Why is monosodium phosphate an effective buffer near pH 7?

Monosodium phosphate is effective near pH 7 because its pKa range is 6.8 to 7.2, which overlaps with the target pH. Buffers are most effective close to their pKa, where the weak acid and conjugate base are present in significant concentrations. This allows the buffer to resist pH changes when small amounts of acid or base are added.

What is the role of NaOH when preparing a monosodium phosphate buffer?

NaOH is added drop-wise to shift the equilibrium toward the conjugate base, disodium phosphate, without altering the overall buffer composition. Unadjusted monosodium phosphate solutions typically have a pH of 4–6, so NaOH raises the pH to the target value of 7.0 while maintaining the buffer's capacity to resist pH changes.

How do the protonated and deprotonated forms of neutral red differ in color and light absorption?

The protonated form of neutral red is red and absorbs green light, while the deprotonated form is yellow-orange and absorbs blue-violet light. These distinct absorption wavelengths allow researchers to track the ionization state of neutral red across different pH values using absorption spectroscopy.

How does protein binding affect the properties of neutral red?

Protein binding alters both the absorbance and pKa of neutral red. When riboflavin-binding protein binds to neutral red, the pKa increases by approximately one unit, indicating that bound-protonated neutral red is a weaker acid than free protonated neutral red by an order of magnitude.

What is the isosbestic point in an absorption spectrum?

The isosbestic point is the wavelength where all absorption spectra cross at a single point, regardless of pH. Recording this wavelength helps verify data quality and indicates a clean transition between the protonated and deprotonated forms of neutral red across the pH range tested.

How do you determine the pKa of neutral red from absorbance data?

Plot absorbance intensity at lambda max versus pH for both free and bound neutral red. Calculate the midpoint between the highest and lowest absorbance values, then find the pH at which this midpoint occurs on the plotted line. That pH value is the pKa of the compound.

Why is it important to handle cuvettes correctly during spectrophotometry?

Cuvettes must be held by their textured sides and have transparent sides wiped clean before measurement to ensure accurate absorbance readings. The transparent sides must align with the spectrophotometer's light beam. Proper handling prevents fingerprints and contamination that would interfere with light transmission and distort results.

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.

Print Procedure Steps

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Source: Smaa Koraym at Johns Hopkins University, MD, USA

Preparing 50 mM NaH₂PO₄ Buffer, pH 7

In the first part of this experiment, you will prepare a sodium phosphate solution buffered at pH 7.0. Monosodium phosphate is a weak acid with the conjugate base, disodium phosphate. Unadjusted monosodium phosphate solutions usually have a pH of about 4 - 6.

Buffers are most effective close to their pKa, which is 6.8 to 7.2 for monosodium phosphate. So, you will use NaOH to push the equilibrium towards the conjugate base without altering the overall composition of the buffer equilibrium.

Before starting the experiment, calculate the mass of monosodium phosphate that you will need to make 200 mL of a 50 mM solution.

Table 1: Preparation of 50 mM NaH₂PO₄ Buffer, pH 7.0

Molar mass of NaH₂PO₄ = 119.98 g/mol
Solution volume (mL)	200
Molarity (mM)	50
Moles of NaH₂PO₄ (mol)
Mass needed (mg)
Initial volume of 1 M NaOH (mL)
Final volume of 1 M NaOH (mL)
Volume of NaOH used (mL)
Volume of DI water needed (mL)

