The objective of this experiment is to determine the buffering region of histidine monohydrochloride by titrating histidine with a base, NaOH. By plotting a suitable graph, the pKa values of histidine can be observed. Normally, a titration curve is constructed to illustrate the relationship between the pH of the mixture and the number of moles of base added to it. However in this experiment, the graph of pH against the number of moles of NaOH per mole of histidine is plotted. This is to ensure that the graph is independent of the volume and concentrations of the solutions used. After determining the pKa values of histidine, the maximal buffering capacity of the histidine-NaOH mixture, as well as the effective buffering range can be determined.
Materials and Methods
To prepare 20mM solution of histidine monohydrochloride, 0.196g of histidine monohydrochloride was dissolved in 46.8mL of water, according to the calculations below:
No. of moles of histidine =
=
9.35 10-4 mol
=
46.8 mL
Upon complete mixing of the 20mM histidine monohydrochloride solution using a magnetic stirrer, 20mL of the solution was transferred into a beaker. The burette was washed with distilled water followed by NaOH and subsequently filled with 0.05M NaOH. The original pH of histidine solution was measured using the pH meter before proceeding with titration. Titration was carried out by adding NaOH to the histidine solution at 0.5mL increments. After each increment, the pH value of the resulting acid-base mixture was recorded. Titration was stopped when the acid-base mixture reached pH 11.5.
Results & Calculations
Calculation of no. of moles of histidine present in solution
=
=
Plotting graph of pH against no. of moles of NaOH per mol of histidine
Table: pH of histidine-NaOH solution with every 0.5mL of NaOH added
Determining pKa values of histidine
(i) Based on Graph 1, the two rectangles indicate the two regions where the curve approaches the point of inflection. The maximum and minimum points of the regions are marked with the yellow circle. By finding the average values of each set of maximum and minimum points, the respective pKa values can be determined.
pKa1 =
= 6.12
pKa2 =
= 9.45
(ii) pKa1 is the point where = 0.5
pKa2 is the point where = 1.5
Based on Graph 1, pKa1 and pKa2 are points marked with the red cross.
pKa1 = 6.16
pKa2 = 9.30
Maximal buffering capacity & Effective buffering range
Based on Graph 1, the acid-base mixture shows maximal buffering capacity at pH 6.12 and pH 9.45. The effective buffering range of a buffer is between ±1 of the maximal buffering capacity. Thus, the effective buffering range of histidine is pH 5.12 to pH 7.12 and pH 8.45 to pH 10.45.
If NaOH has not been accurately prepared, method used in (c)(i) will give a more reliable estimate of the pKa values.
If NaOH has not been accurately prepared, the number of moles of NaOH will be different, changing the ratio of number of moles of NaOH per mole of histidine. Method (c)(ii) depends on this ratio to determine the two pKa values. Hence, inaccurate ratios will cause the resulting pKa values to vary, leading to less reliable estimate of pKa values.
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On the other hand, method (c)(i) does not depend on the ratio between number of moles of NaOH and histidine. Thus, an inaccurate ratio will not affect the pKa values being determined. Instead, method (c)(i) relies on the point of inflection of the graph, which plots pH against the number of moles of NaOH per mole of histidine. Plotting the graph in this manner ensures that it is independent of the volume and concentrations of the solutions used. In other words, even if NaOH has been inaccurately prepared, changing the concentration of the NaOH solution, the shape of the curve remains similar. Since the shape of the curve does not change, the point of inflection will be almost at the same point. pKa values obtained by method (c)(i) will be similar to the original values when NaOH was prepared accurately.
Calculation of pH of the solution after addition of:
5mL of NaOH
No. of moles of NaOH added = Ã- 0.05 = 2.5 x 10-4 mol
NaOH ‰¡ Histidine
No. of moles of histidine reacted = 2.5 x 10-4 mol
Initial no. of moles of histidine = 4 x 10-4 mol
No. of moles of histidine left = 4 x 10-4 – 2.5 x 10-4 mol
= 1.5 x 10-4 mol
pH = pKa + log
pH = 6.12+ log
= 6.34
(ii) 12mL of NaOH
No. of moles of NaOH added = Ã- 0.05 = 6.0 x 10-4 mol
No. of moles of NaOH left = 6.0 x 10-4 – 4 x 10-4
= 2.0 x 10-4 mol
NaOH ‰¡ Histidine
No. of moles of histidine reacted = 2.0 x 10-4 mol
Initial no. of moles of histidine = 4 x 10-4 mol
No. of moles of histidine left = 4 x 10-4 – 2.0 x 10-4 mol
= 2.0 x 10-4 mol
pH = pKa + log
pH = 9.45 + log
= 9.45
(i) Three ionisable groups are present in histidine at the initial pH of the experiment. The three groups are: carboxyl group, amino group and the R group (imidazole group).
