Lab Manual Quantitative Analytical Method by Deepak Pant (best historical fiction books of all time txt) π
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high value, the pH of the solution should change as shown in Fig.1.
It is easy to calculate the concentrations of HA and A- at various stages during a titration. If these concentrations and the pKa of the acid are substituted into Eq. (1), the pH at each stage also can be calculated. The results of such calculations for one titration are given in the table below.
pH CHANGES DURING THE TITRATION OF 50 cm3 OF 0.1 M HA ( Ka= 1 x 10-5) WITH 0.1 M NaOH
NaOH Added (cm3) Total Volume (cm3) pH NaOH Added (cm3) Total Volume (cm3) pH
0.00 50.0 3.0 49.95 95.95 8.00
5.00 55.0 4.05 50.00* 100.0* 8.85*
10.00 60.0 4.40 50.05 100.05 9.70
20.00 70.0 4.82 50.10 100.1 10.00
30.00 80.0 5.18 50.50 100.5 10.70
40.00 90.0 5.60 51.00 101.0 11.00
45.00 95.0 5.95 55.00 105.0 11.68
49.00 99.0 6.69 60.00 110.0 11.96
49.50 99.5 7.00 70.00 120.0 12.23
49.90 99.9 7.70
* Equivalence point
Note that the first addition of base produces a significant rise in pH. This is followed by a region in which the pH changes only slightly. In this region, the solution is buffered by the presence of both weak acid and its salt. As addition of base continues, the acid concentration drops so much that the solution is no longer buffered. Now the pH rises rapidly through the neutralization or equivalence point and slightly beyond. Beyond this region, the acid has been neutralized and the pH of the solution changes only slightly as more base is added. The rise of pH beyond the equivalence point is due to the addition of base to a relatively large volume of solution.
The rapid change in pH near the equivalence point makes a quantitative titration of acid by base a feasible experiment. Note that 2 drops of base solution near the equivalence point (from 49.95 cm3 of base to 50.05 cm3) causes a change in pH of 1.70 units in this solution. Such a change in pH is sufficient to change the color of an indicator from its acid to its basic color, so that we can easily measure to 0.05 cm3 in a 50 cm3 addition, a precision of 1 part in 1000.
The rapid change near the equivalence point is the reason titrant must be added in smaller and smaller amounts as the equivalence point is approached.
Although the data in the table were obtained by calculation using Eq. (1), similar data can be obtained by measuring the pH of the solution during any similar titration. When plotted, it gives a curve similar to that in Fig. 1.
In this experiment, you will collect such data as you titrate a weak acid with a strong base. You will use the data to determine Ka for the acid and to select a suitable indicator for titrations involving this acid.
Calculation of Ka
Because the equilibrium
is re-established after each addition of base during a titration, a value of Ka could be obtained from the data corresponding to any point before the equivalence point on the titration curve. Using the data from two particular points - the first point on the curve and that at the half-neutralization point - makes the calculations easier.
The first point on the curve is the pH of the acid solution before any base is added. Knowing the concentration of the acid in the solution and the pH, you should be able to calculate Ka.
When the acid is exactly half neutralized, the pH of the solution is equal to the pKa of the acid. Eq. (1)
shows why this is so. When half of the acid is neutralized (and present as A-) and the other half is not neutralized (and present largely as HA), the concentrations [A-] and [HA] are equal. At this point the log ([A-] / [HA]) term in Eq. (1) becomes zero.
Eq. (1) then becomes
pH1/2 = pKa .
There are several ways to identify the half-neutralization point. The easiest way is to select the point on the curve that corresponds to one half of the volume of base needed to reach the equivalence point. Ka of the acid can be obtained from the pH corresponding to this point by taking the antilog of -pH (the negative of the pH). Here is an example:
Suppose the pH at the half-equivalence point is 5.2.
Then pH1/2 = pKa = -log Ka = 5.2.
From this, Ka = 10-5.2 = 10-6 x 100.8 = 10-6 x 6.3
or, Ka = 6.3 x 10-6.
Experimental Procedure
1. Calibrate the pH meter with known standard buffers, buffered at pH 4 and 7. (When not in use, always immerse the glass membrane electrode in distilled water. Never allow it to dehydrate.)
2. Pipette 25 cm3 of acid into a 100-cm3 beaker.
3. Titrate the sample with the pH electrode immersed in the solution. This can be done by adding an increment of NaOH, waiting for equilibrium (15 to 30 seconds), then reading and recording the pH of the solution. The size of the increment to be added will change during the course of the titration.
Initially, rather large (3 to 4 cm3) increments of base can be added. Near the equivalence point, rather small increments (3 to 5 drops) will be used. For each increment, the amount of base added should be that amount which will cause a change in pH of approximately 0.2 pH units. Read and record the volume of base added and the pH of the solution. You will observe sharp increases in pH as you approach this point. After the equivalence point, continue adding base, but in 3- to 4- cm3 increments, until a pH reading of 12 is obtained.
