Bt 202 aug 12 2011 ppt1997-2004
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Bt 202 aug 12 2011 ppt1997-2004 Presentation Transcript

  • 1. BT-202 Biochemistry Aug 16, 2011
  • 2.  
  • 3.
    • Acids and Bases: Definitions
    • Ionization of water
    • pH Scale
    • Dissociation of Weak acids and bases
    • Buffers
  • 4. Vinegar and Aspirin
  • 5. Soaps, toothpaste and cleaning agents.
  • 6.  
  • 7.  
  • 8.  
  • 9.  
  • 10.  
  • 11.  
  • 12.  
  • 13.  
  • 14.  
  • 15.  
  • 16.  
  • 17.  
  • 18.  
  • 19.  
  • 20.  
  • 21. HF + H 2 O H 3 O + + F -
  • 22.  
  • 23. Learning check!
    • Label the acid, base, conjugate acid, and conjugate base in each reaction:
    • H 2 SO 4 + H 2 O HSO 4 - + H 3 O +
    • HCl + OH - Cl - + H 2 O
    • HCl + NH 3 Cl - + NH 4 +
    • NH 3 + H 2 O NH 4 + + OH -
  • 24.
    • Identify the conjugate base in each of the pairs below?
    • RCOOH, RCOO -
    • RNH 2 , RNH 3 +
    • H 2 PO - ,H 3 PO 4
    • H 2 CO 3 , HCO 3 -
  • 25.  
  • 26.  
  • 27.  
  • 28. Proton donor
  • 29.
    • Proton Hopping
    • Ionization can be measured by its electrical conductivity.
    • Movement of hydronium ion is very fast due to “proton hopping”.
    • No single proton moves over a long distance but a series of proton hops between water molecules which are hydrogen bonded.
  • 30.  
  • 31.
    • Most of the acid base reactions in the aqueous solution are exceptionally very fast.
    • Proton hopping also plays a role in biological proton transfer reactions.
  • 32.  
  • 33.
    • Reversible ionization is very crucial to the role of water in cellular function.
    • We need to express the extent of ionization in quantitative terms.
    • The position of equilirium for any chemical reaction is given by its EQUILIBRIUM CONSTANT, K eq .
  • 34.
    • For any generalized reaction:
    • Equilibrium constants are dimensionless but biochemists have retained the concentration units (M) or molarity.
  • 35.
    • Equilibrium constant for a reversible reaction of water is:
    • K eq = [H + ] [OH - ]
            • [H 2 O]
            • In pure water at 25°C, the concentration of water is 55.5 M.
  • 36.  
  • 37.  
  • 38.
    • At this pH the concentration of [H + ] and [OH - ] is 1 X 10 -7 .
    • Whenever [H + ] is greater than 1 X 10 -7 [OH - ] is less and visa versa.
  • 39. Learning check!
    • What is the concentration of H + in a solution of 0.1 M NaOH?
    • What is the [H + ] of human saliva if its [OH - ] is 4 x 10 -8 M? Is human saliva acidic, basic, or neutral?
    • What is the [H + ] of a sample of lake water with [OH - ] of 4.0 x 10 -9 M? Is the lake acidic, basic, or neutral?
  • 40. Learning check!
    • What is the concentration of OH - in a solution with an H + concentration of 1.3 X 10 -4 M?
  • 41.
    • If an acidic solution is defined as a solution water where H 3 O + is greater than OH - , what can be said of an acidic solution?
    •  [OH - ]>[H 3 O + ]
    •  The concentration of H 3 O + must be greater than 10 -7.
    •  The concentration of H 3 O + is always 10 -7 if the solution is made in water.
    •  The solution feels slippery and might taste bitter.
  • 42.  
  • 43.  
  • 44.  
  • 45. pH of some aqueous fluids. pH scale ranges from 0-14.
  • 46.  
  • 47.
    • Measuring pH:
    • pH indicator is a substance which changes color around a particular pH Eg. Litmus paper.
    • Universal indicator is a mixture of indicators such that it changes color over a range of pH (Thymol blue, phenolpthalein, bromophenol blue etc).
    • pH meter is used in laboratory.
  • 48. Learning Check!
    • If the [H + ] of a solution is 0.1M, What is the [OH - ]? What is the pH?
    • What is the difference in [H + ] between solution at pH 4 and pH 7?
      • 3M
      • 1000-fold difference in [H + ]
      • 3-fold difference in [H + ]
      • 11M
  • 49.
    • pH in biological system:
    • -catalytic activity of enzymes.
