Alkalinity*

*Source: Water Quality with CBL^TM

INTRODUCTION

Alkalinity should not be confused with pH. The pH of a solution is a measure of the concentration of acid, or H⁺ ions, in the water. Alkalinity is a measure of the water’s capacity to neutralize an acid, or H⁺ ions, thereby keeping the pH at a fairly constant level.

The alkalinity of surface water is primarily due to the presence of hydroxide, OH^–, carbonate, CO₃^2–, and bicarbonate, HCO₃^–, ions. These ions react with H⁺ ions by means of the following chemical reactions:

OH^– + H⁺ H₂O

CO₃^2– + H⁺ HCO₃^–

HCO₃^– + H⁺ CO₂ + H₂O

Expected Levels

Summary of Method

Alkalinity is measured by titrating a water sample with sulfuric acid. The Vernier pH System is used to monitor pH during the titration. The equivalence point will be at a pH of approximately 4.5, but will vary slightly, depending on the chemical composition of the water. The volume of sulfuric acid added at the equivalence point of the titration is then used to calculate the alkalinity of the water.

Materials Checklist

Collection and Storage of Samples

1. This test should be conducted in the lab. Collect at least 300 mL of sample water so that several 100-mL trials could be run, if needed.

2. It is important to obtain the water sample from below the surface of the water and as far away from the shore as is safe. If suitable areas of the stream appear to be unreachable, samplers consisting of a rod and container can be constructed for collection. Refer to page Intro-2 of the Introduction of this book for more details.

3. If the testing cannot be conducted within a few hours, place the samples in an ice chest or a refrigerator.

Testing Procedure

1. Obtain and wear goggles.

2. Using a 100-mL graduated cylinder, carefully add 100.0 mL of sample water to a clean 250-mL beaker. Note: The sample water should be near room temperature when this test is conducted.

3. Place the beaker on a magnetic stirrer and add a magnetic stirring bar. Set the stirrer to a speed that mixes the sample well, but does not splash. If no magnetic stirrer is available, you will need to stir the solution with a stirring rod during the titration.

4. Prepare the pH System for data collection.

Plug the pH Amplifier into the adapter cable in Channel 1 of the CBL System. Connect the pH Electrode to the pH Amplifier.

Use the link cable to connect the CBL System to the TI Graphing Calculator. Firmly press in the cable ends.

5. Obtain a clean buret and rinse it with a few mL of the 0.0100 M H₂SO₄ solution. CAUTION: Sulfuric acid, H₂SO₄, is corrosive. Avoid spilling it on your skin or clothing. Dispose of the rinse solution as directed by your teacher. Use a utility clamp to attach the buret to the ring stand, as shown. Fill the buret a little above the 0-mL level with the H₂SO₄ solution. Drain a small amount of the solution so it fills the buret tip and leaves the H₂SO₄ solution at the 0-mL level.

6. Turn on the CBL unit and the calculator. Start the CHEMBIO program and proceed to the MAIN MENU.

7. Set up the calculator and CBL for pH measurement.

Select SET UP PROBES from the MAIN MENU. Enter "1" as the number of probes. Select PH from the SELECT PROBE menu. Enter "1" as the channel number. Select USE STORED from the calibration menu.

8. Remove the pH electrode from the storage bottle, rinse the tip with distilled water from the wash bottle. Use the second beaker to catch the rinse water. Clamp the electrode in place and position it in the sample water so that it is not struck by the stirring bar.

9. Set up the calculator and CBL for data collection.

Select COLLECT DATA from the MAIN MENU.

Select TRIGGER/PROMPT from the DATA COLLECTION menu.

Allow the system to warm up for 30 seconds, then press .

10. Monitor the pH value on the CBL screen. Once it has stabilized, press on the CBL and enter "0" (the buret volume in mL) on the calculator. You have just saved the first data pair for this experiment.

11. You are now ready to perform the titration. This process goes faster if one person manipulates and reads the buret while another person operates the calculator and enters volumes.

