《英语教学》（英文版）The Preparation of Buffers and Other Solutions-A Chemist's Perspective.pdf_大学文库

Reagents. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 39 Which Grade of Reagent Does Your Experiment Require?. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 39 Should You Question the Purity of Your Reagents?. . . . . . . . 39 What Are Your Options for Storing Reagents? . . . . . . . . . . . 40 Are All Refrigerators Created Equal? . . . . . . . . . . . . . . . . . . . 41 Safe and Unsafe Storage in Refrigerators . . . . . . . . . . . . . . . . 41 What Grades of Water Are Commonly Available in the Lab? . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 42 When Is 18MW Water Not 18MW Water? . . . . . . . . . . . . . . 44 What Is the Initial pH of the Water?. . . . . . . . . . . . . . . . . . . . 44 What Organics Can Be Present in the Water? . . . . . . . . . . . 45 What Other Problems Occur in Water Systems?. . . . . . . . . 46 Bibliography . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 47 BUFFERS Why Buffer? The primary purpose of a buffer is to control the pH of the solution. Buffers can also play secondary roles in a system, such as controlling ionic strength or solvating species, perhaps even affecting protein or nucleic acid structure or activity. Buffers are used to stabilize nucleic acids, nucleic acid–protein complexes, proteins, and biochemical reactions (whose products might be used in subsequent biochemical reactions). Complex buffer systems are used in electrophoretic systems to control pH or establish pH gradients. Can You Substitute One Buffer for Another? It is rarely a good idea to change the buffer type—that is, an amine-type buffer (e.g., Tris) for an acid-type buffer (e.g., phosphate). Generally, this invites complications due to secondary effects of the buffer on the biomolecules in the system. If the purpose of the buffer is simply pH control, there is more latitude to substitute one buffer for another than if the buffer plays other important roles in the assay. How Does a Buffer Control the pH of a Solution? Buffers are solutions that contain mixtures of weak acids and bases that make them relatively resistant to pH change. Conceptually buffers provide a ready source of both acid and base to either provide additional H+ if a reaction (process) consumes H+ , or combine with excess H+ if a reaction generates acid. 32 Pfannkoch

The most common types of buffers are mixtures of weak acid and salts of their conjugate bases, for example, acetic acid/sodium acetate. In this system the dissociation of acetic acid can be written CH3COOH→CH3COO+H where the acid dissociation constant is defined as Ka=HI [CH_ COO[ COOH Rearranging and taking the negative logarithm gives the more familiar form of the Henderson-Hasselbalch equation DH=pk+log ICHsCOo- ICH3 COOH Inspection of this equation provides several insights as to the functioning of a buffer When the concentrations of acid and conjugate base are equa log(1)=0 and the pH of the resulting solution will be equal to the pKa of the acid. The ratio of the concentrations of acid and con jugate base can differ by a factor of 10 in either direction, and the resulting pH will only change by l unit. This is how a buffer main- tains ph stability in the solution To a first approximation, the ph of a buffer solution is inde- pendent of the absolute concentration of the buffer; the pH depends only on the ratio of the acid and conjugate base present However, concentration of the buffer is important to buffer capac ity, and is considered later in this chapter When Is a Buffer not a Buffer? Simply having a weak acid and the salt of its conjugate base present in a solution doesnt ensure that the buffer will act as a buffer Buffers are most effective within t 1 pH unit of their pK Outside of that range the concentration of either the acid or its salt is so low as to provide little or no capacity for pH control Common mistakes are to select buffers without regard to the pKa of the buffer. Examples of this would be to try to use K,HPO/KH2PO4(pKa=6.7)to buffer a solution at pH 4, or to use acetic acid(pKa=4.7)to buffer near neutral pH. What Are the Criteria to Consider When Selecting a Buffer? Target pH Of primary concern is the target pH of the solution. This narrows the possible choices to those buffers with pka values within 1 pH unit of the target pH The Preparation of Buffers and Other Solutions