Click Here to download Table 1

To begin, put on the necessary personal protective equipment, including a lab coat, splash-resistant safety glasses, and gloves. Note: This section of the lab uses NaOH, which is corrosive and toxic. Use caution when pouring and transporting NaOH.
Fill a 250-mL plastic wash bottle with deionized water.
Label two 100-mL beakers, one for neutral aqueous waste and one for basic aqueous waste.
Calibrate your pH meter using the provided buffers. Store the probe in its storage solution when you are done.
Tare weighing paper, and use a clean spatula to measure out the required amount of monosodium phosphate according to your calculations. Record the exact amount of monosodium phosphate in your lab notebook. Note: Monosodium phosphate absorbs moisture from the air, so work quickly to obtain an accurate measurement and close the container tightly when you are done with it.
Place the monosodium phosphate in a 400-mL beaker and clean the spatula with a laboratory wipe. Throw out the weighing paper and the laboratory wipe before returning to your fume hood.
Use a graduated cylinder to measure out 175 mL of deionized water. Pour the water into the beaker of monosodium phosphate and add a magnetic stir bar.
Stir the solution until the salt has dissolved completely and the solution appears homogeneous. This usually takes 2 - 3 min.
Measure 15 mL of deionized water and pour it into the beaker. Continue stirring the solution until it appears homogeneous again (about 1 – 2 min). Turn off the stir motor.
Rinse the pH probe with deionized water, and then clamp it in the solution with the sensor above the stir bar.
Bring a 10-mL graduated cylinder and watch glass to the dispensing hood and measure out 10 mL of 1 M NaOH. Note the exact volume in the graduated cylinder.
Cover your NaOH with the watch glass, and carefully bring it to your fume hood.
Resume stirring the monosodium phosphate solution. While monitoring the pH reading, use a disposable pipette to slowly add NaOH to the stirring solution in a drop-wise manner. Once the buffer pH reaches 7.0, return any NaOH still in the pipette to the graduated cylinder.
Calculate the volume of NaOH that you added to the buffer. Subtract that volume from 10 mL to determine how much deionized water you must add to your buffer to reach a total solution volume of 200 mL.
Measure the deionized water with another 10-mL graduated cylinder and add it to the stirring solution to finish making the buffer.
Rinse the pH sensor with deionized water, cover it with the storage solution filled cap, and unplug or turn off the probe.
Label a 250-mL polyethylene bottle as '50 mM monosodium phosphate buffer, pH 7.0'. Retrieve the magnetic stir bar from the solution.
Use a funnel to pour the buffer solution into the labeled bottle and cap it tightly.

Absorption Spectroscopy of Free and Bound Neutral Red
Buffers can be used to evaluate compounds at specific pH values. In this section, you will record the absorbance spectrum of the indicator neutral red in various buffers. The protonated form of neutral red is red, and the deprotonated form is yellow-orange, meaning that they absorb green and blue-violet light, respectively. Thus, the acidic and basic forms have distinct absorption wavelengths. Protein binding alters the properties of neutral red, changing both its absorbance and its pKa.

After measuring the absorbance spectrum of free neutral red at several pH values, you will add riboflavin-binding protein, or RP, to the solutions and measure the absorbances again. Absorbance intensity is related to concentration, so you will use the spectra to determine the pKa of free and bound neutral red after the lab.
- Draw the following table in your lab notebook. Number the cuvettes 1 through 10 and list the nine buffer pH values that you will use. The 10th cuvette will be a deionized water blank. Note: Always hold cuvettes by the textured sides and wipe the transparent sides just before putting the cuvette in the spectrophotometer. Remember to align the transparent sides with the beam of light in the spectrophotometer.
  Table 2: Absorbance of Free and Bound Neutral Red
  
