficient, then reverse changes occur which lead to a decreased ferritin production that decreases iron storage so that iron can be utilized in the body to a maximum.

Factors increasing iron absorption from the intestine:

1. Conditions associated with increased rate of erythropoiesis.

2. Low body stores of iron

3. Taking ascorbic acid, succinic add, fructose and sorbitol along with iron – Ascorbic acid favors reduction of Fe³⁺ to Fe²⁺; the latter is more readily absorbed. The other compounds make complexes with iron and increase its absorption.

4. Intake of inorganic iron.

Factors inhibiting iron absorption:

1. Malabsorption syndromes.

2. Diarrheal diseases.

3. An excess of phosphates, oxalates and phyfic acid – These form complexes with iron which are insoluble and cannot be absorbed. Vegetable foods have an excess of phosphates and interfere with iron absorption.

4. Subtotal gastrectomy.

5. Surgical removal of the upper small intestine.

6. Food intake along with iron.

7. Antacid therapy.

8. Chronic Infections.

The iron deficiency results in anemia – this is of hypochromic, microcytic type. It is the most common type of anemia being specially present in women of child-bearing age and infants below 1 year of age. The RBCs are smaller in size and have less hemoglobin as well as less mean corpuscular volume RBC count is low but hemoglobin level of blood is proportionately still lower. The color index which is hemoglobin as °o of the normal divided by RBC count as % of the normal is therefore below one. In addition to the symptoms which are common to all anemias such as a pale appearance and breath lessness on minor exertion, the patient shows some characteristic features. These include a derangement of epithelial surface such as abnormal nail growth (spoon shaped nails or koilonychia), glossitis, fissures around the corners of the mouth and localized thickening of the mucous lining of the esophagus causing dysphagia (Plummer-Vinson syndrome)

Biological role of iodine , fluoride, copper, zinc, selenium and cobalt.

Iodine

The total body contains about 20mg iodine, most of it (80%) being present in the thyroid gland. The only known function of iodine is its requirement for the synthesis of thyroid hormone mainly thyroxin (T4) and triiodothyronin (T3).

Dietary requirements: 100-150micrograms per day.

Sources: Sea food, drinking water, iodized salt.

Diseases states: Toxic goiter.

Fluoride

Functions:

1.It prevents the development of dental caries.

2.It is necessary for the proper development of bones .

3.It inhibits the activities of certain enzymes.

Dietary requirements: 1-2 mg per day.

Sources: Drinking water.

Diseases states: dental caries, fluorosis.

Copper

Functions:

1. Its an essential constituent of several enzymes (cytochrome oxidase, catalase, superoxide dismutase etc.)

2. Its necessary for the synthesis of hemoglobin, melanin and phospholipids.

3. Ceruplasmin has oxidase activity and thereby facilitates the incorporation of ferric iron into transferrin.

4. Development of bone and nervous system (myelin requires Cu).

Dietary requirements:2-3 mg per day.

Sources: Liver, kidney, meat, egg yolk, nuts and green leafy vegetables.

Disease status:

1. Copper deficiency (anaemia).

2. Menke’s disease (defect in the intestinal absorption of copper).

3. Wilson’s disease

Zinc

Functions:

1. It is an essential component of several enzymes (carbonic anhydrase, alcohol dehydrase etc.)

2. The storage and secretion of insulin from the beta – cells of pancreas requires zinc.

3. It is require for wound healing.

4. It is essential for the proper reproduction.

Dietary requirements: 10-15g per day.

Sources: Meat , fish, eggs, milk, nutts.

Disease status:

Zinc deficiency: poor wound healing, anaemia, loss of appetite, loss of taste sensation.

Cobalt

Cobalt is only important as constituent of vit-B12. The functions of cobalt is same as that of vit B12.

Selenium

Functions:

1. Selenium along with vit E, prevents the development of hepatic necrosis and muscular dystrophy.

2. Selenium is involved in maintaining structure integrity of biological membranes.

3. Selenium prevents lipid peroxidation and protect the cells against the free radicals.

4. Selenium binds with certain heavy metals and protects the body from their toxic effects.

Dietary requirements: 60-250 micrograms.

Sources: Liver, kidneys, seafood.

Toxicity: Selenosis is toxicity due to very excessive intake of selenium. The manifestation of selenosis includes weight loss, emotional disturbances, diarrhea, hairloss and garlic odour in breath.

Chlorine is contained in all biological liquids of the organism.