The most common types of buffers are mixtures of weak acids and salts of their conjugate bases, for example, acetic acid/sodium acetate. In this system the dissociation of acetic acid can be written as CH3COOH Æ CH3COO- + H+ where the acid dissociation constant is defined as Ka = [H+ ] [CH3COO- ]/[H3COOH]. Rearranging and taking the negative logarithm gives the more familiar form of the Henderson-Hasselbalch equation: Inspection of this equation provides several insights as to the functioning of a buffer. When the concentrations of acid and conjugate base are equal, log(1) = 0 and the pH of the resulting solution will be equal to the pKa of the acid. The ratio of the concentrations of acid and conjugate base can differ by a factor of 10 in either direction, and the resulting pH will only change by 1 unit. This is how a buffer maintains pH stability in the solution. To a first approximation, the pH of a buffer solution is independent of the absolute concentration of the buffer; the pH depends only on the ratio of the acid and conjugate base present. However, concentration of the buffer is important to buffer capacity, and is considered later in this chapter. When Is a Buffer Not a Buffer? Simply having a weak acid and the salt of its conjugate base present in a solution doesn’t ensure that the buffer will act as a buffer. Buffers are most effective within ± 1 pH unit of their pKa. Outside of that range the concentration of either the acid or its salt is so low as to provide little or no capacity for pH control. Common mistakes are to select buffers without regard to the pKa of the buffer. Examples of this would be to try to use K2HPO4/KH2PO4 (pKa = 6.7) to buffer a solution at pH 4, or to use acetic acid (pKa = 4.7) to buffer near neutral pH. What Are the Criteria to Consider When Selecting a Buffer? Target pH Of primary concern is the target pH of the solution. This narrows the possible choices to those buffers with pKa values within 1 pH unit of the target pH. pH pK CH COO CH COOH = + [ ] [ ] - log 3 3 The Preparation of Buffers and Other Solutions 33

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 important. 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 solution 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 HendersonHasselbalch example above]) resulting in less than a 0.1 pH unit change, which is probably tolerable in the system. While most biomolecules can withstand the level of hydrolysis that might accompany 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 corresponding 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

essary enzyme cofactor) precipitate as Ca3(PO4)2. Incompatibili ties can also arise when a buffer component interacts with a surface. One example is the binding of amine-type buffers (i.e Tris) to a silica-based chromatography packing Biochemical Compatibility Is the buffer applied at an early stage of a research project com- patible with a downstream step? A protein isolated in a buffer containing 10mM Mg* appears innocuous, but this cation con centration could significantly affect the interaction between a reg- ulatory protein and its target DNA as monitored by band-shift assay(Hennighausen and Lubon, 1987; Band Shift Kit Instruction Manual, Amersham Pharmacia Biotech, 1994). Incompatible salts can be removed by dialysis or chromatography, but each manipu- lation adds time, cost, and usually reduces yield. Better to avoid a problem than to eliminate it downstream What Can Generate an Incorrect or Unreliable buffer Buffer Salts All buffer salts are not created equal. Care must be exercised when selecting a salt to prepare a buffer. If the protocol calls for an anhydrous salt, and the hydrated salt is used instead the buffer concentration will be too low by the fraction of water present the salt. This will reduce your buffer capacity, ionic strength, and can lead to unreliable results. Most buffer salts are anhydrous, but many are hygroscopic they will pick up water from the atmosphere from repeated opening of the container. Poorly stored anhydrous salts also will produce lower than expected buffer concentrations and reduced buffering capacity. It is always wise to record the lot number of the salts used to prepare a buffer, so the offending bottle can be tracked down if an error is suspecte If a major pH adjustment is needed to obtain the correct pH of your buffer, check that the correct buffer salts were used the ratios of the two salts werent switched, and finally verify the calculations of the proper buffer salt ratios by applying the Henderson-Hasselbalch equation. If both the acid and base com ponents of the buffer are solids, you can use the Henderson Hasselbalch equation to determine the proper mass ratios to blend and give your target pH and concentration. When this ratio is actually prepared, your ph will usually need some minor adjust ment, which should be very minor compared to the overall con- centration of the buffer The Preparation of Buffers and Other Solutions