Treatment of Data
Hand in the write-up of this experiment at the start of the next practical session. The marks are shown in brackets.
Record in the data table the volume of unknown acid used, the molarity of standard base used, and the burette reading and the pH after each addition of base.
1. Complete the data table by calculating for each increment from the burette readings the total volume of base added to that point in the titration. For example, if 3 cm3 of base were added, then 3 more and then 3 more, the numbers in the second column would be 3, 6 and 9 cm3. [1]
2. Prepare a graph with pH on the vertical axis (ordinate) and volume of base added on the horizontal axis (abscissa). [3]
3. Identify the equivalence point in the titration. This point can be approximated as the mid-point of the sharply increasing, nearly vertical, portion of the titration curve. [1]
More precisely, the equivalence point is the point at which the number of moles of base added equals the moles of acid originally present. On your titration curve, it is the point at which the slope changes from increasing to decreasing values.
4. Knowing the equivalence point, determine from your data the volume of base added to reach this point. This is the equivalence volume.
5. Using the equivalence volume and molarity of base used, calculate the concentration of acid in your unknown solution.
6. Determine Ka by the two methods described in the background section. Show your calculations. [2]
7. From your curve, select the pH color change range of indicators suitable for use in titrations of your unknown acid. From the table of indicators provided, select two or more indicators that could be used in titrating your unknown acid. 1]
Acid-Base Indicators
Color of
Name Range Acidic Form Basic Form
Methyl Violet 0.0 β 1.6 Yellow Blue
Crystal Violet 0.0 β 1.6 Yellow Blue
Malachite Green 0.2 β 1.8 Yellow Blue-green
Cresol Reda 0.4 β 1.8 Red Yellow
Thymol Bluea 1.2 β 2.8 Red Yellow
Methyl Yellow 2.8 β 4.0 Red Yellow
Bromphenol Blue 3.0 β 4.6 Yellow Purple
Methyl Orange 3.1 β 4.4 Red Yellow
Bromcresol Green 3.8 β 5.4 Yellow Blue
Methyl Red 4.4 β 6.2 Red Yellow
Chlorophenol Red 4.8 β 6.4 Yellow Red
Bromcresol Purple 5.2 β 6.8 Yellow Purple
Alizarin 5.5 β 6.8 Colorless Yellow
Bromthymol Blue 6.0 β 7.6 Yellow Blue
Phenol Red 6.6 β 8.0 Yellow Red
Neutral Red 6.8 β 8.0 Red Yellow-brown
Cresol Redb 7.2 β 8.8 Yellow Red
Cresol Purple 7.4 β 9.0 Yellow Purple
Thymol Blueb 8.0 β 9.6 Yellow Blue
Phenolphthalein 8.0 β 9.8 Colorless Pink
o-Cresolphthalein 8.2 β 9.8 Colorless Red
Thymolphthalein 9.3 β 10.5 Colorless Blue
Alizarin Yellow R 10.2 β 12.0 Yellow Violet
a The acid range of this diprotic indicator.
b The basic range of this diprotic indicator.
Experiment: Titration of An Antacid
Materials and Usage Directions
1. Materials:
Antacids Tables (various brands)
0.2 M HCl
0.2 M NaOH
methyl red indicator
two 250-mL Erlenmeyer Flasks
two 50-mL burets
2. Buret Use and Titration Technique:
Typically, a special piece of glassware is used to measure out one or both of the solutions used in a titration. It is called a buret, and quickly and accurately measures the volume of the solutions delivered. A diagram of a buret, and instructions on how to use it are given below.
Before use, the buret must be rinsed with the solution it is to contain. Close the stopcock and use a small beaker to pour about 10 mL of solution into the buret. Tip the buret sideways and rotate it until all of the inside surfaces are coated with solution. Then open the stopcock and allow the remaining solution to run out. Again close the stopcock, and pour enough solution into the buret to fill it above the "0" mark. With the buret clamped in a vertical position, open the stopcock and allow the liquid level to drop to "0" or below. Check the buret tip. It should not contain air bubbles! If it does, see your instructor. Adjust the buret so that the liquid surface is at eye level, and take the initial buret reading as shown below:
After the buret is filled, and the initial reading taken, the solution from the buret is added to the solution to be titrated (which is normally in an Erlenmeyer flask). The solution in the flask also contains an indicator, which will change color when neutralization has occurred. At first, the solution from the buret is added rapidly and the flask swirled to mix the solutions. As the endpoint is approached (or you think it is being approached), the stopcock is partially closed to slow down the addition.
