Biological Buffer Systems
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Biological Buffer Systems Kim C. Abian Hassien Alromali George Christopher L. Cabatay Arizaldo E. Castro Weena Ross P. Dygico De La Salle University-Dasmariñas Dasmariñas, Cavite ABSTRACT Phosphate buffer preparation was accomplished via dissolution of 2.7185 g KH2PO4 and 5.2230 g K2HPO4 in distilled H2O enough to make a 250 ml buffer solution with a theoretical pH of 7.40. The solution was tested for its initial pH and further subjected to addition of 5 ml and 10 ml 0.2 M HCl and NaOH to verify resistance to extreme change in pH. Results show an experimental initial pH of 6.89 with 6.89% error versus the theoretical pH of 7.40. Apparently, HCl and NaOH addition demonstrated drastic changes in pH of the buffer system with computed percent errors of 9.62% (5 ml HCl), 13.57% (10 ml HCl), 3.51% (5 ml NaOH) and 2.70% (10 ml NaOH). From the data gathered, it can be inferred that buffer solutions are capable of maintaining a constant environment for body and biochemical processes to continue without disturbance. Avoiding erroneous results calls for consistency in weighing and mixing the specified amount of reagents and thorough washing of the pH meter electrode after each use. INTRODUCTION A buffer is a solution composed of a weak acid and its conjugate base or a weak base and its conjugate acid (1). These are solutions that materialize the importance of acid-base balance in the body via physiological regulating systems. Buffers have the capacity to resist changes in pH during addition of H+ or OH- to maintain a constant pH. It works by removing hydrogen ions when they are in excess and donates them when the solution is low of it. A pertinent example is the Bicarbonate buffer system that maintains pH of blood plasma and Extracellular fluid in its functional value (1). The resistance of buffer systems to change in pH when added with small amounts of acidic or basic substances is attributed to the existing chemical equilibrium between the weak acid and its conjugate base (2). Chemical equilibrium is defined as a condition where the rate of the forward reaction is the same with the rate of the reverse reaction (2). It is mathematically expressed as: K = [Product Concentration]a [Reactant concentration] Figure 1. Equation for Equilibrium constant, K Moreover, the pH of a buffer system can be calculated using the Henderson-Hasselbalch equation exemplified by the following formula (3):
Figure 2. Henderson-Hasselbalch Equation
where, pH is the pH of the buffer solution, pKa is the negative logarithm of the acid equilibrium constant, [A-] is the concentration of the base, and [HA] is the concentration of the acid. Since weak acids and bases have reversible reactions, they can achieve a certain condition where there is chemical equilibrium. Given the example of dissociating Acetic acid to produce acetate ion and hydrogen ion, the equilibrium constant is represented in the following equation: Ka = [CH3COO-][H+]/[CH3COOH]. Weak acids and bases are utilized for buffer systems since they do not dissociate fully and reverse reaction is possible. Strong acids and bases are not applicable since once dissociated, they are 100% dissolved in the solution (2). In addition to this, another group of buffers include synthetic organic buffer systems. The most well-known is Norman Good and fellows’ twelve synthetic organic buffers that have been selected in basis with pKa, solubility, Optical absorbance, Stability, and the like factors (4). The human body owes a lot to buffer systems since medical complications like acidosis and alkalosis can occur if these systems are absent. Irregular heartbeat, Tachycardia, Muscular cramps and muscle weakness are secondary effects of metabolic and respiratory acidosis and alkalosis. MATERIALS AND METHODS Calculation of the needed reagents for the Phosphate buffer was first accomplished using the Henderson-Hasselbalch equation and it was determined that 2.72 g of KH2PO4 and 5.22 g of K2HPO4 are needed to produce the desired buffer system with a pH of 7.40. Using the analytical balance