  Cuvette # Buffer pH Abs. at λ_max(Free NRH⁺) Abs. at λ_max (Bound NRH)
  1 5.0
  2 5.5
  3 6.0
  4 6.5
  5 7.0
  6 7.5
  7 8.0
  8 8.5
  9 11
  10 blank
  Click Here to download Table 2
- Obtain ten 1.5-mL cuvettes and caps and label the caps 1 through 10 to match the table in your lab notebook.
- Label a 25-mL beaker as ‘DIH2O’ and fill it with deionized water.
- Flush your neutral aqueous waste down the drain and relabel the beaker as 'aqueous buffer waste'.
- Label a 400-mL beaker for used micropipette tips.
- Attach a tip to a 1-mL micropipette and use it to dispense 1000 µL of deionized water into cuvette 10. This is your solvent blank.
- Place 925 µL of your pH 7.0 monosodium phosphate buffer in cuvette 5.
- Bring the remaining empty cuvettes to the buffer table. Guided by the table in your lab notebook, dispense 925 µL of each buffer into the appropriate cuvette using the labeled micropipettes. Be careful not to use the same pipette tip for different buffers.
- Once you have finished, bring the buffers back to your workbench. There, set a 200-µL micropipette to dispense 75 µL, and attach a tip to the micropipette.
- Dispense 75 µL of neutral red into each of the nine buffer cuvettes. Invert each cuvette several times to thoroughly mix the solutions. Note: Replace the pipette tip if it touches a buffer.
- Once you have mixed neutral red with every buffer, take a picture, or write down the colors of the solutions.
- Turn on a hand-held spectrophotometer and allow the light source to warm up. Once it is ready, create a new experiment to measure absorbance versus wavelength for the free neutral red.
- Insert the cuvette of deionized water and acquire a spectrum of the deionized water. Set it as a solvent background or blank.
- Then, remove the blank and insert the cuvette of neutral red in the lowest pH buffer, which will have the highest concentration of protonated neutral red.
- Acquire a spectrum, which should show one strong peak. Identify the wavelength corresponding to the highest point of this peak, or the maximum. Record this wavelength in your lab notebook as lambda max for protonated free neutral red.
- Save the data and remove the cuvette. Record the absorbance in your notebook, then collect spectra for cuvettes 2 - 9 using the same procedure.
- Record the wavelength of the isosbestic point in your lab notebook. That is where all the spectra cross at a single wavelength. If a spectrum does not cross at that point, empty and clean the cuvette, prepare a fresh sample, and try again.
- Once you finish collecting spectra for all nine cuvettes, save and export the data.
- Create a new experiment to measure absorbance versus wavelength for the bound neutral red.
- Fit a new tip to the 200-µL micropipette and add 75 µL of riboflavin-binding protein solution to each cuvette, including the blank. Eject the tip, secure the cuvette caps, and invert the cuvette several times to mix the solutions.
- Insert cuvette 10 into the spectrophotometer and set it as the solvent blank.
- Acquire a spectrum of the lowest pH sample, and identify the wavelength corresponding to the maximum of the most intense peak. Record it in your lab notebook as the lambda max of protonated-bound neutral red.
- Acquire spectra for cuvettes 2 - 9 in the same way as you did for the free neutral red samples. Record the intensities at lambda max for protonated-bound neutral red in your table, along with the wavelength at the isosbestic point.
- When you are done, save your data, export it for later analysis, and shut down the spectrophotometer.
- Put away all equipment and dispose of your used pipette tips in an approved waste container or trash bin.
- Empty the cuvettes into the aqueous buffer waste beaker and rinse the cuvettes into the beaker with deionized water.
- Empty the excess NaOH into the basic waste and rinse the graduated cylinder with water.
- Flush the basic aqueous waste down the drain with copious tap water, along with the other aqueous waste.
- Wash your glassware, cuvette caps, and stir bar per your lab standard procedures. Clean your work surfaces with a damp paper towel and throw out used paper towels and lab wipes in the lab trash.
Results
Now, let's analyze our absorbance data to determine the pKa's of neutral red.

Table 3: Determination of Neutral Red pKa's

Free NRH⁺ Bound NRH⁺
λ_max (nm)
ΔA (nm)
Midpoint (nm)
pKa

Click Here to download Table 3
- Plot both sets of intensity data with absorbance intensity at lambda max on the y-axis and pH on the x-axis, with the points joined by smooth lines.
- Calculate the difference between the starting and ending absorbance intensities for each data series.
- Determine the absorbance midway between the starting and ending absorbances for each series, or the absorbance midpoints.
- Find where the midpoint occurs on each line and identify the pH value at that point. These pH values are the pKa's of free and bound neutral red.
- Here, we see an increase in pKa of about one between free and bound neutral red, indicating that bound-protonated neutral red is a weaker acid than free protonated neutral red by an order of magnitude.