If the flask is not swirled during the addition of a few drops, the portion of solution surrounding the incoming drops will change color if the endpoint is near. If this happens, slow down the solution addition, and swirl the flask gently
It is easy to calculate the concentrations of HA and A- at various stages during a titration. If these concentrations and the pKa of the acid are substituted into Eq. (1), the pH at each stage also can be calculated. The results of such calculations for one titration are given in the table below.
pH CHANGES DURING THE TITRATION OF 50 cm3 OF 0.1 M HA ( Ka= 1 x 10-5) WITH 0.1 M NaOH
NaOH Added (cm3) Total Volume (cm3) pH NaOH Added (cm3) Total Volume (cm3) pH
0.00 50.0 3.0 49.95 95.95 8.00
5.00 55.0 4.05 50.00* 100.0* 8.85*
10.00 60.0 4.40 50.05 100.05 9.70
20.00 70.0 4.82 50.10 100.1 10.00
30.00 80.0 5.18 50.50 100.5 10.70
40.00 90.0 5.60 51.00 101.0 11.00
45.00 95.0 5.95 55.00 105.0 11.68
49.00 99.0 6.69 60.00 110.0 11.96
49.50 99.5 7.00 70.00 120.0 12.23
49.90 99.9 7.70
* Equivalence point
Note that the first addition of base produces a significant rise in pH. This is followed by a region in which the pH changes only slightly. In this region, the solution is buffered by the presence of both weak acid and its salt. As addition of base continues, the acid concentration drops so much that the solution is no longer buffered. Now the pH rises rapidly through the neutralization or equivalence point and slightly beyond. Beyond this region, the acid has been neutralized and the pH of the solution changes only slightly as more base is added. The rise of pH beyond the equivalence point is due to the addition of base to a relatively large volume of solution.
The rapid change in pH near the equivalence point makes a quantitative titration of acid by base a feasible experiment. Note that 2 drops of base solution near the equivalence point (from 49.95 cm3 of base to 50.05 cm3) causes a change in pH of 1.70 units in this solution. Such a change in pH is sufficient to change the color of an indicator from its acid to its basic color, so that we can easily measure to 0.05 cm3 in a 50 cm3 addition, a precision of 1 part in 1000.
The rapid change near the equivalence point is the reason titrant must be added in smaller and smaller amounts as the equivalence point is approached.
Although the data in the table were obtained by calculation using Eq. (1), similar data can be obtained by measuring the pH of the solution during any similar titration. When plotted, it gives a curve similar to that in Fig. 1.
In this experiment, you will collect such data as you titrate a weak acid with a strong base. You will use the data to determine Ka for the acid and to select a suitable indicator for titrations involving this acid.
Calculation of Ka
Because the equilibrium
is re-established after each addition of base during a titration, a value of Ka could be obtained from the data corresponding to any point before the equivalence point on the titration curve. Using the data from two particular points - the first point on the curve and that at the half-neutralization point - makes the calculations easier.
The first point on the curve is the pH of the acid solution before any base is added. Knowing the concentration of the acid in the solution and the pH, you should be able to calculate Ka.
When the acid is exactly half neutralized, the pH of the solution is equal to the pKa of the acid. Eq. (1)
shows why this is so. When half of the acid is neutralized (and present as A-) and the other half is not neutralized (and present largely as HA), the concentrations [A-] and [HA] are equal. At this point the log ([A-] / [HA]) term in Eq. (1) becomes zero.
Eq. (1) then becomes
pH1/2 = pKa .
There are several ways to identify the half-neutralization point. The easiest way is to select the point on the curve that corresponds to one half of the volume of base needed to reach the equivalence point. Ka of the acid can be obtained from the pH corresponding to this point by taking the antilog of -pH (the negative of the pH). Here is an example:
Suppose the pH at the half-equivalence point is 5.2.
Then pH1/2 = pKa = -log Ka = 5.2.
From this, Ka = 10-5.2 = 10-6 x 100.8 = 10-6 x 6.3
or, Ka = 6.3 x 10-6.
Experimental Procedure
1. Calibrate the pH meter with known standard buffers, buffered at pH 4 and 7. (When not in use, always immerse the glass membrane electrode in distilled water. Never allow it to dehydrate.)
2. Pipette 25 cm3 of acid into a 100-cm3 beaker.
3. Titrate the sample with the pH electrode immersed in the solution. This can be done by adding an increment of NaOH, waiting for equilibrium (15 to 30 seconds), then reading and recording the pH of the solution. The size of the increment to be added will change during the course of the titration.
Initially, rather large (3 to 4 cm3) increments of base can be added. Near the equivalence point, rather small increments (3 to 5 drops) will be used. For each increment, the amount of base added should be that amount which will cause a change in pH of approximately 0.2 pH units. Read and record the volume of base added and the pH of the solution. You will observe sharp increases in pH as you approach this point. After the equivalence point, continue adding base, but in 3- to 4- cm3 increments, until a pH reading of 12 is obtained.
