Acid Base Balance in the body Maintenance of pH in body fluids Specially the blood Normal Blood pH is 7.35-7.45 Metabolism constantly produce acids Pose a threat to blood pH

Types of acids in the body Volatile acid Can leave solution and enter the atmosphere (e.g. carbonic acid) Fixed acids Acids that do not leave solution (e.g. sulfuric and phosphoric acids) Organic acids Participants in or by-products of aerobic metabolism

Common Acids produced in the body Carbonic Acid Lactic Acid Phosphoric Acid Sulphuric Acid Aceto Acetic acid Beta hydroxy butyric acid Pyruvic Acid Oxalic Acid

Medicine used in daily life Are mostly in acid forms. for examples Aspirin Ponston Antibiotics HCl Anti acid medicine

These produce a change in pH Immediate correction of pH is required. If not done , poses threat to life Change on either side of not greater than 0.2 pH unit is desirable and tolerable Slight fall than 6.8 as seen in DKA may cause death Slight rise than 7.8, may also cause death Correction is done by body buffers.

Learning objectives Buffer, conjugate acids & bases. Buffering actions & buffering capacity. Factors which determine buffering capacity. pKa values of some important physiological buffers and their role.

Buffers Buffers are important in biochemical processes, whether in plasma or in the cytosol of cells. Buffers assure biological reactions occur under conditions of optimal pH. They do this by controlling the hydrogen ion concentration of solutions.

Importance of Buffers Control of pH is an essential property of biological systems pH of Human blood plasma is maintained between 7.35 – 7.45 within 0.2 pH. unit. Values outside not compatible with life Most enzyme catalyzing various intracellular and intercellular reaction, do so at some definite pH / within a very narrow range.

Importance of Buffers Metabolic reactions constantly produce acids / bases but not in balanced amount. Which causes a pH shift, not suitable for optimum activity. This shift needs proper compensation. This is accomplished in Biochemical system mainly by Buffers.

What is buffer? A buffer solution is one that on the addition of an acid or alkali resists a change in pH. Buffer solution consists of a mixture of weak Bronsted Acid and a salt of its conjugate base, e.g. CH3COOH and CH3COONa or NH4OH and NH4Cl

Principle Behind Buffer Action Buffers are composed of weak acids and their salts. A salt is the acid minus its proton. Weak acids and their salts have two properties that are important for buffering action. First, weak acids are a reserve of the protons that neutralize OH- and prevent the solution from becoming alkaline.

Principle Behind Buffer Action Salts of weak acids are strong bases and prevent the solution from becoming acidic. Both components are needed and both are interchangeable through the loss (or gain) of a single proton.

How Buffer solution controls change in pH When alkali is added to a buffer solution, the un-dissociated Bronsted acid present in mixture ionizes and produces H+ ions. These H+ ions reacts with OH- of added alkali to form H2O CH3COOH + OH-  CH3COO- + H2O CH3COOH  CH3COO- + H+

A buffer’s power lies in its reserves. A buffer is at optimal strength when there is an equal amount of HA and A- in solution. This will only occur when the pH of the solution equals the pKa of the acid’s group. Adding OH- causes the buffer to respond by calling on the reserve pool of HA. A- is formed at the expense of HA.

Buffer’s Power This continues until all the excess OH- is neutralized. At the end the salt pool has increased (and the acid pool has decreased) by the same number of moles of base that were added.

OH- Neutralized OH- Reserve acid Reserve salt 11 moles 11 moles HA Reserve acid 11 moles Reserve salt 11 moles A- OH- Neutralized A- OH- A- A- A- A- 6 moles 16 moles

How Buffer solution controls change in pH (Contd.) Addition of alkali facilitates further dissociation of the available CH3COOH to furnish additional proton, which counters the effect and thus no change in pH.

When acid is added to an acetate buffer. H+ + CH3COO-  CH3COOH Proton added (i.e. HCl) combines exceedingly rapidly with the CH3COO- ion present in buffer mixture (as Potassium Acetate) and net change in pH is very negligible. This change in pH would have been very large if conjugate base were absent.

Buffer action or capacity of a buffer depend upon two factors Molar concentration of buffer component have direct relationship with buffer capacity. Buffer conc. is defined as the sum of the conc. of the weak acid and its conjugate base. 0.1M acetate buffer 0.05M of CH3COOH and 0.05 M of CH3COONa in one litre of H2O. It could also contain 0.065M of CH3COOH and 0.035M of CH3COONa.

Ratio of the conc. of conjugate base to the conc. of the weak acid. Buffer action or capacity of a buffer depend upon two factors. (Contd.) Ratio of the conc. of conjugate base to the conc. of the weak acid. Greater the conc. of weak acid, more the buffer will be effective to alkali addition and vice versa.

Buffer selection Selection of a buffer to be effective at the desired pH value, depends on pKa value of the conjugate acid (weak acid), therefore For a buffer to be effective at pH 5.0, we should select a weak acid having pKa value of 5.0 or close to 5.0 Buffer should be prepared with molar conc. of component high but compatible with other feature of the system.

Buffer selection Too high conc. of salt frequently inhibits the activity of enzyme or other Physiological systems. Solubility of the buffer components may also limit the conc. that can be employed.

pKa for some buffers commonly used in Biochemistry. S.No Compound pKa 1 Pyrophosphoric Acid 0.9 2 Histidine 1.8 3 E.D.T.A 2.0 4 Citric Acid 3.1 5 Acetic Acid 4.7 6 Carbonic Acid 6.1 7 Triethanolamine 7.8 8 Tris HMAE 8.0 9 Diethanolamine 8.9 10 Ethanolamine 9.5

Important Physiological buffers Many biological compounds have the ability to ionize These include organic acids, amino acids, proteins, nucleotide etc. pH of more biological fluid is near 7.0 and dissociation of most of these compounds is complete there.

Important Physiological buffers pH control in animals is a complex phenomenon. Buffers and active pH controlling elements are present in circulating blood. These are : CO2 and HCO3, NaH2PO4 and Na2HPO4 oxygenated and non oxy. forms of Hb., Plasma proteins.

Buffer Systems in Body Fluids Figure 27.7

H2CO3  CO2 + H2O  H2CO3  H+ + HCO3- It is the main buffer system of plasma and is responsible for removal of CO2 produced in the tissues during metabolism.

The Carbonic Acid-Bicarbonate Buffer System Figure 27.9a, b

HPO4-2 / H2PO4- It is important buffer system of ICF. The pKa is 7 HPO4-2 / H2PO4- It is important buffer system of ICF. The pKa is 7.21, very near to physiological pH, very effective buffering capacity.

HHbO2  H+ + HbO2 HHb  H+ + Hb- Main buffer system in RBCs, H+ ions liberated by H2CO3 are taken up by Hb in RBCs and change in pH is prevented.

Amino Acid Buffers Figure 27.8