Revision notes on chemical equilibrium - Application of buffer solution uses - for Advanced A/AS Level Theoretical-Physical Chemistry

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Theoretical–Physical Advanced Level Chemistry – Equilibria – Chemical Equilibrium Revision Notes PART 6.5

6.5 Case studies of uses of buffers in aqueous media

Several examples of using of buffers are described and explained e.g. buffers for calibrating pH meters, controlling the aqueous media supporting enzyme action, the carbonate buffering action in blood.

Chemical Equilibrium Notes Parts 5 & 6 Index

6.5 Case studies of the function and uses of buffers in aqueous media

  • Case study 6.5.1 Other common buffer solutions and their use in the laboratory.

    • (c) doc bPotassium hydrogen benzene–1,2–dicarboxylate is an 'all in one' buffer solution of pH 4.0

      • (i) H+(aq) + OOC–C6H4–COOH(aq) (c) doc b HOOC–C6H4–COOH(aq)

      • (ii) –OOC–C6H4–COOH(aq) + OH(aq) (c) doc b OOC–C6H4–COO(aq) + H2O(l)

      • (i) removes hydrogen ions and (ii) removes hydroxide ions.

      • Buffers can be made by mixing the salt with the original benzene–1,2–dicarboxylic acid to give buffers in the range 2.2–3.8

    • A mixture of salts of a polybasic/polyprotic acid e.g. the salts KH2PO4, Na2HPO4 and Na3PO4 from phosphoric(V) acid (a tribasic/triprotic acid) can give buffer solutions in the range pH 6–12 e.g.

      • from Na2HPO4: HPO42–(aq) + H+(aq) (c) doc b H2PO4(aq)  (removes hydrogen ions)

      • from Na3PO4: PO43–(aq) + H+(aq) (c) doc b HPO42–(aq) + H2O(l) (removes hydrogen ions)

      • from KH2PO4: H2PO4(aq) + OH(aq) (c) doc b HPO42–(aq) + H2O(l) (removes hydroxide ions)

      • from Na2HPO4: HPO42–(aq) + OH(aq) (c) doc b PO43–(aq) + H2O(l) (removes hydroxide ions)

      • The HPO42– ion is amphoteric, acting both as a proton donor and acceptor and phosphate(V) ions are important in the buffering of intracellular fluids in living organisms (see Case study 6.5.2 below).

    • Buffer solutions are used to accurately calibrate pH meters.

  • Case study 6.5.2 The importance of buffering in biological systems

    • enzyme2Case study 6.5.2a

      The importance of a stable pH cannot be over emphasised for the efficient function of enzyme–proteins which control most of the chemistry of life. Most enzymes can only operate at their maximum catalytic capacity over a narrow range of pH and three examples are shown on the left.

      6.5.2b and 6.5.2c are examples of conjugate acid–base pairs that fulfil this important buffering operation are described below.

    • 6.5.2a Optimum pH conditions enzyme function.

      • Pepsin operates best in very acid conditions around pH 1.5 to 2.5

        • Pepsin is a vertebrate endopeptidase enzyme secreted as the zymogen pepsinogen by chief cells  of gastric pits and spontaneously formed from it below pH 5 (by an intramolecular process) when it is active. This type of acid protease enzyme hydrolyses peptides to their constituent amino acids.

      • Carbonic anhydrase operates most efficiently when its close to neutral ~pH 6.5 to 7.5

        • Carbonic anhydrase is a zinc–containing metalloenzyme of e.g. vertebrate red blood cells and catalyses reactions such as

          • H2O + CO2 (c) doc b H+ + HCO3

          • a reversible reaction under different pH conditions and it speeds up carbon dioxide transport.

          • It is also involved in blood pH regulation of the kidney.

      • Trypsin is most efficient in slightly alkaline conditions pH 7.5 to 8.5

        • Trypsin is a protease enzyme secreted in an inactive form (trypsinogen) in the pancreas and is converted into an active form by enterokinase. The active form converts proteins into peptides in the digestion systems and the buffering medium is sodium hydrogen carbonate (see section 6.5.2c), and the reaction is essentially a partial hydrolysis breakdown of proteins.

      • For more on this See section 2b.6 in Enzyme structure, function and denaturing

    • 6.5.2b Inside cells hydrogenphosphate(V) ions act as the major intracellular buffer system, with contributions from organic phosphates such as glucose–6–phosphate and ATP.

      • HPO42–(aq) + H+(aq) (c) doc b H2PO4(aq) or H2PO4(aq) + OH(aq) (c) doc b HPO42–(aq)

    • 6.5.2c The major extracellular buffer is the 'carbonic acid'–'bicarbonate' or hydrogencarbonate system which enables e.g. blood, to function as an extraordinary effective buffer operating at about pH 7.

      • A whole series of linked equilibria control the pH of blood ...

        • (i) CO2(g, lungs) + aq (c) doc b CO2(aq)

          • the reversible dissolving of carbon dioxide in water

        • (ii) CO2(aq) + H2O(l) (c) doc b H2CO3(aq)

          • the formation of 'carbonic acid'

        • (iii) H2CO3(aq) (c) doc b HCO3(aq) + H+(aq)

          • the dissociation of carbonic acid into the hydrogencarbonate ion and the hydrogen ion – so decreasing pH

          • OR, the reverse reaction to remove hydrogen ions – so increasing pH

        • (v) H2CO3(aq) + OH(aq) (c) doc b HCO3(aq) + H2O(l)  

          • carbonic acid removing hydroxide ions, so lowering the pH

          • OR, the reverse, the production of hydroxide ion to increase the pH

        • (aq) = intra/extracellular fluids
      • Note the negative anions must be counter balanced by positive ions such as sodium, Na+ otherwise an unwanted potential would be set up.

      • (iv) to (iii) removes hydrogen ions and (v) to (vi) removes hydroxide ions.

      • The effectiveness of the system depends on the reservoir of dissolved carbon dioxide in the blood plasma and the gas in the lungs.

        • If hydroxide ions are removed via reaction (v) to (iv), the depleted H2CO3 is readily replaced via the reaction sequence (i) to (ii) and (ii) to (iii).

        • If hydrogen ions are removed via (iv) to (iii) the reverse sequence of 1. can restore the system to the original pH.

      • The ability of mammals to maintain a fairly constant [HCO3]/[H2CO3] ratio in blood plasma is reflected in the rate of CO2 production in the cell oxidation reactions of respiration and the rate of CO2 loss by expiration.

      • The blood plasma of man is about 7.4 and any deviation below 7.0, or above 7.8, as can happen in disease, can cause irreparable damage.

      • Intracellular and extracellular systems are very pH sensitive and small changes in pH can produce ill–effects in living organisms, hence, e.g. the bodies irritation by all except the very weakest of acids and alkalis in contact with the skin.

  • Case study 6.5.3 Shampoos

    • to complete

  • Case study 6.5.4: *

Chemical Equilibrium Notes Parts 5 & 6 Index


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