buffer and_ ph _meter_biochemistry .pptx

Buffer chemistry

1. BI11.2: Preparation of buffers and pH meter

2. pH & its significance

3. All biochemical reactions in vivo and in vitro are greatly influenced by hydrogen ions of surrounding medium. The convenient way of expressing hydrogen ions concentration is by the term pH which is defined as the negative logarithm of hydrogen ion concentration. Enzymes are optimally active at a particular H + ion concentration.

4. Enzymes are optimally active at a particular H + ion concentration. Water is dissociated to H + and OH - ions. pH = Log [H + ] Eg : pH of a solution having hydrogen ion concentration 0.00000001 ( 10 -7 ) is 7.

5. The pH values of some important biological fluids are as follow. Body fluids pH Pancreatic juice 8.8 Bile 7.6 Blood (at 38°C) 7.35 Saliva & human milk 6.7 Urine 6.0 Gastric juice 1.77

6. What is buffer ?

7. The H + ion concentration varies very little in any biochemical fluid or environment. This variation is arrested by some bases or acids which respectively absorb or donate H + ions depending upon situation. This phenomenon is called Buffering.

8. A buffer is a mixture of a weak acid and its alkali salts or a weak base and its acid salts. The buffer prevents a change in the pH on the addition of an acid or a base. eg : Bicarbonic acid buffer On addition of strong acid such as HCl bicarbonate reacts with the acid forming carbonic acid which is a weak acid.

9. HCl + NaHCO 3 NaCl + H 2 CO 3

10. The capacity of the buffer decreases as the ratio deviates from 1. In general buffers should be used at a pH ± 1 from pk. If the ratio is beyond 50:1 or 1:50, the buffer system is considered to have lost its buffering capacity.

11. In aqueous solutions the pH ranges from 0-14. Molar concentration of H + and OH - ions in pure water is 1 X 10 -7 mol /L. With water, at neutral point [H + ] = [ OH - ] = 1 X 10 -7 mol /L Therefore; pH at neutral point is : –log [ 10 -7 mol /L ], so neutral pH = 7. More than 7 pH indicates the solution is alkaline.

12. L ess than 7 pH indicates the solution is acidic. The pH 0 would be given by 1M HCl . The pH 14 would be given by 1M NaOH . Many biochemical substances posses functional groups that are weak acids or weak bases eg : -COOH (Carboxylic acid), phosphates present on proteins, nucleic acids, coenzymes, intermediary metabolities . Their dissociation behavior influences pH and the structure and function of the entire molecule.

13. Their dissociation behavior influences pH and the structure and function of the entire molecule. The pH of a buffer solution can be calculated by Handerson Hasselbalch equation.  

14. Determination of pH

15. By using indicator: Indicators are the compounds which change the color with changes in the pH of the solution to which they are added. They are weak acids of weak bases. In unionized forms, the indicators show one color while in their ionized forms they have different colors.

16. The color of the solution in presence of an indicator which in turn depends upon the relative proportions of ionized and unionized forms of the indicators which in turn depends on the hydrogen ion concentration. For such indicators there is a definite pH range in which it is present as mixture of the ionized and unionized forms.

17. Indicator pH range color change Litmus 4.5-8.8 Red-blue Thymol blue (acid range) 1.2-2.8 Red-yellow Thymol blue (alkaline range) 8.0-9.6 Yellow-blue Methyl orange 3.0-4.4 Red-yellow Congo red 3.0-4.5 Blue-red Bromophelon blue 3.0-4.6 Yellow-blue Bromophelon green 3.8-5.4 Yellow-blue Phenolphthalein 8.3-10.0 Colorless-red Methyl red 4.4-6.2 Red-yellow Bromocresol purple 5.2-6.8 Yellow-purple Cresol red 7.2-8.8 Yellow-red Bromothymol blue 3.0-4.6 Yellow-blue

18. By using pH paper: The indicator paper given to you is accompanied by a color chart which shows different colors which the indicator exhibits at different pH values. Take a strip of indicator paper & moist it with small drop of solution whose pH has to be determined.

19. Remove the excess fluid adhering to the indicator paper and compare with the color chart of the indicator and thus determine the pH. Do not let the paper dry before comparing with indicator.

20. By using pH meter: The most accurate method for the routine measurement of pH is by pH meter in which a change in [H + ] is measured as a change in electrical potential. If a metal rod is palaced in a solution of its salts, it acquires potential.

21. If 2 different metals are dipped into the solutions of their own salts, the difference can be measured or calculated from two separate potentials . A standard electrode is thus required against which the potential of all other electrodes can be compared.

22. This is the standard hydrogen electrode consisting of platinum rod dipped in an aqueous solution with the given H + activity in which hydrogen gas is bubbled continuously at 1 atmosphere pressure. But as this is too cumbersome to be used as a reference electrode routine use, other secondary reference electrodes of known potential in relation to the standard hydrogen electrodes are used.

23. The most commonly used secondary electrode is the calomel electrode consisting of mercury- mercurous chloride in contact with a saturated solution of KCl . Its potential is pH independent.

24. In the pH meter the most commonly used pH dependent unit is the glass electrode. Certain types of borosilicate glass are permeable to H + but not to other cation and anions. Therefore; if a thin glass membrane separates 2 solutions of different pH, a potential difference is generated across the membrane, the magnitude of which is given by the equation

25. Where E = Potential, R= gas constant, T= absolute temperature, F= Faraday constant, [H + ] I = concentration on the inside which is fixed (0.1 N HCl ), [H + ] o= concentration on the outside (test sample) E  

26. The voltage measured by such a system is primarily the difference between that of the glass and the reference electrodes and it is linearly related to the pH of the test solution. The system therefore consists of the glass electrode in contact with the solution to be measured, the calomel reference electrode, a KCl bride (the KCL should flow slowly into the sample) and the measuring device (meter).

27. These are designed so that pH 7 gives a zero potential. High resistances are used so that very little current is drawn from circuit (a large current flow could change the ion concentration)

28. Precautions The glass electrode is fragile and must be handled with care. It must nit be left dry; usually dipped in 3M KCl . The temperature compenstion dial must be set before it is calibrated as potential produced is dependent on temperature.

29. The meter must be calibrated first with a standard buffer pH 7, and then with standard of pH 4 (if the test sample is expected to be acidic) or with a standard of pH 9-10 (if the test sample is expected to be alkalic ).

30. Setting of instrument Keep the SELECTOR in zero position Switch on the instrument and wait for 10 minutes. Connect the cleaned electrode the proper terminals. Adjust ZERO control to the pointer read 7 Ph.

31. Steps for pH measurement using digital pH meter The pH instrument is calibrated with commercially prepared buffers of known pH values; 3 buffers are used pH -4 (acidic), -7 (neutral), and pH 10 (alkaline). Caliberation of the system in verified using the same buffers as sample buffers (buffer of unknown pH).

32. Upon verifying the caliberation , the sample buffer is estimated for its pH. The electrode should be carefully washed after each pH determination.

33. Presented by: Dr. Krattika Singhal