and weighing paper, the desired quantity of the reagents were obtained (2.7185 g KH2PO4 and 5.2230 g K2HPO4) and mixed in a 250 ml volumetric flask. Distilled water was added to the mixture to facilitate dissolution of the compounds. A stopper was placed on the open end of the flask and to speed up the process, swirling action was employed. When the powder reagents are no longer visible, distilled water was added until the 250 ml mark was reached. The flask was inverted 20 times to ensure mixing (1). The pH of the phosphate buffer was measured using a digital pH meter. Afterwards, 2 sets of 25 ml buffer were prepared using 2 beakers. 5 ml 0.2 M HCl was added to the first set and its pH was also measured. Afterwards, another 5 ml of the same acid (same molarity of 0.2 M) was added. Using the digital pH meter, pH was obtained. On the second beaker, 5 ml 0.2 M NaOH was added. Just like the acid addition, extra 5 ml of the same concentration of NaOH was further added. pH of both instances where addition of the basic compound were measured (1). 25 ml of distilled water was also subjected to the procedure done with the previous two sets of phosphate buffer. The pH readings were noted and % errors of the experimental data versus their theoretical counterparts were computed. Afterwards, the remaining buffer was transferred to a clean plastic bottle, labeled, and stored for next activities’ use (1). RESULTS AND DISCUSSIONS
Initial +5.00 ml 0.200 M HCl +10.00 ml 0.200 M HCl + 5.00 ml 0.200 M NaOH +10.00 ml 0.200 M NaOH
Theoretical 7.40 7.38 7.37 7.40 7.40
Table 1. Addition of Acid and Base
pH Phosphate Buffer Experimental % Error 6.89 6.89% 6.67 9.62% 6.37 13.57% 7.14 3.51% 7.60 2.70%
Distilled Water 5.25 4.30 3.87 9.38 9.47
Table 1 shows the calculated theoretical pH of the initial Phosphate buffer and the succeeding pH values of the system after addition of varying volumes of 0.2 M HCl and NaOH. Computed percent errors are ranging from a minimum value of 2.70% to a maximum of 13.57%. It can be observed that theoretical computed values are clustered in the range of 7.37-7.40 showing the ideal minimal effect of acid/base addition to the solution’s pH. The reason why buffer can resist changes in pH is because it contains both acidic and basic constituents. Its acid and base neutralizes any acidic or basic substance added to the system. Secondary to this fact is the existing chemical equilibrium in a buffer system. Weak acids and bases have reversible reactions and thus when balance is disturbed, shifting can happen to restore equilibrium. Comparing the characteristic of buffer to distilled water, the distilled water has a neutral pH. It doesn't have any acidic or basic properties. As shown by the experimental pH values obtained, significant pH change can be observed after acid and base addition to the distilled water sample. Percent errors between the experimental and theoretical values of the measured pH can be attributed to two factors. First, the amounts of the reagents mixed. As calculated using Henderson-Hasselbalch’s equation, 2.72 g of KH2PO4 and 5.22 g of K2HPO4 are the precise amounts to be combined for the buffer formulation. Apparently, 2.7185 g KH2PO4 and 5.2230 g K2HPO4 were mixed instead of the specified amounts calculated. Another possible source of discrepancy is the pH meter’s electrode. Non-thorough washing of electrode after each use will yield to erroneous pH readings. Aside from inorganic buffer solutions, like the Phosphate buffer formulated, there are also synthetic organic buffers that are used in the field of Biology and Biochemistry (5). In 1966, Norman Good and his colleagues selected 12 synthetic organic buffers in basis with the buffers’ properties. Examples of these buffers include MOPS or 3-morpholinopropane-1-sulfonic acid and PIPES or 1,4-piperazinediethanesulfonic acid. Synthetic organic buffers impedes premature reduction of reducible compounds and some are also more efficient in maintaining pH even under changes in Carbon dioxide concentration (6) (ex. 4-(2-hydroxyethyl)-1-piperazineethanesulfonic acid or commonly abbreviated as HEPES).