Select MORE DATA to collect another data pair. Add a small quantity of H₂SO₄ titrant (enough to lower the pH about 0.2 pH units). When the pH stabilizes, press and enter the current buret reading (to the nearest 0.01 mL). You have now saved the second data pair for the experiment.

Select MORE DATA. Continue adding H₂SO₄ solution in increments that lower the pH by about 0.2 pH units and enter the buret reading after each increment. When the graph shows the pH value beginning to drop more quickly (at approximately pH 5.5), change to one-drop increments. Enter a new buret reading after each addition. Note: It is important that all additions of acid in this part of the titration be exactly one drop in size.

When the pH values start to flatten out (approximately pH 4), again add larger increments that lower the pH by about 0.2 pH units, and enter the buret level after each increment.

Continue for two or three more additions, or until the graph clearly shows that the pH has leveled off again.

12. Select STOP AND GRAPH from the DATA COLLECTION menu when you have finished collecting data.

13. Dispose of the beaker contents as directed by your teacher. Rinse the pH electrode with distilled water from the wash bottle. Use the second beaker to catch the rinse water. Return the electrode to the storage solution bottle and tighten the cap.

Calculations

1. Determine the volume of H₂SO₄ added at the equivalence point of the titration. The equivalence point is the point where the titration curve makes the steepest drop in pH.

Examine the data points along the displayed graph of pH vs. H2SO4 volume. As you move the cursor right or left, the volume (X) and pH (Y) are displayed below the graph. To determine the equivalence point, go to the region of the graph with the steepest drop in pH. In the screen shot at the right, the cursor is shown at the equivalence point.

Find the H₂SO₄ volume just before this jump.

Find the H₂SO₄ volume just after this jump.

Calculate the average of these points by adding them together and dividing by two. Record this number, which represents the exact volume of H₂SO₄ added at the equivalence point, on the Data & Calculations sheet (round to the nearest 0.1 mL).

2. Calculate the alkalinity of the sample by multiplying the volume of H₂SO₄ added at the equivalence point by a conversion factor of 10.0. Record this value on the Data & Calculations sheet (round to the nearest 1 mg/L CaCO₃).

Optional Calculations

3. Calculate the moles of H₂SO₄ used to reach the equivalence point.

4. The reaction occurring in this titration is

H₂SO₄ + CaCO₃ H₂O + CO₂ + CaSO₄

Based on the mole ratio of H₂SO₄ to CaCO₃, calculate the moles of CaCO₃ reacted at the equivalence point.

5. Calculate the mass in grams of CaCO₃ in the sample. Convert to milligrams.

6. Calculate alkalinity in mg/L CaCO₃. Compare this value to your answer in Problem 2.

DATA & CALCULATIONS

Alkalinity

 Column Procedure:

1.  Record the volume of H₂SO₄ at the equivalence point.

2. Multiply Column A by 10.0 as described in the procedure to obtain alkalinity (B = A 5 10.0).

ADDITIONAL INFORMATION

Tips for Instructors

1. The 0.0100 M H₂SO₄ solution can be prepared from concentrated sulfuric acid by following these steps:

Prepare 1 liter of 0.500 M H₂SO₄ by adding several hundred milliliters of distilled water to a 1-L volumetric flask. Carefully add 28.0 mL of concentrated H₂SO₄.

HAZARD ALERT: Concentrated sulfuric acid, H₂SO₄, is very corrosive. Avoid spilling it on your skin or clothing. Remember to always add acid to water—not the other way around. This solution will get very warm.

Bring the level of solution up to the 1-L mark with distilled water. Mix well.

Prepare 1 liter of 0.0100 M H₂SO₄ by adding several hundred milliliters of distilled water to a 1-L volumetric flask. Carefully add 20.0 mL of 0.500 M H₂SO₄.

Bring the level of solution up to the 1-L mark with distilled water. Mix well.

