Water PDF - Chemical Compound in Living Systems

Water:  The most abundant chemical compound in living system 60-70% of our body is water.  2/3 of the total body water is intracellular.  The rest is interstitial fluid of 25% is in the blood plasma.  Acts as a solvent for the substances we need such as Na+ ,K+ , gluco...

Water:  The most abundant chemical compound in living system 60-70% of our body is water.  2/3 of the total body water is intracellular.  The rest is interstitial fluid of 25% is in the blood plasma.  Acts as a solvent for the substances we need such as Na+ ,K+ , glucose , adinosine tri phosphate(ATP) and proteins.  The body controls both the volume and the PH of water.  Maintain a constant environment for the cells called homeostasis ( the same environment ).  Regulates the body temperature.  A reactant and product in many chemical reactions that take place in the living system.  Transporting load and oxygen to the cell and carrying away wastes.  A medium for movement of molecules into and out of cellular compartments.  Important to the structure and function of bio molecules. Buffers: A buffer is one which resists change in pH when small quantities of an acid or a base are added to it. 1. Acidic buffer : is one which has a pH less than 7. Acidic buffer solutions are made from a weak acid and one of its salts-often a sodium salt. Example : a mixture of ethnoic acid and sodium ethanoate. 2. Basic buffer : is one which has a pH greater than 7. Alkaline buffer solutions are made from a weak base and one of its salts. Example : a mixture of ammonia solution and ammonium chloride. How do buffer solutions work?  Acid buffer : example: Ethnoic acid & sodium ethanoate. Ethnoic acid is a weak acid and the position of equilibrium will be to the left. CH3COOHaq ↔ CHECOO-aq +H+ Adding CH3COONa adds more CH3COO- ions. According to the chatelaines principle ( if a system at equilibrium is disturbed by changing the conditions , the position of equilibrium will be shifted to counteract the change) , the position of equilibrium will be shifted further to the left. Adding an acid ti this buffer , the buffer removes most of the new hydrogen ions by combination of H+ ions with CH3COO- ions to form CH3COOH. Since CH3COOH is a weak acid most of new H+ ions removed by this way. CH3COO-aq + H+aq ↔ CH3COOHaq since most of the new H+ ions are removed , the PH will not change very much. Adding an alkali to this buffer , the OH- will react with CH3COOH and position of equilibrium will be shifted to right. CH3COOHaq + OH-aq↔ CH3COO-aq + H2O  Alkaline buffer solutions : consist of weak base and its conjugate acid. Example : NH+4 & NH4CL. NH3aq+H2O ↔ NH4+aq + OH- Adding NH4CL to this adds more NH4+ ions. According to Le chatelier,s principle , the equilibrium will be shifted to the left. Adding an acid , the H + ions of the acid are removed by reacting with ammonia and the equilibrium will be shifted to right. Also the H+ ions are removed by reacting with OH- form from the reaction of NH3 with H2O. NH3 + H2O ↔ NH+4 + OH- Adding alkali to this buffer, the OH- of the alkali are removed by reaction with NH+4 ions according to the following reaction. NH+4 + OH- ↔ NH3 + H2O Chemical equilibrium : A chemical reaction is in a state of equilibrium when the amount of products lost per second by the forward reaction exactly equals the amounts of product lost per second by the reverse reaction.the quantities of products & reactants present at equilibrium are related to each other by equilibrium constant expression. 3H2 + N2 ↔ 2NH3 K = [ NH3]2 / [H2]2 [N2] K= equilibrium constant When the equilibrium constant is greater than 102 , more of the reactants have been converted to products. Thus the products are favored in an equilibrium reaction whose k > 102. When the equilibrium constant is less than 10-2 , only a very small amount of product is formed. Therefore , the reactants are favored in an equilibrium reaction whose K is < 10-2. If the K is between 10-2 - 102 neither product nor reactant is greatly favored. The le - chatelier principle  “ if a system at equilibrium is disturbed by an applied stress , the system changes in such way that this external stress is minimized “ Many examples of this principle are found in living system. 1. Oxygen needed by the body is carried in the blood stream to the cells as oxyhemoglobin Hemoglobin + 4O2 ↔ oxyhemoglobin 2. Glucose - 1 – phosphate enzyme G- 6 –P K = [ G – 6 - P ] / [ G – 1 P] = 2 The important equilibrium in aqueous solute is the one involved in the ionization of water. 