Buffers Preparation

Did you know that baby lotion contains a solution that acts as a buffer to keep the pH of the cream around 6, thus preventing bacteria from multiplying? Buffers play significant roles in our lives. Some of them are even responsible for maintaining our blood pH at 7.4.

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Table of contents

So, if you are interested in learning more about buffers, let's dive into the world of preparing buffers!

• This article is about buffers preparation.
• First, we will talk about how to prepare a buffer and also look at titration.
• Then, we will learn the procedure involving buffer preparation.
• After, we will look at buffer preparation and standardization.
• Lastly, we will learn about the preparation of neutral buffered formalin.

Definition of a Buffer

First, let's define what a buffer solution is.

A buffer is a solution that is capable of resisting changes in pH when small amounts of acid or bases are added to it.

A buffer solution usually consists of a mixture of a weak acid and its salt, or a weak base and its salt. Acidic buffers are those have a pH that is less than 7, whereas a basic buffer has a pH that surpasses 7.

Buffers work by removing any hydrogen (H+) or hydroxide (OH-) ions that are added to it, preventing the pH from changing! But, how exactly do buffers accomplish this? Let's take a look.

Let's say that you have a CH3COOH/CH3COO- buffer solution. Acetic acid is a weak acid, so this is a type of acidic buffer.

$$CH_{3}COOH_{(aq)}+H_{2}O_{(l)}\rightleftharpoons CH_{3}COO^{-}+H^{+}_{(aq)}$$

If you add a small amount of strong acid to this buffer, the acetate ions (CH3COO-) in the buffer will combine with the recently added hydrogen ions (H+) and create more acetic acid. When this happens, the buffer solution removes most of the new hydrogen ions, preventing the pH from undergoing significant changes.

Now, you add a small amount of a strong base to this buffer, the buffer will remove the majority of the added OH- ions by reacting with H+ to form water!

Basic buffers work a similar way. A common example of a basic buffer is a NH3/NH4+ buffer solution.

$$NH_{(3\ aq)}+H_{2}O_{(l)}\rightleftharpoons NH_{(4\ aq)}^{+}+OH^{-}_{(aq)}$$

If you add a small amount of strong acid, the hydroxide ions (OH-) in the buffer will combine with the recently added hydrogen ions (H+) to form water. If you add a small amount of a strong base to this buffer, the buffer will remove the majority of the added OH- ions by reacting NH4+ to form NH3.

pH, and pKa

pH is defined as a measurement of the [H+] ion concentration in a solution. When dealing with buffers, we can use the Henderson-Hasselbalch equation to calculate pH.

$$pH=pKa+log_{10}\frac{[A^{-}]}{[HA]}$$

where, the concentration of the anionic (negatively charged) species is, [A-], the concentration of the acid is, [HA], and the negative logarithm of acid dissociation constant is, pKa. Let's look at an example!

Calculate the pH of a buffer containing made of 0.10 M HA and 0.075 M NaA. (Ka = 2.5 x 10-6); where, an acid is symbolized by, HA, the salt of the acid is symbolized by, NaA, and the acid dissociation constant is, Ka:

First, we need to use the Ka value to calculate pKa.

$$pKa=-log_{10}(2.5\cdot 10^{-6})=5.6$$

Now, we can use the formula above to calculate the pH of the buffer solution.$$pH=5.60+log_{10}\frac{[0.075]}{0.10}$$$$pH=5.48$$

pKa is the negative log of the acid dissociation constant, known as Ka. This acid dissociation constant is used by chemists to determine the strength of an acid, as it measures its ability to dissociate in water. The higher the Ka of an acid, the stronger it will be!

$$Ka=\frac{[products]}{[reactants]}=\frac{[H^{+}]\cdot [A^{-}]}{[HA]}$$

$$pKa=-logKa$$

To learn more about this, check out “Henderson-Hasselbalch Equation” and “pH and pKa

Buffer Preparation and Titration

Now that we know how buffers work, let's look at buffer preparation. There are three ways to prepare a buffer solution:

 Buffer Solution Weak Acid / Conjugate Base Weak base/Strong acid Weak Acid / Strong Base

The first way is by mixing an acid with its conjugate base while keeping the concentrations of both components equal. For example, a buffer containing 0.10 M HClO and 0.10 M NaClO (aq).