(ii) The amino group is responsible for the observed pKa value of 6.12 and the imidazole group is responsible for the pKa value of 9.45.
Structures of ionic species of histidine that participate in cellular buffering
Discussion
Histidine is an amino acid that acts as a buffer and it has three ionisable groups: carboxyl group, amino group and imidazole group. In this experiment, the focus is on the dissociation constant of the amino and imidazole group. The titration curve (as shown in Graph 1) has two ‘steps’, or two points of inflection because the amino group dissociates first followed by the dissociation of imidazole group. Hence, the amino group is responsible for the observed pKa value of 6.12 and the imidazole group is responsible for the pKa value of 9.45. Two methods were used to determine the pKa values of histidine. However these calculated values are only estimates and may deviate from the actual values due to the following experimental errors:
Parallax error occurs during the reading of the burette, resulting in inconsistent increment of NaOH added to the histidine solution. In other words, each increment of NaOH was not maintained at 0.5mL. This directly affects the precision of the experiment.
Possible solution to minimise error:
To avoid parallax error, ensure that the burette reading is taken from eye level at the bottom of the meniscus. The burette should also be placed in an upright position, perpendicular to the table. For a more precise burette reading, a black burette reading card can be placed behind the burette so as to get a clearer view, especially when colourless solutions are used.
The beaker containing the histidine-NaOH mixture is placed on the magnetic stirrer throughout the titration to ensure a homogenous mixture for more accurate pH readings. After every 0.5mL of NaOH added to the mixture, the pH of the resulting mixture is recorded by using the pH meter. However, it takes time for the pH meter to generate a final pH reading that does not fluctuate. If the pH value is recorded too quickly after the addition of NaOH, the pH reading may be inaccurate.
Possible solution to minimise error:
To obtain greater accuracy in pH reading, ensure that an appropriate waiting time (about 2min) is maintained between the addition of NaOH and the recording of pH value.
Conclusion
From this experiment, it can be concluded from the titration curve that the amino group of histidine is responsible for the observed pKa value of 6.12 and the imidazole group is responsible for the pKa value of 9.45. These two pKa values correspond to the pH at which the acid-base mixture shows maximal buffering capacity. The effective buffering range of histidine is pH 5.12 to pH 7.12 and pH 8.45 to pH 10.45.
EXPERIMENT 2: Effect of Buffer pKa on Buffering Capacity
Introduction
Buffers are solutions that are able to maintain a fairly constant pH when a small amount of acid or base is added. This experiment examines the effect of buffer’s pKa on buffering capacity by studying how well the two buffers of different pKa resist pH changes when acid or base is added. In scientific experiments, it is advisable to choose a buffer system in which the pKa of the weak acid is nearer to the pH of the interest. It will be ineffective for a buffer to resist pH changes if its pKa value is more than 1 pH unit from the pH of interest. Thus the study of the effect of pKa on buffering capacity is important in making a suitable choice of pH buffers for a specific experiment.
Materials and Methods
We study the effect of buffer’s pKa on buffering capacity by using 2 different buffers, potassium phosphate buffer and Tris-HCl, with pKa value 6.8 and 8.1 respectively. 3mL of 0.01M potassium phosphate buffer was pipetted into two test tubes, labelled A and B. 3mL of 0.01M Tris-HCl was also pipetted into two test tubes, labelled C and D. Three drops of universal pH indicator were added into each test tube, causing the solutions to turn green in colour (pH 7.0). HCl was added to test tubes A and C until the solutions turned pink (pH 4.0). KOH was added to test tubes B and D until the solutions turned purple (pH 10.0). The number of drops required for the solutions on each test tube to turn pink or purple in colour is recorded. The pH colour chart is used as it shows the colours of the solution at each pH level.
