Ph and buffers


Published on

  • Be the first to comment

No Downloads
Total views
On SlideShare
From Embeds
Number of Embeds
Embeds 0
No embeds

No notes for slide

Ph and buffers

  2. 2. TOPICS TO BE DISCUSSED  Measurement of compounds  Solution and its strength  Preparation of imp. solutions  Acids and bases  pH  pH meter  Henderson Hasselbalch reaction  BUFFERS 2
  3. 3. MEASUREMENT OF COMPOUNDS MOLE: One mole of a compound is grams of that compound equal to its molecular weight.  EQUIVALENT WEIGHT: One gram equivalent is weight of an element or compound represent its capacity to bind or replace 1 mole of Hydrogen. EQIVALENT WEIGHT=Molecular weight/Total no. of +ve valency 3
  4. 4. CAN YOU TELL THE NUMBER OF MOLECULES IN ONE MOLE OF ANY SUBSTRATE?  WHY ARE WE LEARNING EQUIVALENT WEIGHT ?  Because every chemical reaction what ever the type may be always occurs in equivalence.At any point of time same equivalent of reactant/s react to form same equivalent of product. 4
  5. 5. STRENGTH OF SOLUTIONS  % CONC.- Parts per 100 Most frequently used is W/V. e.g. Normal saline contains 0.9% NaCl.It means 100ml of NS contains 0.9 gram NaCl. Parts Per Million(ppm): Part of substance in one million part of the solution. e.g. 10 ppm chloride solution means 10 microgram of chlorine in 1 gram of water. 5
  6. 6.  MOLARITY: No. of MOLES of solute per litre solution. e.g.NaCl has mol.wt of 58.5. So to get 1(M) solution of NaCl 58.5 gram of it should be dissolved in water to make final volume of 1 litre. Q.Di Abietes had an elevated blood glucose level of 684 mg/dL. What is the molar concentration of glucose in Di’s blood? (Hint: the molecular weight of glucose (C6H12O6) is 180 grams per mole.) 6
  7. 7.  MOLALITY: No . Of moles of solute per 1000 gram solvent. So it is easy to prepare molal solutions.  NORMALITY: It depends on equivalent weight of the solute. It is the gram equivalent no. of solute dissolved in one litre of solution. e.g. 1 (N) NaOH contains one gram equivalent i.e. 40 gram NaOH in one litre solution. 7
  8. 8. APPLICATION OF THIS FUNDAMENTALS IN LABORATORY PREPARATION OF 1N HCL Requirements 1. Conc. HCl. AR 2. Indicator- Methyl Orange 3. 1N Sodium Carbonate 4. Glasswares:250ml beaker, 25ml conical flask, 5.0ml graduated pipette,1.0ml of volumetric pipette,graduated cylinder, volumetric flask, analytical balance,reagent bottle of 125ml. 8
  9. 9.  Procedure: 1) By using a measuring cylinder transfer 100ml of 2) 3) 4) 5) 6) 7) 8) distilled water in 250ml beaker. Add 10ml of conc.HCl ,mix well. Titration: Pipette2.0ml of diluted HCl in conical flask ,by using a volumetric pipette. Add one drop of methyl orange. Titrate against 1N sodium carbonate till colour of the reaction mixture changes from red to yellow. Note the titration reading. Take two more reading. Calculate the Normality by using the formula N1V1=N2V2 9
  10. 10.  PREPARATION OF NORMAL SALINE Requirements: 1. 2L conical flask 2. 2L measuring cylinder 3. 1L volumetric flask 4. Sodium Chloride AR 5. Distilled Water 6. Analytical Balance 7. Butter paper 8. 1L reagent bottle 9. Magnetic Stirrer 10
  11. 11.  Procedure: 1. Weigh 9.0 g NaCl on an analytical balance by 2. 3. 4. 5. 6. 7. using a butter paper. Transfer it to 2L conical flask. Add 900ml of distilled water. Mix it by using Magnetic Stirrer. Transfer it to 1L volumetric flask & add distilled water upto the mark. Mix well & store in a clean and dry reagent bottle. Label it appropriately and store at room temperature.[25±5]0c 11
