How do you prepare a 1 M NaOH solution in the laboratory?

Measure 4 g of solid NaOH and add it to a 125-mL Erlenmeyer flask with 50 mL of deionized water. Stir on a stir plate until dissolved. Once homogeneous, add another 50 mL of deionized water and transfer the solution to a labeled 100-mL polyethylene bottle. Always handle NaOH with caution as it is strongly corrosive.

What steps are involved in adjusting buffer pH to a target value?

Dissolve the buffer salt in deionized water and calibrate a digital pH meter. Clamp the pH sensor in the buffer solution and allow the reading to stabilize. Slowly add 1 M NaOH or 2 M HCl with pauses between additions to let the pH stabilize. Continue until reaching the target pH, then dilute to final volume with deionized water.

How should neutral red and riboflavin-binding protein solutions be prepared and stored?

Weigh the solid into a volumetric flask and add deionized water to approximately two-thirds full. Cover with plastic paraffin film and agitate until dissolved. Remove the film and fill to the marked line. Transfer to labeled vials and store both solutions in the refrigerator to maintain stability.

What personal protective equipment is required when preparing buffer solutions?

Wear a lab coat, chemical splash goggles, and gloves when handling buffer components. Prepare all solutions in a chemical fume hood to protect against vapors and splashes. These precautions are especially important when working with corrosive substances like sodium hydroxide.

How is a monosodium phosphate buffer brought to its final volume?

After adjusting the pH to the target value, pour the buffer solution into a clean 100-mL volumetric flask. Bring the solution up to the 100-mL mark with deionized water. Seal with plastic paraffin film and invert repeatedly until the solution appears homogeneous throughout.

What equipment should be set up at each lab station for buffer experiments?

Each pair of students needs a stir plate, pH meter, handheld spectrophotometer, graduated cylinders, beakers, pipettes, cuvettes, and a magnetic stir bar. Additionally, set out calibration buffers, a lab stand with thermometer clamp, micropipettes with tips, and labeled bottles containing prepared buffers in a central location.

How should prepared buffers and chemical solutions be stored after preparation?

Store prepared buffers in labeled polyethylene bottles capped and placed in the instructor's hood. Keep neutral red and riboflavin-binding protein solutions in labeled vials in the refrigerator. Store 1 M NaOH in a corrosives cabinet. Proper storage maintains solution stability and ensures safety in the laboratory.

Buffers

Concepts Instructor Prep Student Protocol

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Buffers are solutions that work towards maintaining a constant pH in a system regardless of the addition of strong acids or bases. In the absence of buffers, the addition of a strong acid or a base to a solution changes the pH significantly. Buffers are crucial in everyday life. For example, blood acts as a buffer and is able to maintain a pH between 7.35 and 7.45 since a pH above 7.8 or below 6.8 can cause death.

So, how does a buffer work? A buffer is essentially a weak acid or base and its conjugate base or acid in equilibrium with each other. Take, for example, a buffer made of acetic acid and sodium acetate. The buffer solution contains a weak acid and its conjugate salt, which dissociates to form the conjugate base acetate. The solution is acidic because the K_a of acetic acid is higher than the K_b of the conjugate base.

Remember that the dissociation constant, K_a or K_b, defines the strength of the acid or base, respectively, as it is the equilibrium constant for the dissociation of the compound in water. The dissociation constant can also be represented as pK_a, or the negative log of K_a, where the smaller the pK_a, the stronger the acid. The same follows for pK_b, where the smaller the pK_b, the stronger the base.

Returning to the acetate buffer, if we add a strong base, like sodium hydroxide, the hydroxide ions react with the hydronium ions present in the solution. This causes more acetic acid to react with water in order to return to equilibrium, thus forming more hydronium ions. If we add a strong acid, like hydrochloric acid, more hydronium ions are formed in solution. Those ions react with the acetate ions to form more acetic acid. In both cases, there is very little change in the solution pH.

Keep in mind that each buffer has a specific pH range where it can buffer a solution. Here we show a titration curve for our acetic acid buffer. The buffering range is highlighted. This region is determined by the Henderson-Hasselbalch equation, which tells us that the pH of the buffer is determined by the K_a of the weak acid and the ratio of the conjugate base to the weak acid.

Far away from this pH, the weak acid or base of the buffer is depleted, and it cannot buffer pH. So, when selecting a buffer for your application, choose one whose pK_a is close to the pH desired.

Finally, all buffers have a limit to their buffering capacity, meaning the amount of acid or a base that can be added to the buffer solution before the pH changes significantly. For example, if we add too much acid to our acetic acid and sodium acetate buffer solution, we will protonate all of the acetate ions and accumulate a lot of hydronium ions, thereby lowering the pH.

Similarly, if we add too much base, we will deprotonate all of the acetic acid and accumulate a lot of hydroxide ions, thereby increasing the pH. Thus, the buffering capacity is influenced by the concentration of the weak acid and conjugate salt. So, a 1 molar acetate buffer has a higher buffering capacity than a 0.1 molar acetate buffer.

In this lab, you'll prepare and examine buffers with a wide variety of pH ranges. You'll then use your buffers to determine the pK_a of a pH indicator, called neutral red, and explore the change in neutral red's pK_a when riboflavin binding protein is introduced.

At the end of this lab, students should know...

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What is a buffer?

A buffer is a solution that maintains a constant pH, regardless of the addition of strong acids or strong bases as long as the amount of acid or base added is within the buffering capacity of the solution.