Results & Questions
Table : Number of drops of acid or base needed for buffer solution to deviate from its initial neutrality (pH 7.0)
pH Buffer
pKa of buffer
Initial pH
No. of drops of HCl required to become acidic (pH 4.0)
No. of drops of KOH required to become alkaline (pH 10.0)
0.01M potassium phosphate buffer
6.8
7.0
5
11
M Tris-HCl
8.1
7.0
2
20
Conclusions drawn from experiments
According to Table 2, potassium phosphate buffer requires five drops of HCl to reach pH 4.0, compared to Tris-HCl which requires only two drops of HCl to reach pH 4.0. This shows that potassium phosphate buffer is a more effective buffer against acids. Potassium phosphate buffer requires eleven drops of KOH to reach pH 10.0 while Tris-HCl requires twenty drops of KOH to reach pH 10.0.
Based on the results, Tris-HCl behaves as a more efficient buffer under basic conditions as it requires more amount of KOH than that of potassium phosphate to reach pH 10.0. This means that Tris-HCl has greater ability to resist increases in pH but not decreases in pH. On the other hand, potassium phosphate buffer is a more efficient buffer under acidic conditions as it requires lesser amount of HCl to reach pH 4.0. Similarly, this means that potassium phosphate buffer has greater ability to resist decreases in pH but not increases in pH.
It can be deduced that a buffer with greater pKa value is a more efficient buffer in basic conditions while a buffer with smaller pKa value is a more efficient buffer in acidic conditions.
Choosing a suitable buffer to study the properties of a phosphatase which functions optimally at pH 7.2
I would use the 0.01M Tris-HCl to study the properties of a phosphatase.
It is more appropriate to use a buffer with effective buffering range nearer to the pH of phosphatase. Tris-HCl has an effective buffering range of pH 7.1 to 9.1 while potassium phosphatase buffer has an effective buffering range of pH 5.8 to 7.8. Simply by considering the effective buffering range of the two buffers, it can be concluded that both buffers can be used to study the properties of phosphatase which functions optimally at pH 7.2.
However, considering the effective buffering range of the buffers is not sufficient to come to a sound conclusion. In this case, phosphatase is an enzyme that functions to hydrolyse phosphate groups. By adding potassium phosphate buffer to phosphatase, phosphatase will break down the phosphate group in the potassium phosphate buffer. This changes the chemical properties and hence the buffering capability of the potassium phosphate buffer.
Therefore, Tris-HCl is a more suitable buffer for the studying of phosphatase.
Discussion
In Experiment 1, the endpoint of the reactions is determined using a pH meter and construction a titration curve. However in this experiment, the endpoint is visually observed by the help of a pH colour chart. Possible sources of experimental errors arising from this method and ways to improve the experiment are discussed below:
In this experiment, only two types of buffers, Tris-HCl and potassium phosphate buffer, were used. The experiment can be improved by using more types of pH buffers to obtain more data. This will allow more accurate evaluation of the relationship between the pKa value and the buffering capacity, and thus the effect of pKa value on the buffering capacity.
Although the pH colour chart is used to compare the colours of the solutions, personal judgment comes into play when determining the colour change in the chemical reactions.
Possible solution to minimise error:
Be consistent in deciding the point of colour change and the endpoint of the experiment.
Conclusion
From this experiment, it can be concluded that a buffer with greater pKa value is a more efficient buffer in basic conditions and a buffer with smaller pKa value is a more efficient buffer in acidic conditions. Though a buffer’s pKa can affect its buffering capacity, however when choosing a suitable buffer for an experiment, we cannot simply rely on the pKa of a buffer. It is also crucial to consider the chemical properties and structure of the buffer and other reagents to be used in the experiment.
EXPERIMENT 3: Effect of Temperature on the pH of a buffer
Introduction
The aim of this experiment is to examine the effect of temperature on the pH of a buffer. This can be done by observing the changes in pH of two different buffers when temperature of the buffer solution decreases from room temperature to 4°C. pH of the buffers that are used to maintain the pH of the lab samples can change during changes in temperature due to cooling process. Changes in pH of buffers upon temperature changes can be explained by the Le Chatelier’s Principle. The study of the effect of temperature on pH of a buffer is crucial in choosing the right pH buffer that is able to show minimum changes in buffer pH, to maintain the properties of the biological samples that requires specific pH environment.