  12. 12.  PREPARATION OF 1N NaOH  Requirements: 1. 1.0ml volumetric pipette 2. 2.0ml serological pipette 3. 25ml conical flask 4. Phenolophthalein [Indicator] 5. 1.0ml oxalic acid 6. Beaker of 250ml 7. Sodium Hydroxide pellets 8. Polyethylene reagent bottle 9. Analytical balance 10. Watch glass 12
  13. 13.  Procedure: 1. 2. 3. 4. 5. 6. 7. 8. By using watch glass ,weigh 4.0 g of NaOH on an analytical balance. Dissolve it in about 90 ml of distilled water. By placing the pelletes in a glass beaker. Titrate it against standard 1 N HCl solution. Take atleast 3 titration reading. Find out normality of NaOH. If normality appears less than 1.0N then add few more pellets. Adjust normality by titrating against 1.0N HCl. Store at room temperature in a polyethylene bottle 13
  14. 14.  PREPARATION OF 2/3 N SULFURIC ACID Requirements: 1. 2.0ml volumetric pipette 2. 5.0 ml serological pipette 3. Conical flask,25ml 4. Measuring cylinder 5. Beaker 250ml 6. 1N NaOH 7. Phenolophthalein 8. Reagent bottle 14
  15. 15.  Procedure: 1. Take 97ml of distilled water in a beaker. 2. Add slowly with constant stirring 2.7ml of conc. Sulfuric acid 3. 4. 5. 6. 7. 8. 9. and mix thoroughly . Titrate against 1N NaOH. Find out the normality of the prepared solution. If normality is>1.0, then dilute the solution by adding distilled water. If normality is <1.0,then add little more conc. Sulfuric acid. Measure the quantity of 1N sulfuric acid : X ml. Add X/2 ml of distilled water to X ml of 1N sulfuric acid. Store in a reagent bottle at room temperature. 15
  16. 16. ACID & BASE 16
  17. 17. The Color Change of the Indicator Bromthymol Blue Fig. 19.5 17
  18. 18. What are Acid and Bases? • Definition of Svante Arrhenius (Sweden) in 1884 “An Acid is a substance that can release a proton or hydrogen ion (H+) when dissolved in water” HCl H+ + Cl- “ A Base is a substance that can release a Hydroxyl ion when dissolved in water” NaOH Na+ +OH18
  19. 19. • According to Thomas Lowry (England) or J.N. Bronsted (Denmark) working independently in 1923: “An Acid is a material that donates a proton: HCl CH3COOH NH4+ H+ + ClCH3COO- + H+ NH3 + H+ “A Base is a material that can accept a proton OH- + H+ H2O CH3COO- + H+ CH3COOH NH3 + H+ NH4+ Every ion dissociation that involves a hydrogen or 19 hydroxide ion could be considered an acid- base reaction
  20. 20. • The G.N. Lewis (1923) idea of acids and bases is broader than the Lowry- Bronsted model. The Lewis definitions are: “Acids are electron pair acceptors. H+ + eH “Bases are electron pair donors. OHOH + e- 20
  21. 21. Some Properties of Acids  Produce H+ (as H3O+) ions in water (the hydronium ion is a hydrogen ion attached to a water molecule)  Taste sour  Corrode metals  Electrolytes  React with bases to form a salt and water  pH is less than 7  Turns blue litmus paper to red “Blue to Red A-CID” 21
  22. 22.  Produce OH- ions in water  Taste bitter, chalky  Are electrolytes  Feel soapy, slippery  React with acids to form salts and water  pH greater than 7  Turns red litmus paper to blue “Basic Blue” 22
  23. 23. pH 23
  24. 24. KEY FACTS ABOUT pH  What does it mean?  What does it measure?  What ‘p’ denotes?  Why we require pH?  Who did this pioneering job?  Derivation of pH=-log[H+]  Significance of change of pH.  Biomedical importance.  How to measure it of a given fluid? 24
  25. 25. The pH scale is a way of expressing the strength of acids and bases. Instead of using very small numbers, we just use the NEGATIVE power of 10 on the Molarity of the H+ (or OH-) ion. Under 7 = acid 7 = neutral Over 7 = base 25
