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Clinical Analytical Chemistry CLS 231. Buffers of Biological & Clinical Significance. Lecture 4 Lecturer: Amal Abu Mostafa. Revision for Lecture # 3.
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Clinical Analytical Chemistry CLS 231 Buffers of Biological & Clinical Significance Lecture 4 Lecturer: Amal Abu Mostafa
Revision for Lecture # 3 Calculate the pH of a solution prepared by adding 10 ml of 0.1 M acetic acid to 20 ml of 0.1 M sodium acetate. Assume Ka for acetic acid, Ka = 1.75 × 10-5 M. We need to calculate the molarity of the acid and salt in the mixture.
Buffers of Biological & Clinical Significance • What makes a “Good” buffer • In 1966 Norman Good and colleagues developed criteria for buffers for biological systems. • A pKa between 6 and 8. • Solubility in water. • Exclusion by biological membranes. • Minimal salt effects. • Minimal effects on dissociation from changes in temperature and concentration. • Minimal interactions between buffer components and critical reaction components. • Chemical stability. • Light absorption outside of wavelengths used for assay readout. • Components should be easy to obtain and prepare.
1- Optimal buffering at a neutral pH • Most biochemical reactions have an optimal pH in the range of 6–8, so buffers for these reactions need to have pKas that support buffering at these pH values. • 2- Solubility in Water • Most biochemical reactions occur in aqueous conditions, so your buffering components should be soluble in water. • If for some reason, you will be using a solvent other than water, make sure you understand how that solvent affects the dissociation of your buffer components.
4- Minimal salt interactions • If the system to be studied requires salts, appropriate ions can be added. However, using an ionic buffer can adversely affect the reaction if reaction studied is affected by salts. • 5- Minimal effects on dissociation from changes in • A- concentration • Changes in dissociation resulting from changes in concentration are usually small, and most buffers can be made as stock solutions that are diluted to working solutions. However, some buffers do show a significant change in pH upon dilution.
Title text to go here • B- Minimal effects on dissociation from changes in T • Changes in temperature can affect dissociation as well • If you have a Tris buffer prepared at 20°C with a pKa of 8.3, it would be an effective buffer for many biochemical reactions (pH 7.3–9.3), but the same Tris buffer used at 4°C becomes a poor buffer at pH 7.3 because its pKa shifts to 8.8. • So the take home message: Prepare the buffer at the temperature at which you intend to use it.
6- Minimal interactions between buffer components and critical reaction components • If a complex forms between the buffer and a required cofactor, say a metal cation like zinc or magnesium, your reaction might be compromised. • For example calcium precipitates as calcium phosphate in phosphate buffers. Not only would any Ca2+-requiring reactions be compromised, but the buffering capacity of the phosphate buffer also is affected. • 7- Chemical Stability • The buffer should be stable and not break down under working conditions. It should not oxidize or be affected by the system in which it is being used.
The biological buffer systems: • Biochemical reactions are especially sensitive to pH. • Most biological molecules contain groups of atoms that may be charged or neutral depending on pH, and whether these groups are charged or neutral has a significant effect on the biological activity of the molecule. • In all multicellular organisms, the fluid within the cell and the fluids surrounding the cells have a characteristic and nearly constant pH. • This pH is maintained in a number of ways, and one of the most important is through buffer systems. • Two important biological buffer systems are: • The dihydrogen phosphate system. • The carbonic acid system.
The phosphate buffer system: • The phosphate buffer system operates in the internal fluid of all cells (in the cytoplasm of all cells).° • This buffer system consists of dihydrogen phosphate ions H2PO-4 as proton donor (acid), and hydrogen phosphate ions HPO4 -2 as proton acceptor (base): • The value of Ka for this equilibrium is 6.23× 10-8 at 25°C. • The phosphate buffer system is maximally effective at a pH close to its pKa. • Buffer solutions are most effective at maintaining a pH near the value of thepKa . • In mammals, cellular fluid has a pH in the range 6.9 to 7.4 ,and the phosphate buffer is effective of maintaining this pH range.
The carbonic acid system • In blood plasma, the carbonic acid and hydrogen carbonate ion equilibrium buffers the pH. • In this buffer, carbonic acid (H2CO3) is the proton donor(acid) and hydrogen carbonate ion (HCO3-) is the proton acceptor (base): • In blood plasma, the concentration of hydrogen carbonate ion (HCO3-) is about twenty times the concentration of carbonic acid. • The pH of arterial blood plasma is 7.4 • If the pH falls below this normal value, a condition called acidosis is produced.
If the pH rises above the normal value, the condition is called alkalosis. The concentration of (HCO3-) and (H2CO3) are controlled by two independent physiological systems. Carbonic acid concentration is controlled by respiration, that is through the lungs. Carbonic acid is in equilibrium with dissolved carbon dioxide gas. H2CO3 (aq) CO2 (aq) + H2O (l) An enzyme called carbonic anhydrase catalyzes the conversion of carbonic acid to dissolved carbon dioxide. In the lungs, excess dissolved CO2 is exhaled as CO2gas. CO2 (aq) CO2 (g) The carbonic acid system
The carbonic acid system • The concentration of hydrogen carbonate ions is controlled through the kidneys. • Excess hydrogen carbonate ions are excreted in the urine. The much higher concentration of HCO3- over that of H2CO3 in blood plasma allows the buffer to respond effectively to the most common materials that are released into the blood. • Normal metabolism releases mainly acidic materials: carboxylic acids such as lactic acid (Hlac). • These acids react with hydrogen carbonate ion and form carbonic acid. • Hlac (aq) + HCO3- (aq) Lac- (aq) +H2CO3(aq) • The carbonic acid is converted through the action of the enzyme carbonic anhydrase into aqueous carbon dioxide. • H2CO3 (aq) CO2 (aq) + H2O (l)
An increase in CO2 (aq) concentration stimulates increased breathing, and the excess carbon dioxide is released into the air in the lungs. • Human blood plasma normally has a pH close to 7.4 • Should the pH-regulating mechanisms fail or be overwhelmed, as may happen in severe uncontrolled diabetes when an overproduction of metabolic acids causes acidosis, the pH of the blood can fall to 6.8 or below, leading to cell damage and death. In other diseases the pH may rise to lethal levels.
