P h scale, buffers, redox potential

This presentation is very useful for chemistry ,life science as well as medical students.

1. pH Scale, Buffers, Redox potential Ifra M.Sc. 3 rd SEM

2. pH Scale Danish biochemist ‘ Soren Sorensen’ used a logarithmic scale for expressing the hydrogen ion concentration. This scale was called pH scale, where p stands for power and H for Hydrogen ion concentration. He defined pH as “The pH of a solution is the negative logarithm of the concentration(in moles/ litre ) of hydrogen ions” pH= ─ log [H+] or log 1/[H+] pH gives an idea about the acidity or basicity of the solution. The ionization product of water form the basis of pH scale.

3. Since pure water is neutral, it contains equal concentration of hydrogen & hydroxyl ions. At a certain temperature, the product of the concentration of H+ & OH+ ions in pure water is known as ionic product of water at the temperature. The ionic product of water at 25°C is approx. equal to 1*10-14 In neutral solution [H+]=[OH-]= 10-7 (pH = 7). In acidic solution, pH value ranges from 0 to 7. In alkaline(basic) solution, the pH value is between 7 to 14

4. A pH scale shows a range of 0 to 14. Sometimes the expression pOH is used to denote the basicity of OH- concentration of a solution. pH=log 1/[OH-]= ─log [OH-] pH + pOH = 14 The pH of an aqueous solution can be estimated by using various indicator, dyes such as phenolphthalein and phenol red. These dyes ionize at a specific pH to produce coloured ions. In laboratories, the pH meter.

5. pH OF SOME FLUIDS FLUID pH Gastric juice 1.0 2. Lemon juice 2.0 3. Tomato juice 4.0 4. Milk , Saliva 6.5 5. Human blood 7.4 – 7.8

6. BUFFERS SOLUTIONS OF RESERVE ACIDITY AND RESERVE ALKALINITY

7. BUFFERS A buffer solution is that which tends to maintain its pH when small amounts of strong acid or base are added to it. It is a mixture of weak acid or weak base and their conjugate base or conjugate acid respectively. It contains a hydrogen ion donor & a hydrogen ion acceptor form of a weak acids and weak bases. A buffer system is most effective when the concentration of H+ & H+ acceptor is equal. Carbonic acid bicarbonate is a common buffering system in blood plasma. The weak carbonic acid dissociated into H+ & HCO3- as follows- H2CO3 → H+ + HCO3-

8. When a small amount of HCl is added to this system, H+ ions are produce from the acid combine HCO3- ions to form H2CO3. If a small amount of NaOH is added, the OH produced reacts with H+to form water molecules. Thus this system soaks the H+ or OH- produced from strong acid or base & tend to maintain the original pH. Similarly weak bases & their salts also work as buffer system. Buffer systems in the organism help in carrying on most of the biochemical reactions in a narrow pH range of 6 to 8. The blood for example maintains its constant pH of about 7.4 despite the fact that it carries a large number & variety of chemicals. Buffer system provide protection to cells & tissues against sudden change in pH.

9. Buffering in the blood The pH range of blood is normally in the range of 7.35 to 7.45. If pH decreases below this range, the symptoms of acidosis appear & death of the animal may occur at pH 7.8 . This is because the enzymes present in the blood are extremely sensitive to changes in pH. The major buffer systems of the blood are bicarbonate, phosphate, haemoglobin & proteins buffers. Haemoglobin is a good buffer because of its capacity to act as a oxygen acceptor as well as a oxygen donor. Oxyhaemoglobin ( HHbO2) is a stronger acid than than carbonic acid but haemoglobin ( HHb ) is weaker acid. When blood is circulated through the pulmonary veins( in lungs), haemoglobin is converted to oxyhaemoglobin by absorbing oxygen. Because of its acidic nature it reacts with the bicarbonates present in blood. HHbO2 + BHCO3 → BHbO2 + H2CO3 ( B= Na, K, etc )

10. The carbonic acid thus produced is decomposed into carbon dioxide & water by the enzyme carbonic anhydrase. H2CO3 → CO2 + H2O The salt of oxyhaemoglobin formed during reaction with bicarbonate converted to the salt of haemoglobin by deoxygenation in the tissues, which then reacts with the carbonic acid produced from carbon dioxide liberated in the oxidation of carbohydrates, to form bicarbonate salt and haemoglobin . BHbO2 → BHb + O2 BHb + H2CO3 → BHCO3 + HHb The haemoglobin thus liberated goes to the lungs again, where it can be oxygenated. This cyclic oxidation and deoxygenation of haemoglobin between lungs and tissues is represented as Henderson cycle.