Functions:

1. As a component of sodium chloride, it is essential in acid-base equilibrium:

2. As chloride ion, it is also essential in water balance and osmotic pressure regulation.

3. It is also important in the production of hydrochloric acid in the gastric juice.

4. Chloride ion is important as an activator of amylase

Sources: It is mainly available as sodium chloride

Daily requirement: 5 – 10 g

The requirements of NaCI depend on the climate and occupation and on the salt content of the diet. Foods of animal origin contain more NaCI than those of vegetable origin.

Disease state: chloride deficit also occurs when losses of sodium are excessive in diarrhea, sweating; loss of gastric juice by vomiting.

Excretion: it is chiefly eliminated in the urine. Also Cl is excreted in the sweat.

Sulphur

In the organism sulphur exists both as organic and inorganic compounds

Functions:

1. It is present primarily in the cell protein in the form of cysteine and methionine.

2. The cysteine is important in protein structure and in enzymic activity.

3. Methionine is the principal methyl group donor in the body .

4. Sulfur is a constituent of coenzyme A and lipoic acid which are utilized for the synthesis of acetyl-CoA and S-acetyl lipoate, respectively.

5. Sulfur is a component of other organic compounds, such as heparin, glutathione, thiamine, biotin, taurocholic acid, sulfocyanides, indoxyl sulfate, chondroitin sulfate, insulin, penicillin, anterior pituitary hormones and melanin.

Sources:

Sulfur intake is mainly in the form of cystine and methionine present in proteins. Other compounds present in the diet contribute small amounts of sulfur.

Disease state:

1. The serum sulfate concentration is increased in the presence of renal functional impairment, pyloric and intestinal obstruction and leukemia.

2. Marked sulfate retention in advanced glomerulonephritis cause the development of acidosis.

3. An increase in the blood indican concentration (indoxyl potassium sulfate) may occur in uremia.

Excretion: it is excreted in the urine

Complex compound

Coordination compounds are the compounds in which the central metal atom is linked to а number of ions or neutral molecules by coordinate bonds i.е. by donation of lone pairs of electrons by these ions or neutral molecules to the central metal atom е.g. nickel tetracarbonyl, [Ni(CO)₄] in which CO molecules are linked to the central nickel atom by coordinate bonds by donating lone pairs of electrons.

If the species formed by linking of а number of ions or molecules by co-ordinate bonds to the central metal atom (or ion) carries positive or negative charge, it is called a complex ion, е.g. [Fe(СN)₆]^4-, [Cu(NH₃)₄]²⁺, [Ag(CN)₂]^– etc. Hence co-ordination compounds may also be defined as those compounds which contain complexions е.g.

K₄[Fe(СN)₆], [Cu(NН₃)₄]SO₄, Na[Ag(CN)₂] etc.

The branch of inorganic chemistry dealing with the study of co-ordination compounds is known as co-ordination chemistry.

Types of complex compounds There are following three types of complexes:

(i) А complex in which the complex ion carries а net positive charge is called cationic complex, е.g. [Co(NН₃)]³⁺, [Ni(NH₃)₆]²⁺ etc.

(ii) А complex in which the complex ion carries а net negative charge is called anionic complex, е.g.

[Ag(CN)₂]^– [Fe (CN)₆]^4-

(iii) А complex carrying no net charge is called а neutral complex or simply а complex, е.g. [Ni(CO)₄], [CoC1₃ (NН₃)₃] etc.

One central atom:

Ammonia complex [Cu(NH₃)₄]SO₄

Aqua complex [Al(H₂O)₆]Cl₃

Acidic complex K₂[PtCl₄]

Complex with difference ligands K[Pt(NH₃)Cl₃]

Cyclic (chelates)

Hem Chlorophyll

Polycentral compoynds

Chain [Cr(NH₃)₅ – OH – (NH₃)Cr]Cl₃

Chelaes (CO)₅Mn – Mn(Co)₅

Before we take up а study of the different aspects of co-ordination compounds, it is important to know some terms to be used therein. А few of these are briefly described below:

(1) Ligands. In the formation of the coordinate bonds, the anions or the neutral molecules act as the electron-pair donors whereas the central metal ion acts as the electron pair acceptor.

The donor atoms, molecules or anions, which donate а pair of electrons to the metal atom and form и co-ordinate bond with it are called ligands.

The common donor atoms in ligands are nitrogen, oxygen and less common are arsenic and phosphorus. For example, in [Ni(NН₃)₃]²⁺, central ion is Ni²⁺ and ligands are NH₃ molecules.

