Concentration or Buffer Capacity Choosing the appropriate buffer concentration can be a little tricky depending on whether pH control is the only role of the buffer, or if ionic strength or other considerations also are impor tant. When determining the appropriate concentration for pH control, the following rule of thumb can be used to estimate a reasonable starting concentration 1. If the process or reaction in the system being buffered does not actively produce or consume protons(H), then choose a moderate buffer concentration of 50 to 100 mM 2. If the process or reaction actively produces or consumes protons(H,), then estimate the number of millimoles of H* that are involved in the process (if possible)and divide by the solu tion volume. Choose a buffer concentration at least 20x higher than the result of the estimation above The rationale behind these two steps is that a properly chosen buffer will have a 50: 50 ratio of acid to base at the target pH, therefore you will have 10x the available capacity to consume or supply protons as needed. A 10% loss of acid(and corresponding increase in base species), and vice versa, results in a 20% change n the ratio (CH3 COo][CH3COOH from the Henderson- Hasselbalch example above]) resulting in less than a 0.1 pH unit change, which is probably tolerable in the system. While most bio- molecules can withstand the level of hydrolysis that might accom- pany such a change(especially near neutral pH), it is possible that he secondary and tertiary structures of bioactive molecules might be affected Chemical Compatibility It is important to anticipate(or be able to diagnose) problems due to interaction of your buffer components with other solution components. Certain inorganic ions can form insoluble complexes with buffer components; for example, the presence of calcium wi cause phosphate to precipitate as the insoluble calcium phosphate, and amines are known to strongly bind copper. The presence of significant levels of organic solvents can limit solubility of some inorganic buffers. Potassium phosphate, for example, is more readily soluble in some organic solutions than the correspond ing sodium phosphate salt One classic example of a buffer precipitation problem occurred when a researcher was trying to prepare a sodium phosphate buffer for use with a tryptic digest, only to have the Ca(a nec-Concentration or Buffer Capacity Choosing the appropriate buffer concentration can be a little tricky depending on whether pH control is the only role of the buffer, or if ionic strength or other considerations also are impor￾tant. When determining the appropriate concentration for pH control, the following rule of thumb can be used to estimate a reasonable starting concentration. 1. If the process or reaction in the system being buffered does not actively produce or consume protons (H+ ), then choose a moderate buffer concentration of 50 to 100 mM. 2. If the process or reaction actively produces or consumes protons (H+ ), then estimate the number of millimoles of H+ that are involved in the process (if possible) and divide by the solu￾tion volume. Choose a buffer concentration at least 20¥ higher than the result of the estimation above. The rationale behind these two steps is that a properly chosen buffer will have a 50 :50 ratio of acid to base at the target pH, therefore you will have 10¥ the available capacity to consume or supply protons as needed. A 10% loss of acid (and corresponding increase in base species), and vice versa, results in a 20% change in the ratio ([CH3COO- ]/[CH3COOH from the Henderson￾Hasselbalch example above]) resulting in less than a 0.1 pH unit change, which is probably tolerable in the system. While most bio￾molecules can withstand the level of hydrolysis that might accom￾pany such a change (especially near neutral pH), it is possible that the secondary and tertiary structures of bioactive molecules might be affected. Chemical Compatibility It is important to anticipate (or be able to diagnose) problems due to interaction of your buffer components with other solution components. Certain inorganic ions can form insoluble complexes with buffer components; for example, the presence of calcium will cause phosphate to precipitate as the insoluble calcium phosphate, and amines are known to strongly bind copper. The presence of significant levels of organic solvents can limit solubility of some inorganic buffers. Potassium phosphate, for example, is more readily soluble in some organic solutions than the correspond￾ing sodium phosphate salt. One classic example of a buffer precipitation problem occurred when a researcher was trying to prepare a sodium phosphate buffer for use with a tryptic digest, only to have the Ca2+ (a nec- 34 Pfannkoch