How to Prepare 500mL a 50 mM pH 6 Citrate buffer starting from citric acid as the sole source of conjugate base
I used the HH equation and got that citric acid concentration would be 0.035M, when converted to moles would be 3.362g. Then I was planning to add that amount to a beaker and add water (1/3 of what I need) and then add NaOH dropwise until I reach desired pH and then change fill the flask to 500mL mark. Am I right?
Preparing Buffer Solutions with a Specific pH Value
1. Calculate [A-]:[HA] ratio using Henderson-Hasselbalch Equation. pH = pKa + log10([A-]/[HA])
2. Ca. mole fraction of each component. Xi = ni/∑n
3. Ca. Molarity of each component. Mi = (MBuffer)(Xi)
4. Ca. moles of each component. ni = (Mi)(VBuffer)Liters
5. Ca. Stock Volume (Liters) of each component needed. (Vi)Liters = ni/Mi
6. Transfer each component into appropriate Volume Graduated Cylinder and dilute up to needed final volume.
Ratio => Xi => Mi => ni => Vi => Mix & Dilute
Example:
Using the Henderson-Hasselbalch Equation this is a two-component buffer system meaning that the weak acid and its conjugate base are added separately.
How would you prepare 10mL of a 0.01M phosphate buffer, pH 7.40, from stock solutions of 0.10M KH2PO4 and 0.25M K2HPO4? pKa of KH2PO4 = 7.20.
1. Use the Henderson Hasselbalch equation to find the ratio of A- to HA. pH = pKa + log [A- ] / [HA] 7.40 = 7.20 + log [A- ] / [HA] 0.20 = log [A- ] / [HA] 1.584893192 = [A- ] / [HA]* *Since [A- ] / [HA] = 1.584893192, we can say that [A- ] / [HA] = 1.584893192/ 1. In this case [A- ] = 1.584893192; [HA] = 1.
2. Calculate the decimal fraction (part/whole) of each buffer component. A- = 1.584893192 / (1.000 + 1.584893192) = 1.584893192 / 2.584893192= 0.61313682 HA = 1.000 / 2.584893192= 0.38686318
3. Find the molarity (M) of each component in the buffer by simply multiplying the molarity of the buffer by the decimal fraction of each component. MA- = 0.01M x 0.61313682 = 0.006131368M MHA = 0.01M x 0.38686318 = 0.003868632M
4. Calculate the moles of each component in the buffer. Moles = Molarity x Liters of buffer molesA- = 0.006131368M x 0.01L = 6.131 x 10-5 moles molesHA = 0.003868632M x 0.01L = 3.869 x 10-5 moles
5. Calculate the volume of each stock solution required to make the buffer Liters of stock = moles of the buffer component / Molarity of the stock LA- = 6.131 x 10-5 moles / 0.25 M = 2.452 x 10-4 L = 245µL LHA = 3.869 x 10-5 moles / 0.10 M = 3.869 x 10-4 L = 387µL
6. To prepare this buffer, one would use appropriately-sized pipets to measure and transfer each component to a 10mL volumetric flask and bring the solution to volume with DI-H2O.
Yes, your proposed method is almost correct. Let's go over the steps in detail to ensure the proper preparation of the 500 mL, 50 mM pH 6 Citrate buffer starting from citric acid as the sole source of conjugate base.
Step 1: Calculate the required amount of citric acid:
You correctly used the Henderson-Hasselbalch equation to determine the citric acid concentration. The equation is:
pH = pKa + log ([A-]/[HA])
pH = 6 (desired pH)
pKa = pKa of citric acid (pKa1 = 3.13, pKa2 = 4.76, pKa3 = 6.40)
To create a buffer at pH 6, we can assume that only the first dissociation (pKa1) is relevant because the pH is close to this value. So, we have:
6 = 3.13 + log([A-]/[HA])
Simplifying the equation, we get:
log([A-]/[HA]) = 2.87
To convert it to actual values, we need to know the ratio [A-]/[HA]. In this case, since citric acid is the sole source of conjugate base, [A-] = [citric acid]. Thus:
[Citric Acid] / [HA] = 10^2.87
We are given that the final concentration of the citrate buffer is 50 mM, so [Citric Acid] = 50 mM = 0.050 M.
0.050 M / [HA] = 10^2.87
Solving for [HA], we find:
[HA] = 0.050 M / 10^2.87
Step 2: Convert [HA] to grams:
Using the molar mass of citric acid (192.14 g/mol), we can convert the concentration to grams:
[HA] = 0.050 M / 10^2.87 = 1.348 x 10^-3 mol/L
Mass = moles x molar mass:
Mass of HA = 1.348 x 10^-3 mol/L x 192.14 g/mol = 0.259 g
So, you need approximately 0.259 grams of citric acid.
Step 3: Dissolve citric acid in a beaker:
Measure out 0.259 grams of citric acid on a balance and add it to a clean beaker. You may want to tare the balance with an empty container before measuring the citric acid for accuracy.
Step 4: Add water:
Add approximately 1/3 of the total volume of water you need for your final buffer volume to the beaker with citric acid. In this case, add about 166.67 mL of water (500 mL / 3 = 166.67 mL).
Step 5: Adjust pH with NaOH:
Add NaOH dropwise to the citric acid solution while continuously measuring the pH using a pH meter or pH paper. Continuously stir the solution while adding NaOH to ensure proper mixing.
Keep in mind that NaOH is a strong base and can cause harm, so handle it with care. Add the NaOH slowly and mix well after each addition until the desired pH of 6 is reached.
Step 6: Adjust final volume:
Once you have reached the desired pH of 6, transfer the solution from the beaker to a 500 mL volumetric flask and add water to the flask until the solution reaches the 500 mL mark. Mix well to ensure a homogeneous solution.
Congratulations! You have prepared a 500 mL, 50 mM pH 6 Citrate buffer starting from citric acid as the sole source of the conjugate base.