essary enzyme cofactor) precipitate as Ca3(PO4)2. Incompatibilities can also arise when a buffer component interacts with a surface. One example is the binding of amine-type buffers (i.e., Tris) to a silica-based chromatography packing. Biochemical Compatibility Is the buffer applied at an early stage of a research project compatible with a downstream step? A protein isolated in a buffer containing 10 mM Mg2+ appears innocuous, but this cation concentration could significantly affect the interaction between a regulatory protein and its target DNA as monitored by band-shift assay (Hennighausen and Lubon, 1987; BandShift Kit Instruction Manual, Amersham Pharmacia Biotech, 1994). Incompatible salts can be removed by dialysis or chromatography, but each manipulation adds time, cost, and usually reduces yield. Better to avoid a problem than to eliminate it downstream. What Can Generate an Incorrect or Unreliable Buffer? Buffer Salts All buffer salts are not created equal. Care must be exercised when selecting a salt to prepare a buffer. If the protocol calls for an anhydrous salt, and the hydrated salt is used instead, the buffer concentration will be too low by the fraction of water present in the salt. This will reduce your buffer capacity, ionic strength, and can lead to unreliable results. Most buffer salts are anhydrous, but many are hygroscopic— they will pick up water from the atmosphere from repeated opening of the container. Poorly stored anhydrous salts also will produce lower than expected buffer concentrations and reduced buffering capacity. It is always wise to record the lot number of the salts used to prepare a buffer, so the offending bottle can be tracked down if an error is suspected. If a major pH adjustment is needed to obtain the correct pH of your buffer, check that the correct buffer salts were used, the ratios of the two salts weren’t switched, and finally verify the calculations of the proper buffer salt ratios by applying the Henderson-Hasselbalch equation. If both the acid and base components of the buffer are solids, you can use the HendersonHasselbalch equation to determine the proper mass ratios to blend and give your target pH and concentration. When this ratio is actually prepared, your pH will usually need some minor adjustment, which should be very minor compared to the overall concentration of the buffer. The Preparation of Buffers and Other Solutions 35

pH Adjustment Ionic strength differences can arise from the buffer preparation procedure. For example, when preparing a 0. 1 M acetate buffer of pH 4.2, was 0.1 mole of sodium acetate added to 900ml of water, and then titrated to pH 4.2 with acetic acid before bringing to 1L volume? If so, the acetate concentration will be significantly higher than 0. 1M. Or, was the pH overshot, necessitating the addition of dilute Naoh to bring the ph back to target, increas- ing the ionic strength due to excess sodium? The 0.1M acetate buffer might have been prepared by dissolving 0. 1 mole sodium acetate in 1 liter of water, and the ph adjusted to 4.2 with acetic acid. Under these circumstances the final acetate concentration is anyone's guess but it will be different from the first example above The best way to avoid altering the ionic concentration of a buffer is to prepare the buffer by blending the acid and conjugate base in molar proportions based on Henderson-Hasselbalch cal culations such that the pH will be very near the target pH. This solution will then require only minimal ph adjustment. Dilute to within 5% to 10% of final volume, make any final ph adjustment hen bring to volume. Generally, select a strong acid containing a counter-ion already present in the system(e.g, Cl, PO4, and OAc) to adjust a basic buffer. The strength(concentration) of the acid should be chosen so that a minimum(but easily and reproducibly delivered) volume is used to accomplish the ph adjustment. If overshooting the ph target is a problem, reduce the concentration of the acid being used. Likewise. choose a base that contains the cations already present or known to be innocuous in the assay(Na*, K, etc. Solutions of strong acids and bases used for final pH adjustment usually are stable for long periods of time, but not forever. Was the Naoh used for pH adjustment prepared during the last ice age? Was it stored properly to exclude atmospheric CO2, whose presence can slowly neutralize the base, producing sodium bicar bonate(NaHCO3) which further alters the buffer properties and ionic strength of the solution? Buffers from Stock Solutions Stock solutions can be a quick and accurate way to store " buffer precursors:” Preparing10×to100× concentrated buffer salts can simplify buffer preparation, and these concentrated solutions can also retard or prevent bacterial growth, extending almost indefi nitely the shelf stability of the solutions Pfannkoch