The second way involves mixing a weak base with a strong acid. This buffer should have a higher amount of the weak species compared to the strong species. For example, a buffer made up of 1.0 M HCl mixed with 1.25 M NH3.

Another way to prepare a buffer solution is by mixing a weak acid with a strong base. Similarly, this buffer should also have a higher amount of the weak species compared to the strong species. For example, a buffer containing 1.50 moles of the weak oxyacid HNO2 and 1.25 moles of the strong acid NaOH.

When deciding on a buffer to use, the best choice is to prepare a buffer that has the pKa closest to the desired pH. Let's put this into practice.

You need a buffer solution with a pH of around 5. Which of the following combinations would you choose to prepare the buffer?

1. HIO3 + KIO3
2. HCN + KCN
3. HNO2 + KNO2
4. HIO4 + NaIO4

Use the Ka values below.

Acid dissociation constant for different acids, Isadora Santos - StudySmarter Originals.

Although this seems like a complicated problem, it is actually simple to solve! We said before that the best buffer would be the one with a pKa close to the desired pH value (in this case, a pH of 5). So, let's go ahead and use the Ka values to calculate pKa.

$$pKa=-log_{10}Ka$$

$$pKa\ of\ HIO_{3}=-log_{10}(1.7\cdot 10^{-1})=0.77$$

$$pKa\ of\ HIO_{4}=-log_{10}(2.8\cdot 10^{-2})=1.55$$

$$pKa\ of\ HNO_{2}=-log_{10}(4.5\cdot 10^{-4})=3.35$$

$$pKa\ of\ HCN=-log_{10}(6.2\cdot 10^{-10})=9.21$$

In this case, the pKa closest to pH 5 is the pKa of HNO2. Therefore, we would choose to prepare a buffer containing HNO2 + KNO2.

Buffer solutions are also seen in titrations. During a titration, a solution of known concentration is added to the solution of unknown concentration until the endpoint of the titration is reached. You can learn more about titrations by reading "Types of Acid-Base Titrations".

Titration is a process used by chemistry to determine the unknown concentration of a particular solution by combining it to a solution of known concentration.

As an example, let's look at the titration of 50.0 mL of a 0.1 M acetic acid, CH3COOH (aq) with 0.1 M NaOH (aq). Notice that this is an example of a titration of a weak acid with a strong base.

1. At the start of the titration, the solution only contains CH3COOH (aq) and has a pH of 2.89.
2. Between the initial pH and the equivalence point, NaOH is added dropwise to the solution, and the added hydroxide ions (OH-) convert, CH3COOH (aq) into CH3COO- (aq), forming a buffer solution!
3. At the equivalence point, the acid gets completely neutralized by the added base (moles of acid = moles of base), resulting in the formation of the salt solution CH3COONa (aq).
4. After the equivalence point, there is no more acid left to react with excess OH-.

Buffer Preparation Procedures

Now, let's look at the procedure to prepare a buffer solution of known pH.

Step 1. The first step in the preparation of a buffer solution is to choose the conjugate acid-base pair that best works for your needs. For example, if you needed to prepare a buffer solution of pH 6.3, you would choose an acid with a pKa that is closest to that of the desired pH. So, we do this by calculating Ka and then finding an acid with a Ka value close to the calculated Ka, and a pKa value close to 6.3.

$$pKa=-logKa$$

$$Ka=10^{-pKa}$$

$$Ka=10^{-6.3}=5.01\cdot 10^{-7}$$

Carbonic acid (H2CO3) has a Ka value of 4.4 x 10-7 and a pKa1 of 6.37, so it would be a great choice of weak acid for our buffer. To supply the hydrogen carbonate ions, we can choose a salt such as sodium hydrogen carbonate (NaHCO3).