  26. 26.  ‘p’ denotes ‘negative logarithm of’ to the base 10  SORENSEN brought this pioneering concept in the year of 1909.  A small change in pH!!!! See what happens in reality? Q.The laboratory reported that Di Abietes’ blood pH was 7.08 (reference range7.37-7.43) What was the [H] in her blood compared with the concentration at a normal pH of 7.4? 26
  27. 27. From inspection, you can tell that her [H] is greater than normal, but less than 10 times higher. A 10-fold change in [H] changes the pH by 1 unit. For Di, the pH of 7.08-log [H], and therefore her [H] is 1x 107.08. To calculate her [H], Express- 7.08 as- 8+ 0.92. The antilog to the base 10 of 0.92 is 8.3. Thus, her [H] is 8.3x 10-8 compared to 4.0 x10-8 at pH 7.4, or slightly more than double the normal value. 27
  28. 28. MEASUREMENT OF pH pH paper Based upon change of colour of acid/base indicator. Most commonly used for biological fluids: 1.Methyl Red (pH range:4.46.2,color change:Red to yellow) 2.Bromothymol Blue(pH range 8.0-9.6,color change: yellow to blue) Test area is compared with corresponding color chart. pH meter Basically it is a Galvanic cell. Important Component: 1.Glass electrode: Very thin 0.1 mm.It contains 0.1mol/lit HCl connected to a pt. wire via Ag-AgCl combination. 2.Calomel electrode: Glass tube containing saturated KCl connected to pt wire through Hg-HgCl2 paste. 28
  29. 29. Principle of pH meter  When a pair of electrodes or a combined electrode( glass &calomel electrode) is dipped in an aqueous solution,a potential is developed across the thin glass of the bulb.The e.m.f. of complete cell(E) formed by linking of this two electrodes at a given solution temperature is therefore E=Eref – Eglass  Eref is the potential of stable calomel electrode which at normal room temp is +0.250V  Eglass is the potential of glass electrode which depends on the pH of the solution under test. 29
  30. 30. The resultant small e.m.f. can be recorded potentiometrically by using vaccum tube amplifier.Variations of pH with E may be recorded directly on the potentiometer scale graduated to read pH directly. 30
  31. 31. 31
  32. 32. 32
  33. 33. For biological systems:  Ionization of a strong acid is TOO BIG!  Ionization of water itself is way TOO LITTLE!  Ionization of a weak acid is JUST RIGHT! 33
  34. 34. Weak Acids (most acids in nature) CH3COOH Acetic Acid ~ Gels used for UTI, antimicrobial solution plastics, dyes, insecticides ears, H2CO3 Carbonic Acid ~Bicarbonate buffer system, carbonated drinks H3PO4 Phosphoric Acid ~ Drugs, fertilisers, soaps, detergents, animal feed 34
  35. 35. • Weak bases are: – Weak electrolytes – Do not contain OH – but react with H2O numbers of OH – small  Reaction with Water : Weak bases NH3(g) + H 2O HCO3 – (aq) + H2O NH4 + (aq) + OH – (aq) H2CO3 (aq) + OH-(aq) 35
  36. 36. • Weak acids are: – Weak electrolytes – Small % ionisation weak conductors  Dissociation in Water : Weak acids Polar covalent molecules Mainly stay as molecules 36
  37. 37. 37
  39. 39. The Henderson-Hasselbalch Equation Take the equilibrium ionization of a weak acid: HA(aq) + H2O(aq) = H3O+(aq) + A-(aq) Ka = [H3O+] [A-] [HA] Solving for the hydronium ion concentration gives: [H3O+] = Ka x [HA] [A-] Taking the negative logarithm of both sides: ( ) ( ) -log[H3O +] = -log Ka - log [HA] [A-] pH = -log Ka - log [HA] [A-] Generalizing for any conjugate acid-base pair : [base] [acid] ( pH = -log Ka + log ) Henderson-Hasselbalch equation 39
  40. 40. H-A H + A p H = -log [H ] [A-] pH = pKa + log [HA] a useful concept: when [H-A] = [ A ] pKa = pH + log (1) Biological fluids are often buffered (constant pH) and it is useful to know the predominant species 40 present at a given pH.