2H2O ↔ H3O+ + OH- The tendency of H2O to ionize is very small. The concentration of H+ is 1×10-7 M. the concentration of OH- is 1×10-7 M. only one of every 550×106 water molecules is ionized K = [ H3O+] [ OH-] / [H2O]2 K [H2O]2 = K- = [ H3O+] [OH-] The ionization of H2O is simply written as H2 O H+ + OH - [ H3O +] = [ H +] K = KW = [ H+ ] [ OH-] = [ 1x10 -7] [ 1 x10-7] = 1x 10 -14 [ H+ ] = [ OH- ] Neutral solution [ H +] > [ OH - ] Acidic solution [ H + ] < [ OH - ] Basic solution Henderson – Hasselbach equation pH = pKa + log [ salt] / [ acid] pH = pKa ( maximum buffer capacity) Control of pH in body fluids: In the human body the blood plasma has a normal PH of 7.4. If the PH fall below 7 or rise above 7.8 the results would be fatal. The PH of the body is controlled by buffers systems which are very effective in protecting this fluid from large changes in PH. If 1ml of 10M HCL were added to 1L of un buffered normal saline (0.15M NACL) at 7 pH ,the pH would fall to 2. But if 1ml of 10M HCL is added to 1L of blood at PH 7.4, the PH will drop to only 7.2 1 L of ( 0.15 M NaCL pH = 7) + 1 ml of (10 M HCL) ↔ pH = 2 1 L of blood (pH 7.4) + 1 ml of ( 10 M HCL) ↔ pH = 7.2 1) The major buffer system in blood is carbonic acid - bicarbonate system. H2CO3 ↔HCO -3 + H+ adding a strong acid to the system will increase the concentration of H+ , shifting the reaction to the left and forming more carbonic acid. H2CO3→HCO-3 + H+ but carbonic acid is unstable and will decomposes to CO2 and H2O. H2CO3 → CO2 + H2O the carbon dioxide can be removed from the blood and exhaled by lungs. Various factors can cause abnormal increase in acid levels in the blood. such factors are hypoventilation , ingestion of excess acids , excess loss of bicarbonate or deceased excretion of H+ through kidney failure. All of these conditions will cause an increase in H+ levels in the blood and a decrease in the concentration of basic components such as HCO-3. when the PH of blood drops to 7.1 the condition is known as acidosis ( either respiratory system or metabolic , if the origin is other than respiratory , the body has many ways to restore the blood PH to normal)  It can expel the excess CO2 formed from H2CO3 through an increase in the rate of breathing.  It can increase the excretion of H+ and retention of H2CO3 by kidneys resulting in acidic urine. The bicarbonate buffer system also protects against an addition of strong base to the system. A base will react with the H+ ions to produce water, decreasing the H+ ions concentration in the system. This will drive the reaction to the right H2CO3 ↔ HCO-3 + H+ An increase in base in blood can occur in cases of hyperventilation ,excessive ingestion of basic substances such as antacids. The pH of blood can increase to 7.8 resulting in a condition known as alkalosis. the body can return the pH to normal by many ways  Decrease in expulsion of CO2 by the lungs.  An increase excretion of HCO-3 by kidneys , resulting in an alkaline urine Bicarbonate Buffer  The principal extracellular buffer, comprising carbonic acid (the proton donor) and bicarbonate (the proton acceptor). It functions in the same way as other conjugate acid base pairs. However, there are important differences: 1. The base constituent, bicarbonate (HCO3 -) is regulated by kidneys. 2. The acid component (H2CO3) is regulated by pulmonary ventilation. Thus, bicarbonate buffer is subject to regulation by kidneys and lungs. The acid component of this buffer (also called respiratory component) is generated from dissolved carbon dioxide [CO2(d)] and water, by the reaction shown below. The reaction is catalyzed by the enzyme carbonic anhydras The dissolved carbon dioxide in blood circulation is in equilibrium with the gaseous carbon dioxide [CO 2(g)] in the air space of the lungs. As a result, concentration of carbonic acid is ultimately dependent on the partial pressure of carbon dioxide in the gas phase. Carbonic acid can dissociate to yield bicarbonate. Thus, reversible equilibrium exist between the gaseous carbon dioxide in the lungs and the bicarbonate ions in blood plasma, as shown below: Action of the bicarbonate buffer involves these equilibrium. For instance, when an acid is added to blood, concentration of H rises. The latter is taken up by HCO3 – resulting in the rise of concentration of carbonic acid (Step 1). This causes the Step 2 to go forward, and the concentration of carbon dioxide (d) in the blood rises. This in turn results in an increase in the pressure of carbon dioxide in the gas phase in the lungs (Step 3), and the extra carbon dioxide is exhaled through increased rate of breathing. Reverse series of reactions occur when an alkali (OH) is added. It is taken up by carbonic acid to form HCO3 -. Concentration of carbonic acid falls momentarily, but is quickly replenished from large pool of gaseous carbon dioxide. Rate of breathing decreases under these circumstances so that the carbon dioxide is retained and dissolved in water to form carbonic acid. Bicarbonate buffer is highly effective: Bicarbonate buffer system is an effective physiological buffer because of its equilibration with a large reserve of gaseous carbon dioxide in the air space of the lungs. Since pK’ of carbonic acid is 6.1, the bicarbonate buffer should be most effective at or around pH of 6.1 (i.e. 6.1 1) as a buffer is most effective when pH equals pK’. However, bicarbonate buffer is highly effective at the physiological pH of 7.4 also because of its equilibration with gaseous carbon dioxide. Carbonic anhydrase is the principle enzyme that catalyzes generation of HCO3 -. Decreased activity of this enzyme, therefore, results in decreased plasma bicarbonate concentration. Consequently, the ratio of bicarbonate to carbonic acid (normally 20) tends to fall, resulting in a fall of pH 2) Another buffer active mainly with the cell is the phosphate buffer system. H2PO-4↔ HPO-24 Adding strong acid will drive the reaction to the left , increasing the concentration of H2PO-4. Adding strong base will drive the reaction to right as H+ reads with base to form H2O.the kidney remove any excess of HPO-24 and H2PO-4 from the body