The ligand may contain one or more than one donor atom. If only one donor atom is present in its molecule, which can coordinate, then it is called as unidentate (unidentate means one point of attachment or having “one tooth”, uni = one and dent = tooth). This is also referred to as monodentate. А few examples are: NH₃, Н₂О and CN^–. They mayor may not be ionic but the complex part always contains co-ordinate bonds.

The ligand may contain two donor atoms (i.е. coordinating groups) positioned in such а way that а five or а six membered ring is formed with the metal ion, then it is called bidentate chelating ligand and the ring is called chelate ring and the resulting complex is called а metal chelate. The well known examples of the bidentate ligands are

The complexes formed by Cu (II) and Pt (II) ions with ethylenediamine are metal chelates represented as follows:

The ethylene diamine (еn) has two nitrogen atoms and oxalate ion has two oxygen atoms, which, can link to the metal ion.

Similarly, we may have tridentate, tetradentate, hexadentate and polydentate ligands. The hexadentate ligand, edta, (EDTA) has six donor atoms i.е. two nitrogens and four oxygens (of the carboxylic acid groups) capable of bonding to the metal atom.

Some important characteristics of chelates.

These are as follows:

(i) Chelating ligands form more stable complexes than the monodentate analogs. This is called chelating effect.

(ii) Chelating ligands, which do not contain double bonds e.g. ethylenediamine form five membered stable rings. The chelating ligands such as acetylacetone form six membered stable ring complexes.

(iii) Ligands with large groups form unstable rings than the ligands with smaller groups due to steric hindrance.

Importance of chelates. Chelates are widely used in industry and laboratory

(i) in the softening of hard water

(ii) in the separation of lanthanides and actinides

(iii) in the detection of some metals in qualitative analysis

(iv) in the estimation of nickel (II), magnesium (II) and copper (II) ions quantitatively.

(2) Coordination number. The total number of monodentate ligands (plus double the number of bi dentate ligands if any) attached to the central metal ion through coordinate bonds is called the coordinatioumber of the metal ion. In other words, coordinatioumber may be defined as the number of co-ordinate bonds formed with the central metal ion by the ligands.

For example, in the complex ions: [Ag(СN)₂]^–, [Cu(NН₃)₄]²⁺ and [Cr(Н₂О)₆]³⁺, the coordinatioumbers of Ag, Cu and Cr are 2,4 and 6 respectively. Similarly, in the complex ion, [Fe(C₂O₄)₃]^2-, the co-ordinatioumber of Fe is 6 because C₂O₄^2- is а bidentate ligand. The more common coordinatioumbers for metals are 2, 4 and 6 while less common are 3, 5, 7 and 8.

(3) Coordination sphere. The central atom and the ligands which are directly attached to it are enclosed in square brackets and are collectively termed as the coordination sphere. The ligands and the metal atom inside the square bracket behave as а single constituent unit. The ionizable groups are written outside the brackets. For example, in the coordination compounds, [Cu(NH₃)₄]SO₄, the complex ion [Cu(NН₃)₄]²⁺, in which Cu²⁺ is the central metal ion and four NH₃ molecules are ligands forms the coordination sphere. Similarly, in К₂[Pt Cl₆], the complex ion [Pt Cl₆]^2- in which Рt⁴⁺ is the central metal ion and six Сl^– ions are the ligands form the coordination sphere of this complex.

(4) Oxidatioumber or oxidation state. It is а number that represents an electric charge which an atom or ion actually has or appears to have when combined with other atoms е.g., oxidatioumber of copper in [Cu(NH₃)₄]²⁺ is +2 but coordinatioumber is 4.

Similarly, the oxidation number of Fe in [Fe(СN)₆]^3- is + 3 but the coordinatioumber is 6.

The method of calculation of oxidatioumber of а metal in а coordination compound or а complex ion is illustrated below with а few examples:

(i) Oxidatioumber of Cu in [Cu (NН_З)₄]SO₄. Sulphate ion (SO₄ ^2–) carries charge = – 2. As the complex is neutral, charge on the complex ion should Fе = + 2. As NH₃ carries no charge, therefore, charge on copper = + 2 i.е. oxidatioumber of Cu = +2.

(ii) Oxidatioumber of Fe in [Fe (СN)₆]^3-

As each CN ion carries charge = – 1, charge on 6 CN ions = – 6. As total charge on the complex ion = – 4, therefore charge on Fe must be = + 2 i.e. oxidatioumber of Fe = + 2.

(iii) Oxidatioumber of Fe in К₃[Fe(С₂О₄)₃]. As each К⁺ ion carries charge = + 1, charge оn 3 К⁺ ions = + 3. Hence charge on the complex ion = – 3. As each oxalate ion

C₂O₄^2- has charge = – 2, charge on the three С₂О₄ = – 6. Therefore, charge on Fe should be + 3. oxidation number of Fe in the given complex = + 3.