  41. 41. BUFFERS 41
  42. 42. Definition: Buffer is a solution that resists change of pH value on addition of small amount of strong acid or alkali.  Chemically a buffer system consists of a weak acid and its salt mixture, alternatively weak bases and their salt.  e.g. carbonic acid and sodium bicarbonate  One important point to note that buffers do not remove H+ ions from body.It only acts temporarily as shock absorbent to reduce free H+ ion. The H+ ion ultimately eliminated by renal mechanism. 42
  43. 43. Why buffer system required? An average metabolic activity daily produces almost 22000meq of acid. If all the acid is dissolved in body fluids at a point of time, pH of body would be <1.0 However pH of body is strictly maintained between 7.36-7.44 under normal condition, it needs to be buffered in body. Principal buffers of ECF Principal buffers of ICF  Bicarbonate  Phosphate  Phosphate  Protein[Hemoglobin]  Proteins buffer 43
  44. 44. Mechanism of action of buffers (Strong base) AddedOHOHAdded HA HA AddedH+ AH+ B+A44
  45. 45. 45
  46. 46. BICARBONATE BUFFER SYSTEM  Most important buffering system of our body.  Proper lung functioning is important as bicarbonate buffer system is directly linked up with respiration. CO2 (d) + H2O H2CO3 HCO3- + H+ CO2 (d) means dissolved carbon dioxide. Both the reaction is catalyzed by carbonic anhydrase. 46
  47. 47. How bicarbonate buffer acts NaHCO3 Na+ H+ L- HCO3- H2CO3 (weak acid) H+ + L- Na-Lactate (salt) A strong and nonvolatile acid is converted into weak[less dissociable] and volatile acid at the expense of NaHCO3. 47
  48. 48. ALKALI RESERVE Alkali reserve is represented by amount of sodium bicarbonate in blood that has not yet combined with strong and nonvolatile acid. It denotes the buffering strength of bicarbonate at that point of time. Normal ratio of NaHCO3: H2CO3 = 20: 1 48
  49. 49. HENDERSON HESSELBALCH EQUATION CALCULATION OF pH from ABG REPORT pH=pKa + log [base]/[acid] pKa = 3.5 So for bicarbonate buffer pH = pKa + log[HCO3]/[H2CO3] [H2CO3] is best estimated as [CO2]d /400 so pH= 3.5+ log400[HCO3-]/[CO2]d =3.5+ log 400 + log [HCO3]/[CO2]d =6.1 + log [HCO3]/[CO2]d Because only 3% of the gaseous CO2 is dissolved,[CO2]d = 0.03PaCO2;so pH= 6.1+ log[HCO3]/0.03PaCO2 [HCO3-] as mEq/ml PaCO2 as mm of Hg. 49
  50. 50. HEMOGLOBIN in RBC Buffering capacity of Hb ,as of many protein,depends on number of dissociable buffering group viz. acidic –COOH gr ,basic – NH2 gr, guanidino gr,and most importantly IMIDAZOLE gr. With the pH range of 7.0 to 7.8 most of the physiological buffering system of Hb is due to the “imidazole” group of amino acid “HISTIDINE” 50
  51. 51. 1.Fe++ containing group- carrying of O2 2.Imidazole N2 group which can give up or accept H+ depending upon pH of medium. Fe++ N N+ 51
  52. 52. Thus buffering capacity of Hb is due to the presence of “Imidazole” nitrogen group which remains dissociated in acidic medium and conjugate base forms. 52
  53. 53. Association of Hb and HCO3buffer in RBC  cooperate each other in buffering blood and transporting co2 to the lungs.  although no CA is present in plasma and interstitial tissue fluid ,RBC contain high amount of this enzyme and CO2 rapidly converted to H2CO3. 53
  54. 54.  BICARBONATE is transported out of RBC into the blood in exchange of Cl- ions  As the RBCs approach lungs alveoli equilibrium reverses.CO2 is released from RBCs causing more H2CO3 to dissociate into CO2 and H2O. THUS HCO3- BUFFER SYSTEM IS INTIMATELY LINKED TO THE DELIVARY OF O2 TO THE TISSUES. 54
  55. 55. The PHOSPHATE buffer  Acts to regulate intracellular pH Na2HPO4/NaH2PO4=[Alk.PO4]/[AcidPO4] Normal plasma ratio is 4:1 It plays a major role as intracellular buffer in the RBCs and other type of cells where conc. Is much higher than the blood or ICF. 55
  56. 56.  When a strong acid enters; HCl Na2HPO4 Cl- Na+ H+ H+ NaH2PO4 + Na+HPO4= Na+HPO4= NaCl 56
  57. 57.  When an alkali enters; NaH2PO4 NaOH Na+HPO4= Na+ H+ OH- H2O + Na2HPO4 Thus a healthy KIDNEY is necessary . 57
  58. 58. The PROTEIN buffer  Blood contains high amount of extracellular protein e.g. Albumin which contribute to its buffering capacity. Na+Pr- / H+Pr- = [salt]/[acid] In acidic medium: NH2 group takes up H+ ions and protein becomes positively charged. In alkaline medium: here protein acts as an acid.Acidic COOH gr dissociates & gives H+ .It combines with added OH- to form Water. Proteins becomes negatively charged. 58
  60. 60. 60