Using The Henderson-Hasselbalch Solutions Of Specific Ph

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Objectives: To employ the Henderson-Hasselbalch equation for calculations of amounts of a weak acid and its conjugate base required to prepare the buffer solutions of specific pH. Examine pH behavior of buffer solutions using previously studied methods of measuring pH. Reinforce the proper use of volumetric glassware in preparation of solutions of a required concentration. Practice with titration technique in the determination of a buffer capacity. Introduction: A buffer solution is one that is resistant to change in pH value when small amounts of a strong acid or base are added. For example, when 0.01 mole of a strong acid or base are added to distilled water, the pH drops to 2 with the acid and rises to 12 with the base. If the same amount of a strong acid or base is added to an acetic acid – sodium acetate buffer, the pH may only change a fraction of a unit. Buffers are important in many areas of chemistry, biochemistry, and biology. When the pH must be controlled during the course of a reaction, the solutions are often buffered. This is often the case in biochemistry when enzymes or proteins are being studied. Our blood is buffered to a pH of 7.4. Variations of a few tenths of a pH unit can cause illness or death in humans. Acidosis is the condition when pH drops too low. Alkalosis results when the pH is higher than normal. Two species are required to make a buffer solution. One is capable of reacting with OH- and the other will react with

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