Treatment of Data
Hand in the write-up of this experiment at the start of the next practical session. The marks are shown in brackets.
Record in the data table the volume of unknown acid used, the molarity of standard base used, and the burette reading and the pH after each addition of base.
1. Complete the data table by calculating for each increment from the burette readings the total volume of base added to that point in the titration. For example, if 3 cm3 of base were added, then 3 more and then 3 more, the numbers in the second column would be 3, 6 and 9 cm3. [1]
2. Prepare a graph with pH on the vertical axis (ordinate) and volume of base added on the horizontal axis (abscissa). [3]
3. Identify the equivalence point in the titration. This point can be approximated as the mid-point of the sharply increasing, nearly vertical, portion of the titration curve. [1]
More precisely, the equivalence point is the point at which the number of moles of base added equals the moles of acid originally present. On your titration curve, it is the point at which the slope changes from increasing to decreasing values.
4. Knowing the equivalence point, determine from your data the volume of base added to reach this point. This is the equivalence volume.
5. Using the equivalence volume and molarity of base used, calculate the concentration of acid in your unknown solution.
6. Determine Ka by the two methods described in the background section. Show your calculations. [2]
7. From your curve, select the pH color change range of indicators suitable for use in titrations of your unknown acid. From the table of indicators provided, select two or more indicators that could be used in titrating your unknown acid. 1]
Acid-Base Indicators
Color of
Name Range Acidic Form Basic Form
Methyl Violet 0.0 β 1.6 Yellow Blue
Crystal Violet 0.0 β 1.6 Yellow Blue
Malachite Green 0.2 β 1.8 Yellow Blue-green
Cresol Reda 0.4 β 1.8 Red Yellow
Thymol Bluea 1.2 β 2.8 Red Yellow
Methyl Yellow 2.8 β 4.0 Red Yellow
Bromphenol Blue 3.0 β 4.6 Yellow Purple
Methyl Orange 3.1 β 4.4 Red Yellow
Bromcresol Green 3.8 β 5.4 Yellow Blue
Methyl Red 4.4 β 6.2 Red Yellow
Chlorophenol Red 4.8 β 6.4 Yellow Red
Bromcresol Purple 5.2 β 6.8 Yellow Purple
Alizarin 5.5 β 6.8 Colorless Yellow
Bromthymol Blue 6.0 β 7.6 Yellow Blue
Phenol Red 6.6 β 8.0 Yellow Red
Neutral Red 6.8 β 8.0 Red Yellow-brown
Cresol Redb 7.2 β 8.8 Yellow Red
Cresol Purple 7.4 β 9.0 Yellow Purple
Thymol Blueb 8.0 β 9.6 Yellow Blue
Phenolphthalein 8.0 β 9.8 Colorless Pink
o-Cresolphthalein 8.2 β 9.8 Colorless Red
Thymolphthalein 9.3 β 10.5 Colorless Blue
Alizarin Yellow R 10.2 β 12.0 Yellow Violet
a The acid range of this diprotic indicator.
b The basic range of this diprotic indicator.
Experiment: Titration of An Antacid
Materials and Usage Directions
1. Materials:
Antacids Tables (various brands)
0.2 M HCl
0.2 M NaOH
methyl red indicator
two 250-mL Erlenmeyer Flasks
two 50-mL burets
2. Buret Use and Titration Technique:
Typically, a special piece of glassware is used to measure out one or both of the solutions used in a titration. It is called a buret, and quickly and accurately measures the volume of the solutions delivered. A diagram of a buret, and instructions on how to use it are given below.
Before use, the buret must be rinsed with the solution it is to contain. Close the stopcock and use a small beaker to pour about 10 mL of solution into the buret. Tip the buret sideways and rotate it until all of the inside surfaces are coated with solution. Then open the stopcock and allow the remaining solution to run out. Again close the stopcock, and pour enough solution into the buret to fill it above the "0" mark. With the buret clamped in a vertical position, open the stopcock and allow the liquid level to drop to "0" or below. Check the buret tip. It should not contain air bubbles! If it does, see your instructor. Adjust the buret so that the liquid surface is at eye level, and take the initial buret reading as shown below:
After the buret is filled, and the initial reading taken, the solution from the buret is added to the solution to be titrated (which is normally in an Erlenmeyer flask). The solution in the flask also contains an indicator, which will change color when neutralization has occurred. At first, the solution from the buret is added rapidly and the flask swirled to mix the solutions. As the endpoint is approached (or you think it is being approached), the stopcock is partially closed to slow down the addition.
If the flask is not swirled during the addition of a few drops, the portion of solution surrounding the incoming drops will change color if the endpoint is near. If this happens, slow down the solution addition, and swirl the flask gently
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