Good results can be obtained by very carefully preparing the 0.0100 M H₂SO₄ solution as described above. For greatest accuracy, however, the acid should be standardized against a primary standard of sodium carbonate solution, Na₂CO₃. To prepare this solution, dry 3 to 5 g Na₂CO₃ at 250°C for 4 hours and cool in a dessicator. Weigh 2.5 ± 0.2 g (to the nearest mg), transfer to a 1-L volumetric flask, fill the flask to the mark with distilled water and mix. Calculate the molarity of this solution using four significant figures. Do not keep more than one week. Titrate the H₂SO₄ solution with the standard Na₂CO₃ to determine the precise molarity of the H₂SO₄.

2. The procedures given in this test determine total alkalinity. There are three kinds of alkalinity referred to in environmental literature—phenolphthalein alkalinity, methyl orange alkalinity, and total alkalinity.

Phenolphthalein alkalinity measures the buffering action of bases as strong as, or stronger than, the carbonate ion. If the sample water has a pH level higher than approximately 8.3, you will see two equivalence points in the titration curve instead of just the one around 4.5. The first drop in pH at around 8.3 is the phenolphthalein equivalence point, and the amount of acid used to titrate to this point is used to calculate the phenolphthalein alkalinity. If the sample water is initially below pH 8.3, the phenolphthalein alkalinity is zero. Sample reactions that occur during the titration include

OH^– + H⁺ H₂O

CO₃^2– + H⁺ HCO₃^–

Methyl orange alkalinity measures the buffering action of bases as strong as, or stronger than, the bicarbonate ion. This is calculated using the volume of acid needed to titrate to the lower pH equivalence point of 4.5. In addition to the reactions above, the following reaction also occurs

HCO₃^– + H⁺ CO₂ + H₂O

Total alkalinity is the sum of the phenolphthalein alkalinity and the methyl orange alkalinity.

3. The ions discussed above are the major contributors to alkalinity, but there can also be some minor contributions from borates, silicates, and phosphates. These contributions are so small, however, that they are usually negligible.

4. An alternate way of determining the precise equivalence point of the titration is to take the second derivative of the pH-volume data. The CHEMBIO data-collection program contains a sub-program, CMBDERIV, that allows you to view first and second derivative plots of pH-volume data. The CMBDERIV program is set up to analyze volume data in L1 and pH data in L2. To run the program, follow this procedure:

After you have collected pH-volume data, leave the CHEMBIO program by selecting QUIT from the MAIN MENU.

Start the CMBDERIV program and proceed to the GRAPHS menu. Select SECOND DERIV, a plot of @ ²pH/@ vol².

Using the arrow keys, move the cursor to the equivalence point. This will be the point where the curve crosses the zero line.

The x-value shown at the bottom of the screen is the volume of acid at the equivalence point.

5. The second derivative of the pH-volume data could also be calculated using the Graphical Analysis software program after the data has been transferred from the calculator to the computer. A sample of this curve along with the original data curve is shown at the right.

6. It can be interesting to use a pH indicator in addition to the pH System when conducting this titration. Phenolphthalein changes from pink above pH 8.3 to colorless below 8.3. Methyl orange changes from yellow above pH 4.5 to red below 4.5, but the change is not very sharp. A mixed bromocresol green-methyl red indicator will give a sharper equivalence point at the lower pH. It is greenish-blue at pH 5.2, light blue at pH 5.0, light pink-gray at pH 4.8, and light pink at pH 4.5. This bromocresol green-methyl red indicator can be ordered from the Hach Company at 1-800-227-4224.

7. Do not pause for more than a few minutes during the titration process. In water with a high alkalinity, the pH values will steadily climb if left to sit for an extended period of time.

8. The pH of the equivalence point for total alkalinity will vary slightly, depending on the chemical composition of the water. Some sample values are given in Table 2. The correct equivalence point for the sample should be evident from the data.