Materials and Methods
We study the effect of temperature on the pH of a buffer by using two different buffers, 0.01M potassium phosphate buffer and 0.01M Tris-HCl. 3mL of each buffer solution were pipetted into two separate test tubes. The initial pH values of the two buffers at room temperature are measured using the pH meter and recorded. Subsequently, both test tubes were placed into the ice box to cool to 4°C. After 20 minutes, the test tubes were taken out of the ice box and placed in an ice bath to maintain the temperature of the buffer solutions at 4°C. The pH of the cooled buffer solutions were measured again and recorded to obtain the results as seen in Table 3. By evaluating the pH changes (either increase or decrease) and the extent of these changes from the original pH value, we can observe the effect of temperature on the pH of a buffer.
Results & Questions
Table : The changes in the pH of the buffer solution as temperature is decreased to 4°C
Buffer
pH at room temperature
pH at 4°C
Difference in pH change (unit)
0.01M potassium phosphate buffer
7.03
7.49
0.46
0.01M Tris-HCl
7.01
8.16
1.15
Effect of temperature on the pH of Tris-HCl and potassium phosphate buffer
According to Table 3, at low temperature of 4°C, both buffer solutions become more alkaline. As temperature decreased from the room temperature to 4°C, the pH potassium phosphate buffer increased from 7.03 to 7.49, with a difference in pH change of 0.46. With the same change in temperature, the pH of Tris-HCl increased from 7.01 to 8.16, with a difference in pH change of 1.15. This shows that Tris-HCl exhibits greater changes in pH than potassium phosphate buffer, upon a given change in temperature. In conclusion, temperature has a greater effect on the pH of Tris-HCl compared to potassium phosphate buffer.
HA A» + Hº ΔH = -ve
As illustrated by the chemical equation above, the dissociation of buffers are endothermic processes. Being an endothermic process, heat is being absorbed and temperature decreases. Based on Le Chatelier’s Principle, when temperature decreases, the system will react to result in an increase in temperature. Hence, decreasing temperature to 4°C favours the backward reaction, which is an exothermic reaction that produces heat. The position of equilibrium shifts to the left, more Hº reacts with A» to form HA. Thus, the concentration of Hº decreases and causes the pH of the buffer to increase.
Discussion
Based on the experimental results, it is clear that temperature changes the pH of the buffer. Though this is not a complicated experiment, it is still subjected to experimental errors and can be improved by the following ways:
Only two types of buffers, Tris-HCl and potassium phosphate buffer, were used in this experiment. The experiment was also conducted at only one temperature. Using several buffers over a range of temperatures will allow us to observe the pH of a variety of buffers at different temperatures. In addition, both buffers used in this experiment showed an increase in alkalinity. Hence, including more variety of buffers will allow us to evaluate which type of buffer has tendency to become more alkaline or acidic with the changes in temperature.
This experiment was conducted without the use of a thermometer, hence there was uncertainty in determining the temperature of the buffer solutions. It was assumed that by placing the test tubes in the ice box for 20 minutes and then transferring into an ice bath, the buffer solutions would be maintained at 4ËšC. However, it is difficult to maintain ice baths at 4ËšC for a long period of time due to heat gain from the surroundings.
Possible solution to minimise error:
Keep a thermometer in the ice bath and consistently check the temperature of the ice bath. Add in more ice when the ice melts.
It was difficult to identify the endpoint of the experiment. Even after a long period of time (about 30 minutes), the pH reading shown on the pH meter still continued to increase slowly. Hence, stopping the experiment too early may result in an inaccurate pH reading.
Possible solution to minimise error:
Since it is difficult to identify the endpoint of the experiment, it is perhaps more logical to standardise the duration of the experiment for both buffer solutions. For example, 30 minutes for each buffer solution.
Conclusion
From this experiment, it can be concluded that a decrease in temperature will cause a change in pH of a buffer. However, the pH of the buffer does not always increase when temperature decreases. This depends on whether the dissociation process is endothermic or exothermic. In the case of an endothermic dissociation process, pH of the buffer will increase when temperature decreases. This can be explained by Le Chatelier’s Principle which states that the backward exothermic reaction will occur so as to counteract the change. Hence, the Tris-HCl and potassium phosphate buffers become more alkaline as temperature decreases.
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