    • -blood pH changes can be life- threatning. pH blood is 7.4.
    • -acidosis.
    • -alkalosis.
  • 50.  
  • 51. ACIDS AND BASES-STRENGTH
  • 52.
    • Because strong acids completely ionize, the [H + ] of a solution made with a strong acid is easily figured out, since it is equal to the molarity of the solution.
    ACIDS AND BASES-STRENGTH
  • 53. Learning Check!
    • Write out the acid dissociation reaction for hydrochloric acid?
    • Calculate the pH of a solution of 5.0 X 10 -4 M HCl?
    • Write out the acid dissociation reaction for Sodiaum hydroxide?
    • Calculate the pH of a solution of 7.0 X 10 -5 M NaOH?
  • 54. ACIDS AND BASES-STRENGTH
  • 55. ACIDS AND BASES-STRENGTH
  • 56. ACIDS AND BASES-STRENGTH
    • The tendency of any acid to lose proton and form its conjugate base in water is defined by the equilibrium constant ( K eq ) for reversible reaction:
    • HA H + + A -
    • K eq = [H + ] [A - ] = K a
    • [HA]
  • 57. ACIDS AND BASES-STRENGTH
    • The degree of dissociation of a weak acid in water is described by the acid dissociation constant, K a . It is also called the ionization constant.
    • As seen in the figure.
  • 58. ACIDS AND BASES-STRENGTH
  • 59. ACIDS AND BASES-STRENGTH
    • Also included in the figure is p K a which is analogous to pH
    • p K a = log 1/ K a = -log K a
    • eg.
    • K a of acetic acid is 1.74 x 10 –5 .
    • pK a = –log (1.74 x 10 –5 ) = 4.76.
  • 60.
    • Experimental determination of amount of acid in a given solution is done by titration. It also reveals the p K a .
    • Concentration of the acid in original solution can be calculated from volume and concentration of NaOH added.
  • 61. ACIDS AND BASES-STRENGTH
  • 62. ACIDS AND BASES-STRENGTH
  • 63. ACIDS AND BASES-STRENGTH
    • The stronger the acid, the lower its p K a ; the stronger the base the higher its p K a .
    • p K a can be determined experimentally ; it is the pH at the midpoint of a titration curve for the acid and base.
  • 64. Learning Check!
    • Which of the following aqueous solutions has the lowest pH: 0.1 M HCl; 0.1 M acetic acid (pKa = 4.86); 0.1 M formic acid (pKa = 3.75)?
  • 65. Buffer
  • 66. BUFFER
    • Buffers are aqueous systems that tend to resist changes in pH when small amounts of acid (H + ) or base (OH - ) are added.
    • A buffer system consists of weak acid (proton donor) and its conjugate base (proton acceptor).
  • 67. Buffer The acetic acid-acetate pair as a buffer system
  • 68. Buffers
    • Buffering region
    • This is the zone in the titration curve of a weak acid which extends 1 pH unit on either side of the p K a . In this zone additon of small amount of either H + or OH - has much less effect on pH than region outside this zone.
  • 69. THE HENDERSON–HASSELBALCH EQUATION
    • The extent of ionization of a weak acid (the pK a ) influences the final concentration of H + ions (the pH) of the solution, there must be a relationship between pH and the pK a of a weak acid. This relationship is given by the Henderson–Hasselbalch equation :
    • pH  =  pK a   +  log [conjugate base] [acid]
  • 70. THE HENDERSON–HASSELBALCH EQUATION
    • Calculate p K a , given pH and the molar ratio of proton donor and acceptor;
    • Calculate pH, given p K a and the molar ratio of proton donor and acceptor; and
    • Calculate the molar ratio of proton donor and acceptor , given p K a and pH.
  • 71. Learning check!
    • Which of the following compounds would be the best buffer at pH 5.0: formic acid (pKa = 3.8), acetic acid (pKa=4.76) or ethylamine (pka=9.0)?
  • 72. Learning check!
    • (a) What is the pH of a mixture of 0.042 M NaH 2 PO 4 and 0.058 M Na 2 HPO 4 ?
    • (b) If 1.0 ml of 10.0N NaOH is added to a liter of the buffer prepared in (a), how much will the pH change?
    • (c) If 1.0 ml of 10.0 N NaOH is added to a liter of pure water at pH 7.0, what is the final pH? Compare with (c)?
  • 73. Buffers in cells and tissues Cytoplasm of cells contain many amino acids with functional groups that are weak acids or weak bases
  • 74. Buffers in cells and tissues Blood plasma is buffered inpart by the bicarbonate system.
  • 75.