Buffers preparation calculator

M HC1 and 1 M NaOH for use in buffer preparation. Compositions calculated from Eq. 3-8 will usually yield buffers of pH very close to those expected. 

Introduction to fast performance liquid chromatography. Review Buffer preparation, Definitions of pH, Henderson-Hasselbalch equation, and Buffer calculations. 

Be able to manipulate the Henderson-Hasselbalch equation to calculate parameters important for buffer preparations. 

The common amino acids are simply weak polyprotic acids. Calculations of pH, buffer preparation, and capacity, and so on, are done exacdy as shown in the preceding sections. Neutral amino acids (e.g., glycine, alanine, threonine) are treated as diprotic acids (Table l-l). Acidic amino acids (e.g., aspanic. acid, glutamic acid) and basic amino acids (e.g., lysine, histidine, arginine) are treated as triprotic acids, exactly as shown earlier for phosphoric acid. 

Sufficient independent equations are readily written to permit a rigorous evaluation of the hydronium ion concentration for either of these systems. Ordinarily, however, it is permissible to introduce the simplifying assumption that only one of the equilibria is important in determining the hydronium ion concentration of the solution. Thus, for a buffer prepared from H2A and NaHA, the dissociation of HA to yield is neglected, and the calculation is based only on the first dissociation. With this simplification, the hydronium ion concentration is calculated by the method described in Section 9C-1 for a simple buffer solution. As shown in Example 15-4, it is easy to check the validity of the assumption by calculating an approximate concentration of A and comparing this value with the concentrations of H2A and HA-. 

In Exercise 17.15 you calculated the pH of a buffer prepared by adding 20.0 g each of acetic acid and sodium acetate to enough water to make a 2.00-L solution, (a) Use the Calculating pH Using the Henderson-Hasselbalch... 

Calculating the pH of a buffer from given volumes of solution Given concentrations and volumes of acid and conjugate base from which a buffer is prepared, calculate the buffer pH. (EXAMPLE 17.11)... 

Suppose you need to prepare a buffer with a pH of 9.36. Using the Henderson-Hasselbalch equation, you calculate the amounts of acetic acid and sodium acetate needed and prepare the buffer. When you measure the pH, however, you find that it is 9.25. If you have been careful in your calculations and measurements, what can account for the difference between the obtained and expected pHs In this section, we will examine an important limitation to our use of equilibrium constants and learn how this limitation can be corrected. 

Aqueous solutions buffered to a pH of 5.2 and containing known total concentrations of Zn + are prepared. A solution containing ammonium pyrrolidinecarbodithioate (APCD) is added along with methyl isobutyl ketone (MIBK). The mixture is shaken briefly and then placed on a rotary shaker table for 30 min. At the end of the extraction period the aqueous and organic phases are separated and the concentration of zinc in the aqueous layer determined by atomic absorption. The concentration of zinc in the organic phase is determined by difference and the equilibrium constant for the extraction calculated. 

Prepare the solutions and measure the pH at one temperature of the kinetic study. Of course, the pH meter and electrodes must be properly calibrated against standard buffers, all solutions being thermostated at the single temperature of measurement. Carry out the rate constant determinations at three or more tempertures do not measure the pH or change the solution composition at the additional temperatures. Determine from an Arrhenius plot of log against l/T. Then calculate Eqh using Eq. (6-37) or (6-39) and the appropriate values of AH and AH as discussed above. 

Lactic acid, QH C, is a weak organic acid present in both sour milk and buttermilk. It is also a product of carbohydrate metabolism and is found in the blood after vigorous muscular activity. A buffer is prepared by dissolving lactic acid, HLac (ffa = 1.4 X 10-4), and sodium lactate, NaC3H503, NaLac. Calculate [H+] and the pH of the buffer if it is made of... 

It. A buffer is prepared by dissolving 0.0250 mol of sodium nitrite, NaN02, in 250.0 mL of 0.0410 M nitrous acid, HN02. Assume no volume change after HN02 is dissolved. Calculate the pH of this buffer. 

CI8-OOO2. Calculate the concentrations of the major species present in a buffer solution prepared by adding 1.40 g NaOH to 140. mL of 0.750 M NH4 Cl. 

For the preparation of a buffer solution of pH = 5.00, sodium acetate and acetic acid should be added to pure water in a molar ratio of 1.8 1.0. The exact amounts must be calculated using the volume and concentration ofthe solution, as Example Illustrates. 

A practical problem in solution preparation usually requires a different strategy than our standard seven-step procedure. The technician must first identify a suitable conjugate acid-base pair and decide what reagents to use. Then the concentrations must be calculated, using pH and total concentration. Finally, the technician must determine the amounts of starting materials. The technician needs a buffer at pH = 9.00. Of the buffer systems listed in Table 18-1. the combination of NH3 and NH4 has the proper pH range for the required buffer solution. 

C18-0136. A technician accidentally pours 35 mL of 12 M HCI into the I.O L of buffer solution freshly prepared as described in Problem 18.93. (a) Do a calculation to determine whether the buffer has been ruined, (b) Is it possible to bring the buffer solution back to the original pH calculated in Problem 18.93 If so, what reagent, and how much, must be added to restore the buffer ... 

Calculations. For determination of the intrinsic viscosity [ti] the prepared pectins were solved in an 0.1 M phosphate buffer with pH 6.0. The relative viscosity was determined by a glass. Ubbelhode viscometer at 25 0.1 °C. The flow time of solvent (L) was 81.8 seconds. At least six pectin solutions with different concentrations were measured in a way that their flow times (ts) comply the order 1.2to[Pg.528]

Example. A solution of ethyl acetate in pH 10.0 buffer (25°C) 1 hour after preparation was found to contain 3mg/mL. Two hours after preparation, the solution contained 2mg/mL. Calculate the... 

Aerosil was converted into the sodium form by treating it with a buffer solution (pH = 8.4) made of sodium hydroxide and sodium hydrogen carbonate solutions, after which it was filtered, washed free of alkali, and dried. This sodium-exchanged aerosil was then suspended in a solution of Ni(en)3(N03)2 prepared by adding the calculated amount of ethylene-diamine to a solution of nickel nitrate. The suspension was agitated for about 30 min and filtered off. The catalyst was then washed and dried at 100°C. 

In practice, a buffer solution (see Table 3. 4) is prepared by the partial neutralization of the selected weak acid or base with a suitable strong acid or base, or by the addition of the calculated amount of the corresponding salt. The assumption is made that the salt is completely dissociated in solution, e.g. an... 

We can prepare a buffer of almost any pH provided we know the pAa of the acid and such values are easily calculated from the Ka values in Table 6.5 and in most books of physical chemistry and Equation (6.50). We first choose a weak acid whose pKa is relatively close to the buffer pH we want. We then need to measure out accurately the volume of acid and base solutions, as dictated by Equation (6.50). 

The present chapter deals with calculations involving isotonicity, pH, and buffering of topical preparations. The discussion presented here is also relevant to the dosage forms for other routes of administration including parenteral routes. 

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