Figure 3. Structure of HEPES
Figure 4. Structure of PIPES
Going to the medical and physiological importance of buffer solutions, Phosphate buffer system plays a major role in buffering renal tubular fluid and intracellular fluids. Another significant buffer solution is the Carbonic acid-Bicarbonate Buffer. It is the controlling system to maintain our blood plasma and ECF pH at normal rate. If hydrogen ions become available in excess, the system counteracts by shifting the hydrogen ions to Carbon dioxide to lessen blood plasma acidity. Furthermore, if the Phosphate buffer acts with the respiratory system, speed of breathing can be utilized too to expel extra hydrogen ions in form of CO 2 gas. The carbonic acid (H2CO3) is the hydrogen-ion donor (acid) and hydrogen carbonate ion (HCO3-) is the hydrogen-ion acceptor (base). The carbonic acid-bicarbonate buffer system functions in a way that, additional H+ is consumed by HCO3- and additional OH- is consumed by H2CO3. The concentrations of hydrogen carbonate and of carbonic acid are controlled by two independent physiological systems. For instance, the carbonic acid concentration is controlled by the respiratory system, particularly through the lungs. Carbonic acid is in equilibrium with dissolved carbon dioxide gas. On the lungs, the excess dissolved carbon dioxide is exhaled as carbon dioxide gas. On the other hand, the concentration of hydrogen carbonate ions is controlled through the kidneys. Excess hydrogen carbonate ions are excreted in the urine (7). There are specifically two abnormal conditions that may result from an excess of acid or alkali (base) within the blood, namely acidosis and alkalosis. In line with this, it must always be notes that the normal blood pH must be maintained between 7.35-7.45 to ensure proper functioning of metabolic processes and the delivery of the right amount of oxygen to tissues. Any disease or condition that affects the lungs, kidneys, metabolism, or breathing has the potential to cause acidosis or alkalosis (7). Acidosis refers to the condition when excess acid occurs in the blood resulting to a pH level of below 7.35. Acidosis is basically due to increased acid production within the body, consumption of substances that are metabolized to acids, decreased acid excretion, or increased base excretion. Acidosis can be classified into two, namely respiratory acidosis, and metabolic acidosis. Respiratory acidosis develops when there is too much carbon dioxide (an acid) in the body. This type of acidosis is usually caused by a decreased ability to remove carbon dioxide from the body through effective breathing. Its other specific causes include: chest deformities (i.e. kyphosis), chest injuries, chest muscle weakness, chronic lung disease, and overuse of sedative drugs. Metabolic acidosis, on the other hand, develops when specifically the kidneys cannot remove enough acid from the body. Its other causes include poisoning by aspirin, ethylene glycol, or methanol, and severe dehydration. Generally speaking, the effects of acidosis include mental confusion, fatigue and lethargy, breathing difficulty, and even shock or death (7). Alkalosis refers to the condition when levels of alkali become abnormally high resulting to a blood pH greater than 7.45. Alkalosis may also be caused by a low level of carbon dioxide in the blood that results from rapid or deep breathing—a condition normally exemplified by respiratory alkalosis, one of the types of alkalosis. Respiratory alkalosis is commonly caused by fever, lack of oxygen, liver disease, lung disease, salicylate poisoning, and being at high altitude. Metabolic alkalosis is caused by too much bicarbonate in the blood—which is the usual characteristic of the general term of alkalosis. The general effects when alkalosis is experienced include decreased myocardial contractility, arrhythmias, decreased cerebral blood flow, confusion, mental obtundation, neuromuscular excitability, and impaired peripheral oxygen unloading (7). REFERENCES
Legaspi, G.A. 2011. Essentials of Biochemistry Laboratory. Philippines (Evaluation copy)
Lehninger, A.L. 1976. Biochemistry. 2nd Ed. NY: Worth Publishers, Inc.
Buffers and the derivation of Henderson-Hasselbalch Equation. Retrieved http://www.chembuddy.com/?left=pH-calculation&right=pH-buffers-henderson-hasselbalch. Retrieved on June 29, 2011
(4) Good’s Buffers. Retrieved from http://www.chemicalland21.com/info/GOOD%27S %20BUFFERS.htm Retrieved on June 29, 2011
(5) Wu A.L. 1989. Use of organic buffers to reduce dehydroascorbic acid interference in analytical methods. Retrieved from http://www.freepatentsonline.com/4885240.html Retrieved on June 25, 2011 (6)
Bhattacharyya A. Yanagimachi R. 2005. Synthetic organic pH buffers can support fertilization of guinea pig eggs, but not as efficiently as bicarbonate buffer. Retrieved from http://onlinelibrary.wiley.com/doi/10.1002/mrd.1120190203/abstract. Retrieved on June 25, 2011
Acidosis and Alkalosis. Retrieved from http://www.nlm.nih.gov Retrieved on June 29, 2011