9. Calibration of the pH electrode is not needed for this test. The measurement of interest is the volume of acid used to reach the equivalence point and not the exact pH value at the equivalence point. The calibration for the pH electrode that is included in the CHEMBIO program will give excellent results. This is due to the fact that even if the calibration was off, the titration curve would still indicate the correct volume of acid added at the equivalence point. If, however, you still wish to perform a new calibration, follow the procedure below:

First Calibration Point

Select PERFORM NEW from the CALIBRATION menu. Follow the directions on the calculator screen to allow the system to warm up, then press .

Remove the electrode from the bottle by loosening the lid, then rinse the electrode with distilled water.

Place the electrode tip into a pH-10 buffer solution. Wait for the readings displayed on the CBL screen to stabilize, then press on the CBL.

Enter "10" (the pH value of the buffer) on the calculator.

Second Calibration Point

Rinse the pH Electrode with distilled water and place it in a pH-7 buffer solution.

Wait for the readings displayed on the CBL to stabilize and press .

Enter "7" (the pH value of the buffer) on the calculator and press to return to the MAIN MENU.

10. It is important to have an adequate supply of pH buffer solutions available. We recommend having at least 100 mL each of pH-4, pH-7, and pH-10 buffer solutions on hand to calibrate your pH probes. Vernier Software sells a set of pH buffer capsules for preparing buffer solutions with pH values of 4,6,7, and 10 (order code PHB, $10). Simply add the capsule contents to 100 mL of distilled water.

You can also prepare pH buffers using the following recipes:

pH 4.00: Add 2.0 mL of 0.1 M HCl to 1000 mL of 0.1 M potassium hydrogen phthalate.

pH 7.00: Add 582 mL of 0.1 M NaOH to 1000 mL of 0.1 M potassium dihydrogen phosphate.

pH 10.00: Add 214 mL of 0.1 M NaOH to 1000 mL of 0.05 M sodium bicarbonate.

11. The pH electrode can be stored short term (up to 24 hours) in pH-4 or pH-7 buffer solution. For long-term storage (more than 24 hours) the pH electrode should be stored in a pH 4 buffer/KCl storage solution in the storage bottle. The pH electrode is shipped in this solution. You can prepare additional storage solution by adding 1.0 g of solid potassium chloride, KCl, to 100 mL of pH-4 buffer solution.

12. The relationship between alkalinity and hardness is sometimes confusing because they are both reported in units of mg/L CaCO₃. Hardness uses these units because the calcium ion, Ca²⁺, is its primary constituent. Alkalinity and hardness levels often rise and fall together as calcium carbonate levels rise and fall, but their levels can diverge because alkalinity also includes HCO₃^– and OH^– ions and total hardness includes Mg²⁺ ions.

Answers to Calculations

1. Answers will vary. Sample data of 10.50 mL H₂SO₄ will be used in the following sample calculations.

2.

Note: The conversion factor of 10.0 is derived from the following equation:

Answers to Optional Calculations

3. Answers will vary. Sample data of 10.50 mL H₂SO₄ will be used in the following sample calculations.

Moles H₂SO₄ = molarity 5 liters = 0.0100 M 5 0.01050 L = 0.000105 moles H₂SO₄

4. X = 0.000105 mol CaCO₃

5. Mass CaCO₃ (g) = moles 5 molar mass CaCO₃ = 0.000105 mol 5 100 g/mole = 0.0105 g

Mass CaCO₃ (mg) = g 5 1000 mg/g = 0.0105 g 5 1000 mg/g = 10.5 mg CaCO₃

6.

The answer to Problem 6 should be the same as the value obtained in Problem 2.

How the pH System Works

The Vernier pH System is a general-purpose pH measurement system consisting of a pH Electrode and a pH Amplifier. The electrode has a gel-filled reference half cell that is sealed—it never needs to be refilled. The glass bulb at the tip of the electrode measures the H⁺ ion activity and produces a corresponding voltage. This voltage varies linearly with the log of the H⁺ ion activity. Because pH = –log[H⁺], the voltage produced varies linearly with pH.