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FUNDAMENTALS OF ACID – BASE (pH and Buffers)

FUNDAMENTALS OF ACID – BASE (pH and Buffers). Herbert M. San Pedro Lecturer. Acid - Base. An important factor in the proper behavior of many biochemical phenomena (processes)

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FUNDAMENTALS OF ACID – BASE (pH and Buffers)

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  1. FUNDAMENTALS OF ACID – BASE(pH and Buffers) Herbert M. San Pedro Lecturer

  2. Acid - Base • An important factor in the proper behavior of many biochemical phenomena (processes) • Any deviation from the expected acid-base balance may lead to some disorders, most common are the alkalosis and acidosis. • Can be measured by knowing the Hydronium ion concentration [ ] (established by Sorensen) - pH

  3. DISSOCIATION PRINCIPLE • Ability to form ions in solution. • Water, plays a vital role in the determination of the degree of acidity (alkalinity) of many inorganic substances, likewise in the biochemical processes. • Water being a naïve substance, can act as both an acid or base ( amphoterism )

  4. Cont. • Water molecules themselves may react with each other and create ions (auto-ionization of water – autoprotolysis) • Eqn: H2O + H2O ↔ H3O+ + - OH • Or Simplified by: H2O ↔ H+ + - OH

  5. Cont. • The presence of such ions contribute significantly to the over-all properties of water. • Going further analytically: the auto-ionization of water is expressed in terms of equilibrium expression: K = [H+ ][- OH] / [H2O} K = dissociation constant

  6. Cont. • To further simplify the equation: it is considered that the concentration of pure water being experimentally to high (55.56 mol/L) is doubtedly affected by dissociation (making it constant). • Eqn becomes: K[H2O] = Kw (water const.) Kw = [H+ ][- OH ] = ion product of water • Simplifying things out, it is from the above eqn. wherein the concept pH was introduced and also applied to acidic and basic solutions.

  7. Cont. • pH – degree of acidity, value is affected by the amount of hydronium ion present in solution. • Low pH – High concentration of hydronium ion. (below 7) • High pH – Low concentration of hydronium ion. (above 7) • The formula: pH = - log [H+ ] : UE

  8. Sample Calculations: What is the pH of a solution whose hydrogen ion concentration is 3.2 x 10-4 mol/L. Solution: pH = - log [H+ ] = - log (3,2 x 10-4 ) = - (-3.5) pH = 3.5

  9. Cont: 2. What is the pH of a solution whose hydroxide ion concentration is 4,0 x 10-4 mol/L: (in similar fashion pOH = - log [- OH]) Considering analytically the ion product of water Kw = [H+ ][- OH] = 1 x 10-14 taking the log of both sides this becomes: pH + pOH = 14 : UE

  10. Cont. To solve the problem: [- OH] = 4.0 x 10-4 pOH = - log [- OH] = - log (4.0 x 10-4 ) = - (-3.4) pOH = 3.4 pH + pOH = 14 pH = 14 – pOH = 14 – 3.4 pH = 10.4

  11. Cont. Considering the vice-versa: [H+ ] = 10-pH : UE Given the pH of a substance to be equal to 5. Calculate its hydrogen ion concentration. [H+ ] = antilog (-pH) = antilog ( - 5) [H+ ] = 1 x 10-5 mol/L

  12. The Weak Acid / Weak Base • Many biochemicals possess functional groups that are weak acids or bases. Thus knowledge of the dissociation of WA/ WB is basic to understanding the influence of intracellular pH on structure and biologic activity. • And these therefore require a different approach in calculating for their acidity and basicity (acidity / basicity dissociation constant; Ka or Kb) : HA ↔ H+ + - A (weak acid) Ka = [H+ ][- A] / [HA]

  13. HENDERSON-HASSELBACH EQUATION • A very useful equation (UE) in calculating for pH involving weak acids/weak base and their conjugates. • Derivation: HA ↔ H+ + - A Ka = [H+ ][- A]/[HA] Ka[HA] = [H+ ][- A] [H+ ] = Ka[HA]/[- A]

  14. Cont. Log [H+ ] = Log (Ka[HA]/[- A]) = Log Ka + Log [HA]/[- A] Multiplying both sides by -1 -Log [H+ ] = - Log Ka – Log [HA]/[- A] pH = pKa – Log [HA]/[- A] pH = pKa + log [- A]/[HA] : UE

  15. BUFFERS • Solution of Weak Acid/Base and their conjugate. (salts) • Has the ability to resist change in pH after the addition of strong acids/bases. • To measure the pH changed that occurred after the addition of either strong acid/base the HHE is being used.

  16. Cont. • There are so many buffers in the human body and all of these are responsible to maintain the proper pH the body requires (physiologic buffers) • a. Bicarbonate, Phosphate, Protein Buffers • Any deviation from the desired pH may lead to disorders like acidosis or alkalosis.

  17. Acid-Base Balance • Blood place a crucial role In the over-all acid-base balance, It must retain (through the help of buffers) its normal pH (even after the addition of strong acid/bases) so as not jeopardize the biochemical processes. • The Kidney and the Lungs are the two major organs that help the blood to regulate its pH

  18. Cont. • The exhaling of CO2 by the lungs and regulation of HCO3- by the kidney, are the responsible processes how the blood is capable of maintaining its pH. > Any trouble acquired by either of the organs or conditions/situations again acquired by any persons may suggest possible imbalance to the equilibrium

  19. Cont. CO2 + H2O ↔ H2CO3 ↔ H+ + HCO3-

  20. Cont.

  21. Cont.

  22. Cont.

  23. THANK YOU!!! Long Examination on Thursday (July 15, 2010) from the Beginning –up to Last Topic -- No Meeting on Saturday (July 10, 2010)

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