pH Adjustment Ionic strength differences can arise from the buffer preparation procedure. For example, when preparing a 0.1M acetate buffer of pH 4.2, was 0.1 mole of sodium acetate added to 900 ml of water, and then titrated to pH 4.2 with acetic acid before bringing to 1L volume? If so, the acetate concentration will be significantly higher than 0.1M. Or, was the pH overshot, necessitating the addition of dilute NaOH to bring the pH back to target, increasing the ionic strength due to excess sodium? The 0.1 M acetate buffer might have been prepared by dissolving 0.1 mole sodium acetate in 1 liter of water, and the pH adjusted to 4.2 with acetic acid. Under these circumstances the final acetate concentration is anyone’s guess but it will be different from the first example above. The best way to avoid altering the ionic concentration of a buffer is to prepare the buffer by blending the acid and conjugate base in molar proportions based on Henderson-Hasselbalch calculations such that the pH will be very near the target pH. This solution will then require only minimal pH adjustment. Dilute to within 5% to 10% of final volume, make any final pH adjustment, then bring to volume. Generally, select a strong acid containing a counter-ion already present in the system (e.g., Cl- , PO4 3+ , and OAc- ) to adjust a basic buffer. The strength (concentration) of the acid should be chosen so that a minimum (but easily and reproducibly delivered) volume is used to accomplish the pH adjustment. If overshooting the pH target is a problem, reduce the concentration of the acid being used. Likewise, choose a base that contains the cations already present or known to be innocuous in the assay (Na+ , K+ , etc.) Solutions of strong acids and bases used for final pH adjustment usually are stable for long periods of time, but not forever. Was the NaOH used for pH adjustment prepared during the last ice age? Was it stored properly to exclude atmospheric CO2, whose presence can slowly neutralize the base, producing sodium bicarbonate (NaHCO3) which further alters the buffer properties and ionic strength of the solution? Buffers from Stock Solutions Stock solutions can be a quick and accurate way to store “buffer precursors.” Preparing 10¥ to 100¥ concentrated buffer salts can simplify buffer preparation, and these concentrated solutions can also retard or prevent bacterial growth, extending almost indefi- nitely the shelf stability of the solutions. 36 Pfannkoch

The pH of the stock solutions should not be adjusted prior to dilution; the pH is the negative log of the H* ion concentration so dilution by definition will result in a pH change. Always adjust the ph at the final buffer concentrations unless the procedure explicitly indicates that the diluted buffer is at an acceptable p nd ionic concentation, as in the case with some hybridization and electrophoresis buffers(Gallagher, 1999) Filtration n many applications a buffer salt solution is filtered prior to mixing with the other buffer components. An inappropriate filter can alter your solution if it binds with high affinity to one of the solution components. This is usually not as problematic with polar buffer salts as it can be with cofactors, vitamins, and the like. This effect is very clearly demonstrated when a solution is prepared with low levels of riboflavin. After filtering through a PTFE filter the filter becomes bright yellow and the riboflavin disappears from the solution Incomplete Procedural Information If you ask one hundred chemists to write down how to adjust the pH of a buffer, you'll probably receive one hundred answers, d only two that you can reproduce. It is simply tedious to describe in detail exactly how buffer solutions are prepared. When ading procedures, read them with an eye for detail: Are all details of the procedure spelled out, or are important aspects left out? The poor soul who tries to follow in the footsteps of those who have gone before too often finds the footsteps lead to a cliff Recognizing the cliff before one plunges headlong over it is a learned art. A few prototypical signposts that can alert you of an impending large first step follow Which salts were used to prepare the"pH 4 acetate buffer"? Sodium or potassium? What was the final concentration? Was pH adjustment done before or after the solution was brought to final volume? If the solution was filtered, what type of filter was used? What grade of water was used? What was the ph of the starting water source What Is the Storage Lifetime of a Buffer? a stable buffer has the desired ph and buffer capacity intended when it was made. The most common causes of buffer failure are The Preparation of Buffers and Other Solutions