(iv) Oxidatioumber of Ni in [Ni(CO)₄].

Total charge on the complex = 0. As CO carries no charge, charge on Ni should be = 0, oxidatioumber of Ni in the given complex =0.

(5) Charge on the complex ion. The charge carried by а complex ion is the algebraic sum of the charges carried by central metal ion and the ligands coordinated to the central metal ion. For example, in the complex ion, [Ag (СN)₂]-, Ag⁺ ion carries а charge of + 1 and each CN^– ion carries а charge of –1. Therefore, the net charge on the complexion [Ag(СN)₂]^– is +1 +(– 2)= – 1.

Similarly, in the complex ion [Cu (NH₃)₄]²⁺, Cu²⁺ ion carries а charge equal to + 2 and as NH₃ molecules are neutral, therefore, the net charge on the complex is + 2.

Various terms discussed above may be illustrated by taking an example of the formation of а complex of CoC1₃ with NН₃ as shown below:

Nomenclature: Coordination compounds are formulated and named according to the system set up by Inorganic Nomenclature Committee of the International Union of pure and Applied Chemistry (IUPAC). According to the latest (1990) IUPAC system, the following rules are observed while writing formulas and naming coordination compounds.

Rules for Formula Writing:

(1) Formula of the cation whether simple or complex is written first followed by that of the аinon.

(2) The coordination sphere is written in square brackets.

(3) The sequence of symbols within the coordination sphere is first the metal atom followed by anionic ligands, theeutral ligands and finally cationic ligands.

[Metal atom, anionic, neutral, cationic ligands]

If there are а number of anionic ligands, they are listed alphabetically according to the first symbol of their formulae. Same principle is followed for neutral ligands or positive ligands.

The formulae of а few complexes are given below:

Na[PtBrCl (NO₂)(NН₃)]

[Co(H₂O)₂(NН₃)₄] Сl

(4) Polyatomic ligands are enclosed in parentheses but all ligands are formulated without any space in between.

(5) The number of cations or anions to be written in the formula is calculated on the basis that total positive charge must be equal to the total negative charge, as the complex as а whole is electrically neutral.

2. Rules for Naming the Coordination Compounds:

(1) Order of naming ions: The positive ion (cation) whether simple or complex, is named first followed by the negative ion (anion). The name is started with а small letter and the complex part is written as one word, е.g.

[Co(NН₃)₆] C1₃, hexaamminecobalt (III) chloride.

K₂[Pt C1₆], potassium hexachloroplatinate (IV).

But the non-ionic and molecular complexes are given one word name

[Co(NO₂)(NH₃)₃], triamminetrinitrocobalt (III)

[PtC1₄(NH₃)₂], diamminetetrachloroplatinum (IV).

(2) Naming of ligands: Different types of ligands are named differently as follows:

(i) Negative ligands end in – 1, е.g., СN^– (cyano), Сl^– (chloro), Br^– (bromo), F^– (fluorо), NO₂^– (nitro), ОН^– (hydroxo), О^2- (охо), SO₄^2-(sulphato), С₂О₂^2- (oxalato), NН₂^– (amido ), NH^2- (imido), ONO^– (nitrito), NO₃^– (nitrato), SCN^– (thiocyanato), NCS^– (isothiocyanato), СН₂(NН₂)COО^– (glycinato)etc.

(ii) Neutral ligands have no special ending: NН₃ (ammine), Н₂О (aqua), CO (carbonyl), CS (thiocarbonyl) and NO (nitrosyl)

(iii) Positive ligands (which are very few) end in -ium, е.g., NН₃⁺ (hydrazinium), NO⁺ (nitrosonium), NО₂⁺ (nitronium).

(iv) Organic ligunds. Organic free radicals are given their owames. For example, СН₃ (methyl), С₂Н₅ (ethyl), С₆Н₅ (phenyl), С₅Н₅ (cyclopentadienyl).

For organic neutral molecules, their names are used. For example, Р(С6Н₅)₃, triphenylphosphine;

(v) Unidentate ligands with more than one co-ordinating atoms. It is essential to designate the point of attachment of а ligand by placing the symbol of the donor atom attached after that; name of the groups separated by hyphen. These ligands are called ambidentate ligands е.g., in thiocyanate and nitrite ions, we have two options each.

— SCN, thiocyanate – NО₂^– nitro

— NCS, isothiocyanate – ONO, nitrite