11. Haemoglobin as buffer CO2 +H2O H2CO3 H Hb HCO3- + H+ Hb Carbonic anhydrase HCO3- Cl - Cl - CO2 ERYTHROCYTES PLASMA

12. Renal regulation of blood pH PHOSPHATE BUFFER Renal tubular cell Na+ HCO3- + H+ H2CO3 CO2 + H2O Na2HPO4 pH=7.4 Na+ NaHPO4- H+ NaH2PO4 pH=4.5 EXCRETED Na+ HCO3- BLOOD TUBULAR LUMEN

13. PROTIEN BUFFER SYSTEM RENAL TUBULAR CELL Glutamine Glutaminase Glutamate Na+ HCO3- + H+ CA H2CO3 CO2 + H2O BLOOD Na+ HCO3- TUBULAR LUMEN NH3 Na+ H+ NH4+ EXCRETED

14. pH=7.4 CO2 (H2CO3) HCO3- Lungs ( CO2 exhaled) Metabolism ( CO2 generated) Kidneys ( HCO3- generated , H* lost) Erythrocytes ( CO2 transported , HCO3- generated) pH Regulation In Blood

15. DISEASES OCCUR DUE TO pH DISTURBANCE ALKALOSIS – It is a rise in pH METABOLIC- due to increase in bicarbonate. RESPIRATORY- due to decrease in carbonic acid. Occur due to vomiting, anemia, hypokalemia and at high altitude. Compensated by hypoventilation and HCO3- excretion by kidney. ACIDOSIS – It is a decline in pH METABOLIC- due to decrease in bicarbonate. RESPIRATORY- due to an increase in carbonic acid . Occur due to diabetes, heart, liver, lungs and kidney problems. Compensated by hyperventilation and HCO3- retained by kidney.

16. Other Examples Buffering Agent pKa Useful pH Range CITRIC ACID 3.13,4.76,6.40 2.1- 7.4 ACETIC ACID 4.8 3.8-5.8 KH2PO4 7.2 6.2-8.2 CHES 9.3 8.3-10.3 BORATE 9.24 8.25-10.25

17. USES In determining pHof unknown solutions. In studying the rate of chemical reactions. In the manufacture of ethyl alcohol from molasses (pH 5-6.8). In paper manufacture , leather tanning etc. In preparing cultures in biological specimens.

18. pH of the buffer system pH of the buffer solution can be calculated if the composition of the mixture as well as the ionization constant of the weak electrolyte is known. For example in a buffer mixture of acetic acid and sodium acetate, the pH can be determined if the ionization constant of acetic acid is known. The formula for such a determination can be derived as follows-

19. Hendorson -Hassel Equation For acid , pH= pKa + (salt)/(acid) For base , pOH = pKb + (salt)/(base) pH + pOH = pKw =14 In blood, pH= pKa +log(base)/(acid)= pKa+log (HCO3-)/(H2CO3) H2CO3 H+ + HCO3-

20. REDOX POTENTIAL The process of electron transfer accompanied with the oxidation – reduction of the system. A compound losing electron is oxidized while a compound gaining it is reduced. For example ferrous ion is oxidized to ferric ion is oxidized to ferric ion by losing one electron & vice-versa. The quantitative measure of the affinity of a compound to lose or gain electron is the redox potential. Fe++ → Fe+++ + e-

21. A redox system can be compared to a dry electric cell. In a cell, the e- are transferred through a wire from one electrode to the other. This generates an electric an electric current. The capacity to gain or lose e- in such a system is electrode potential. It can be measured through a standard hydrogen potential of zero, at 1N concentration & 1 atmospheric pressure. The electrode potential of a reducing- oxidizing system can also measured in a similar way. Electrode potential in this case will be the redox potential. Redox potential of an organic compound can be measured in the laboratory by using a standard platinum electrode. E = E• + RT/ nF In [oxidant]/[ reductant ] where E= redox potential E•= Redox potential of mixture containing equimolal concentration of oxidant and reductant .

22. R= Gas constant T= absolute temperature F= Faraday number= 96500 coloumbs n= number of e- This equation is called Peter’s equation. Under normal conditions of temperature i.e 30◦C, valency change of 2 & converting into log10 [oxidant]/[ reductant ]

23. Redox potential of hydrogen involving system E=E• + RT/ nF In[oxidant]/[ reductant ]+ RT/ nF In[H+]