The pH of the stock solutions should not be adjusted prior to dilution; the pH is the negative log of the H+ ion concentration, so dilution by definition will result in a pH change. Always adjust the pH at the final buffer concentrations unless the procedure explicitly indicates that the diluted buffer is at an acceptable pH and ionic concentation, as in the case with some hybridization and electrophoresis buffers (Gallagher, 1999). Filtration In many applications a buffer salt solution is filtered prior to mixing with the other buffer components. An inappropriate filter can alter your solution if it binds with high affinity to one of the solution components. This is usually not as problematic with polar buffer salts as it can be with cofactors, vitamins, and the like. This effect is very clearly demonstrated when a solution is prepared with low levels of riboflavin. After filtering through a PTFE filter, the filter becomes bright yellow and the riboflavin disappears from the solution. Incomplete Procedural Information If you ask one hundred chemists to write down how to adjust the pH of a buffer, you’ll probably receive one hundred answers, and only two that you can reproduce. It is simply tedious to describe in detail exactly how buffer solutions are prepared.When reading procedures, read them with an eye for detail: Are all details of the procedure spelled out, or are important aspects left out? The poor soul who tries to follow in the footsteps of those who have gone before too often finds the footsteps lead to a cliff. Recognizing the cliff before one plunges headlong over it is a learned art. A few prototypical signposts that can alert you of an impending large first step follow: • Which salts were used to prepare the “pH 4 acetate buffer”? Sodium or potassium? What was the final concentration? • Was pH adjustment done before or after the solution was brought to final volume? • If the solution was filtered, what type of filter was used? • What grade of water was used? What was the pH of the starting water source? What Is the Storage Lifetime of a Buffer? A stable buffer has the desired pH and buffer capacity intended when it was made. The most common causes of buffer failure are The Preparation of Buffers and Other Solutions 37

H changes due to absorption of basic(or acidic)materials in the storage environment, and bacterial growth. Commercially pre- pared buffers should be stored in their original containers. The storage of individually prepared buffers is discussed below. The importance of adequate labeling, including preparation date. composition,pH, the preparer's name, and ideally a notebook number or other reference to the exact procedure used for the preparation, cannot be overemphasized. Absorption of bases The most common base absorbed by acidic buffers is ammonia. Most acidic buffers should be stored in glass vessels. The common indicator of buffer being neutralized by base is failure to achieve he target pH. In acidic buffers the pH would end up too high Absorption of acids Basic buffers can readily absorb CO2 from the atmosphere forming bicarbonate, resulting in neutralization of the base. This is very common with strong bases(NaoH, KOH), but often the effect will be negligible unless the system is sensitive to the pres ence of bicarbonate(as are some ion chromatography systems)or he base is very old. If high concentrations of acids(e.g, acetic neutralized by these as well. a similar common problem is improper storage of a basic solution in glass. Since silicic materi als are acidic and will be attacked and dissolved by bases, long term storage of basic buffers in glass can lead to etching of the glass and neutralization of the base Microbial contamination Buffers in the near-neutral pH range can often readily sup port microbial growth. This is particularly true for phosphate containing buffers. Common indicators of bacterial contamination are cloudiness of the solution and contamination of assays or plat Strategies for avoiding microbial contamination include steril zing buffers, manipulating them using sterile technique, refriger- ated storage, and maintaining stock solutions of sufficiently high ionic concentration. A concentration of 0.5 M works well for phos- phate buffers. For analytical chemistry procedures, phosphate buffers in target concentration ranges(typically 0. 1-0.5M) should be refrigerated and kept no more than one week. Other buffers could often be stored longer, but usually not more than two weeks