Step 2. After choosing the components for our buffer solution (H2CO3/NaHCO3), we need to calculate the ratio of buffer components by rearranging the Henderson-Hasselbalch equation.

$$pH=pKa+log\frac{[HCO_{3^{-}}]}{H_{2}CO_{3}}$$

$$6.30=6.37+log\frac{[HCO_{3^{-}}]}{H_{2}CO_{3}}$$

$$6.30-6.37=log\frac{[HCO_{3^{-}}]}{H_{2}CO_{3}}$$

$$10^{-0.07}=log\frac{[HCO_{3^{-}}]}{H_{2}CO_{3}}$$

$$\frac{[HCO_{3^{-}}]}{H_{2}CO_{3}}=\frac{0.85}{1}$$

So, for every 1 mole of H2CO3 in 1 L of solution, we will need 0.85 moles of NaHCO3.

Step 3. Now that have the ratio of components, we can use it to calculate the buffer solution concentration. Let's say that in your laboratory you have a 0.5 M H2CO3 in stock. We can use it to calculate the amount of NaHCO3 required to mix with the weak acid to create a buffer with a pH of 6.3.

First, calculate the moles of carbonic acid.

$$Mol\ H_{2}CO_{3}=1.0L\cdot \frac{0.5mol\ H_{2}CO_{3}}{1.0L}=0.5\ moles\ H_{2}CO_{3}$$

Then, use the buffer ratio from step 2 to find moles of NaHCO3 from the moles of H2CO3.

$$Mol\ NaHCO_{3}=0.5\ moles\ H_{2}CO_{3}\cdot \frac{0.85\ mol\ NaHCO_{3}}{1.0\ mol\ H_{2}CO_{3}}=0.43NaHCO_{3}$$

After, convert moles of NaHCO3 to mass (in grams).

$$g\ NaHCO_{3}=0.43\ mol\ NaHCO_{3}\cdot \frac{84.006g\ NaHCO_{3}}{1.0mol\ NaHCO_{3}}=36.12g\ NaHCO_{3}$$

Now, we know exactly how much NaHCO3 we need to add to our H2CO3 to make a 1-L buffer solution!

Step 4. The final step in the buffer preparation procedure is to mix the components! First, we have to weigh 36.12g NaHCO3 and dissolve it in enough 0.5 M H2CO3. After it dissolves, we have to add more 0.5 M H2CO3 solution until we get a volume of 1.0 L.

To make sure that the pH is correct, we can use a pH meter. If the pH is slightly higher, add a couple of drops of a strong acid until the desired pH is reached. If the pH is a bit lower than expected, add a few drops of a strong base until the desired pH is reached.

Congratulations! You just made an H2CO3/HCO3- buffer solution with a pH of 6.3!

Buffer preparation and Standardization

Let's talk about standardization. Standard buffer solutions are buffer solutions of standard pH. Standard buffer solutions have different pH values and can be prepared by combining different solutions. For example, we can make a hydrochloric acid buffer with a pH of 1.2 by mixing 50mL of 0.2M KCl with 85.0 mL of 0.2M HCl in a 200mL volumetric flask, and then add water to volume!

We can use standard buffer solutions to calibrate a pH meter. For example, during the calibration of a pH meter using standardized buffers with pH 7 and pH 2, the pH meter electrode is placed in the pH 7 buffer solutions and then allowed to calibrate. The same thing is done with the pH 2 buffer after!

Neutral Buffered Formalin Preparation

Have you heard of formalin before? Formalin is a solution that is used as a fixative for the preservation of biological specimens. At low concentrations, formalin is bacteriostatic, meaning that it inhibits the growth of microorganisms, but does not kill them.

One way of preparing a 10% Neutral Buffered Formalin from stock solutions is by mixing 100 .0mL of a 37% formaldehyde stock solution, 900.0 mL of distilled water, 4.0 g of NaH2PO4 , and 6.5 g of Na2HPO4.