pH changes due to absorption of basic (or acidic) materials in the storage environment, and bacterial growth. Commercially prepared buffers should be stored in their original containers. The storage of individually prepared buffers is discussed below. The importance of adequate labeling, including preparation date, composition, pH, the preparer’s name, and ideally a notebook number or other reference to the exact procedure used for the preparation, cannot be overemphasized. Absorption of Bases The most common base absorbed by acidic buffers is ammonia. Most acidic buffers should be stored in glass vessels. The common indicator of buffer being neutralized by base is failure to achieve the target pH. In acidic buffers the pH would end up too high. Absorption of Acids Basic buffers can readily absorb CO2 from the atmosphere, forming bicarbonate, resulting in neutralization of the base. This is very common with strong bases (NaOH, KOH), but often the effect will be negligible unless the system is sensitive to the presence of bicarbonate (as are some ion chromatography systems) or the base is very old. If high concentrations of acids (e.g., acetic acid) are present in the local environment, basic buffers can be neutralized by these as well. A similar common problem is improper storage of a basic solution in glass. Since silicic materials are acidic and will be attacked and dissolved by bases, longterm storage of basic buffers in glass can lead to etching of the glass and neutralization of the base. Microbial Contamination Buffers in the near-neutral pH range can often readily support microbial growth. This is particularly true for phosphatecontaining buffers. Common indicators of bacterial contamination are cloudiness of the solution and contamination of assays or plates. Strategies for avoiding microbial contamination include sterilizing buffers, manipulating them using sterile technique, refrigerated storage, and maintaining stock solutions of sufficiently high ionic concentration.A concentration of 0.5M works well for phosphate buffers. For analytical chemistry procedures, phosphate buffers in target concentration ranges (typically 0.1–0.5M) should be refrigerated and kept no more than one week. Other buffers could often be stored longer, but usually not more than two weeks. 38 Pfannkoch

REAGENTS Which Grade of Reagent Does Your Experiment Require? Does your application require top-of-the-line quality, or will technical grade suffice? A good rule of thumb is that it is safer to substitute a higher grade of reagent for a lower grade, rather than vice versa. If you want to apply a lower grade reagent, test the sub- stitution against the validated grade in parallel experiments. Should You Question the Purity of Your Reagents? A certain level of paranoia and skepticism is a good thing in a New from the manufacturer The major chemical manufacturers can usually be trusted when providing reagents as labeled in new, unopened bottles. Mistakes do happen, so if a carefully controlled procedure fails, and you eliminate all other sources of error, then consider the reagents as a possible source of the problem Opened Container Here's where the fun begins. Once the bottle is opened, the manufacturer is not responsible for the purity or integrity of the chemical. The user must store the reagent properly, and use it correctly to avoid contamination, oxidation, hydration, or a host of other ills that can befall a stored reagent. How many times have you been tempted to use that reagent in the bottle with the faded label that is somewhere over 40 years old? A good rule of thumb is if the experiment is critical, use a new or nearly new bottle for which the history is known. If an experiment is easily repeated should a reagent turn out to be contaminated, then use your judg ment when considering the use of an older reagent. How can you maintain a reagent in nearly new condition? Respect the manufacturers instructions. Storage conditions (freezer, refrigerator, dessicator, inert atmosphere, etc )are often provided on the label or in the catalog. Improper handling is more kely than poor storage to lead to contamination of the reagent. It is rarely a good idea to pipette a liquid reagent directly from al bottle; this invites contamination. Instead, pour a portion into a second container from which the pipetting will be done. Solids are less likely to be contaminated by removing them directly from the bottle, but that is not always the case. It's usually use greater care handling a critical enzyme cofactor nstance,but satisfactory to transfer buffer salts from a bottle, for The Preparation of Buffers and Other Solutions

REAGENTS Which Grade of Reagent Does Your Experiment Require? Does your application require top-of-the-line quality, or will technical grade suffice? A good rule of thumb is that it is safer to substitute a higher grade of reagent for a lower grade, rather than vice versa. If you want to apply a lower grade reagent, test the substitution against the validated grade in parallel experiments. Should You Question the Purity of Your Reagents? A certain level of paranoia and skepticism is a good thing in a scientist. But where to draw the line? New from the Manufacturer The major chemical manufacturers can usually be trusted when providing reagents as labeled in new, unopened bottles. Mistakes do happen, so if a carefully controlled procedure fails, and you eliminate all other sources of error, then consider the reagents as a possible source of the problem. Opened Container Here’s where the fun begins. Once the bottle is opened, the manufacturer is not responsible for the purity or integrity of the chemical. The user must store the reagent properly, and use it correctly to avoid contamination, oxidation, hydration, or a host of other ills that can befall a stored reagent. How many times have you been tempted to use that reagent in the bottle with the faded label that is somewhere over 40 years old? A good rule of thumb is if the experiment is critical, use a new or nearly new bottle for which the history is known. If an experiment is easily repeated should a reagent turn out to be contaminated, then use your judgment when considering the use of an older reagent. How can you maintain a reagent in nearly new condition? Respect the manufacturer’s instructions. Storage conditions (freezer, refrigerator, dessicator, inert atmosphere, etc.) are often provided on the label or in the catalog. Improper handling is more likely than poor storage to lead to contamination of the reagent. It is rarely a good idea to pipette a liquid reagent directly from the original bottle; this invites contamination. Instead, pour a portion into a second container from which the pipetting will be done. Solids are less likely to be contaminated by removing them directly from the bottle, but that is not always the case. It’s usually satisfactory to transfer buffer salts from a bottle, for instance, but use greater care handling a critical enzyme cofactor. The Preparation of Buffers and Other Solutions 39

Reagents Prepared by Others Never blindly trust a reagent prepared by someone other than yourself, especially for critical assays. It’s a lot like packing your own parachute—it’s your responsibility to prepare your important solutions. If you want to trust the outcome of an important experiment to something someone else may have prepared while thinking about an upcoming vacation, it’s up to you. Prepare critical solutions yourself until you have a solid working relationship with whomever you plan to share solutions with. Even then, don’t get offended if they don’t trust your solutions! Reagents Previously Prepared by You How reliable are your solutions? Your solutions are probably fine to use if: • Your labeling and record-keeping are contemporary and accurate. • You don’t share solutions with anyone who could have mishandled and contaminated them. • Your material is within it’s expected shelf life. What Are Your Options for Storing Reagents? Storage is half the battle (handling is the other half) in keeping reagents fit for use. Follow the manufacturer’s recommendations. Shelf (Room Temperature) Solids, like buffer salts, are usually stored on the shelf in sealed bottles. Sometimes it is appropriate (e.g., for hygroscopic materials) to store them in a dessicator on a shelf. Many nonflammable liquid reagents can be also stored on a shelf. Care should be taken to store incompatible chemicals separately. For example, store acids and bases separated; store strong oxidizers away from other organics. Vented Flammables Cabinet Flammables or reagents with harmful vapors (e.g., methylene chloride) should be stored in ventilated cabinets designed for chemical storage. These cabinets are designed to minimize the chance of fire from flammable vapors; they often are designed to contain minor leaks, preventing wider contamination and possible fire. It is a good practice to use secondary spill containers (e.g., polypropylene or TeflonTM trays) in the flammables cabinet if they are not already built into the design. 40 Pfannkoch