Now, I hope that you feel more confident when it comes to preparing buffers!

Preparing Buffers - Key takeaways

• A buffer is a solution that is capable of resisting changes in pH when small amounts of acid or bases are added to it.
• Buffers work by removing any hydrogen (H+) or hydroxide (OH-) ions that are added to it, preventing the pH from changing.
• When deciding on a buffer to use, the best choice is to prepare a buffer that has the pKa closest to the desired pH.
• There are three ways to prepare a buffer solution: 1) Weak Acid / Conjugate Base, 2) Weak base/Strong acid, and 3) Weak Acid / Strong Base.

References

1. “Inquiry Labs for AP® Chemistry: Buffers in Household Products.” FlinnPREP, www.flinnprep.com/Chapter?courseFriendlyUrl=Chemistry_Inquiry_Lab&unitFriendlyUrl=Buffers_in_Household_Products&chapterTitle=Buffers_in_Household_Products. ‌
2. “Buffers - Analytical Chemistry Video.” Clutch Prep, www.clutchprep.com/analytical-chemistry/buffers. ‌
3. “How Do You Prepare a Buffer Solution of Known PH from Scratch?” Master Concepts in Chemistry, masterconceptsinchemistry.com/index.php/2018/11/01/how-do-you-prepare-a-buffer-solution-of-known-ph-from-scratch/. ‌
4. “Preparation of Buffer Solutions : Pharmaceutical Guidelines.” Pharmaguideline, www.pharmaguideline.com/2010/09/preparation-of-buffer-solutions.html. ‌
5. “Calibrating the PH Meter.” Purdue.edu, 2019, chemed.chem.purdue.edu/genchem/lab/equipment/phmeter/use.html. ‌
6. Rosenthal, Patrick R. Medical Microbiology. S.L., Elsevier - Health Science, 2020. ‌
7. “Chemistry of Formalin Fixation · QED Bioscience Inc.” QED Bioscience Inc, 14 June 2018, www.qedbio.com/blog/chemistry-of-formalin-fixation/. ‌
8. Surgical Pathology Staining Manual - Formalin Fixation. Webpath.med.utah.edu. ‌
9. Theodore Lawrence Brown, et al. Chemistry : The Central Science. 14th ed., Harlow, Pearson, 2018. ‌
10. Moore, John T, and Richard Langley. McGraw Hill : AP Chemistry, 2022. New York, Mcgraw-Hill Education, 2021. ‌

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Frequently Asked Questions about Buffers Preparation

How are buffers prepared?

There are three ways to prepare a buffer solution.

1. By mixing an acid with its conjugate base.
2. By mixing a weak base with a strong acid.
3. By mixing a weak acid with a strong base

How do you write a buffer equation?

The buffer equation of a weak acid/conjugate base buffer solution is written by adding the weak acid in the reactant base, and the conjugate base plus hydronium ions on the product side.

How do you prepare a buffer solution for pH meter calibration?

Buffer solutions used for pH meter calibrations are standard solutions, meaning that they have a standard pH.

We can use standardized buffers with pH 7 and pH 2 to calibrate a pH meter. In this process, the pH meter electrode is placed in the pH 7 buffer solution and then allowed to calibrate. The same thing is done with the pH 2 buffer after.

What is the first step for preparation of buffer solution?

The first step in the preparation of a buffer solution is to choose the conjugate acid-base pair that better works for your needs. As a general sure, you should choose an acid with a pKa that is closest to that of the desired pH.

How do you prepare a buffer with a specific pH?

There are 4 steps involved in the preparation of a buffer solution with a specific pH.

1. Choose an acid with a pKa that is closest to that of the desired pH.
2. Calculate the ratio of buffer components using the Henderson-Hasselbalch equation.
3. To calculate the buffer solution concentration.
4.  Mix the components to create your buffer!

Test your knowledge with multiple choice flashcards

What is the correct definition of a Buffer?

_____ buffers are those that have a pH of less than 7.

_____ buffers are buffers that have a pH greater than 7.

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