Alkalinity and Buffering Capacity of Water

Reference Chapters: 6, 8


After performing this experiment, the student shall be able to:

• Determine the alkalinity and buffering capacity of several types of water samples: surface water, groundwater (mineral water) and sea water.

• Prepare different solutions or mixtures of acids and their conjugate bases (i.e., buffers), and measure their buffering capacity by titration with acids and bases.

• Calculate the concentrations of an acid and of its conjugate base to create a buffer for a desired buffering capacity at a specific pH.

• Prepare a buffer system.


This experiment will allow the determination of the alkalinity and buffering capacity of water samples from different natural sources. The buffering capacity is the ability to neutralize the pH and the resistance to change in it due to the small acidic or basic inputs or discharges. When a system is poorly buffered, the addition of even small amounts of an acid or a base will noticeably alter its pH, but when a system is well buffered, the same addition barely modifies its pH (i.e., it becomes relatively insensitive to the addition of small amounts of acids or bases). The buffering capacity of a system is defined as the moles/L of strong acid (or strong base) needed for a change in one pH unit of a solution. A typical buffer is formed by a combination of a weak acid

(or base) with its corresponding salt. For example: HA(aq) ^ H+q) + A~aq) (la)

The equilibrium (acidity) constant is: [H+][A]

from which we can derive the equation of p Ka (= — log Ka) with respect to the pH:

This is known as Henderson-Hasselbalch equation and it is built under the assumption that [H+] or [OH-] « [HA] and [A~], where HA = weak acid, and A~ = the corresponding anion generated from the salt. In a well-buffered system, the greatest resistance to changes in pH will occur when the ratio of concentrations of the acid and its salt are approximately equal and therefore the yKa will be equal to its pH. From the above equation it is clear that this occurs at [A~]/[HA] = 1.

By knowing the pKa of the buffering acid, one can estimate the pH at which its greatest buffering capacity will be centered. The pH of a buffer solution is affected by two factors: the concentration ratio, [A-]/[HA] (i.e., the inverse ratio of the acid to the conjugate base), and the strength of the parent acid or base. The stronger the parent acid or base in the buffer solution, the more extreme will the buffer's pH value be.

The buffering capacity depends on the concentration of the buffer, and on the type and concentration of the acid or base to be added to the buffered solution. In selecting the right working buffer for a specified pH, it is common to consider that its pKa must be at least one pH unit above or below the working pH.

The buffering capacity in natural waters is mainly due to the carbonate system and its equilibria. Therefore, it is important to know the alkalinity of the system, because this will provide the capacity for neutralizing an acid. The expression for alkalinity (i.e., dissolved species only) or acid neutralizing capacity (ANC) (i.e., the whole sample) is generally based on the carbonate system:

and this property is expressed as mg/L (or in eq/L, in the case of ANC) of the equivalent calcium carbonate.

The ANC of natural water systems depends on the composition of the watershed. If there are minerals with poor solubility in the surrounding soil, the ANC will be low, whereas if calcareous minerals are present, there will be a high ANC. Some dissolved organic substances derived from decaying plant materials may also contribute to the ANC capacity of the water.

In this experiment the student will use samples as those considered in Experiment 1 and determine their alkalinity and buffering capacity. The student will also prepare a series of solutions where the concentration of a known conjugate salt will vary, measure how the buffering capacity changes with the proportions of the salt, and calculate the buffering range or limits. In this case, the proportion of salt that gives the highest buffering capacity or range (and its corresponding pH) will be determined. Finally, the student will calculate the composition of a buffer solution required to give a specified pH; prepare it based on theoretical concepts and equations, and measure its pH and buffering capacity.

Experimental Procedure

Estimated time required to complete the experiment: This depends on the number of samples analyzed (approx. 5-20 min per sample per analysis).

This experiment may be carried out in one or two sessions.

The first step is to obtain water samples. The student should obtain the samples from the original sources similar to what was done for Experiment

1. Appropriate samples include river or lake water, groundwater (if possible from a well, and if not, a sample of bottled mineral water—preferably from a natural source), seawater (if not available, it can be prepared synthetically or obtained from a commercial source). The samples must be collected in clean polyethylene bottles and analyzed immediately after sampling; otherwise, the values may vary substantially.



1 pH meter with a thin test tube

mineral or groundwater

pH combination electrode

river or lake water

polyethylene bottles


4 50-mL beakers

D.I. water

1 1-mL volumetric pipet

0.25 M, 0.1 M NaOH

2 2-mL microburet

Phenolphthalein indicator

1 three-finger clamp

methyl orange indicator

2 Beral pipets

bromocresol green indicator

5 25-mL beakers

0.25 M, 0.1 M, 0.01 M HCl

2 10-mL volumetric pipets

0.01 M Na2S203

1 2-mL volumetric pipet


M Na2C03

3 10-mL Erlenmeyer flask


M NaHC03

2 propipet bulbs or adapted


M KH2P04



M K2HP04

1 universal stand


M H3P04

1 buret clamp

2 5-mL volumetric pipets

1 2-mL graduated pipet (1/100)

1 stirring plate

1 micro magnetic stir bar

5 25-mL volumetric flasks

1 50-mL volumetric flask

2 25-mL burets

The first part of the experiment (Part A) measures the pH and alkalinity of the water samples and compares the values among them as well as with those of tap and distilled or DI water. The alkalinity will be measured by titration with dilute hydrochloric acid up to three specific pH values: 8.3, 7 (i.e., neutrality), and 4.5. This will make it possible to calculate the different contributions of the species responsible for the alkalinity. The acid-base indicator phenolph-thalein is used to indicate the 8.3 endpoint and to obtain the P-alkalinity. Results can be best interpreted in the light of the discussion of Section and Example 6.5.

To identify the amount of acid needed to reach the 4.5 endpoint, one may use methyl orange or bromocresol green indicators. This titration indicates the presence of the rest of the carbonate ions present (when it reaches a pH near 6.5) and of all the bicarbonate ions present up to pH 4.5. The total amount of acid indicates total alkalinity. For a more accurate titration, the pH must be followed with a pH meter. The student will compare the amount of each titrant added and the pH of each sample, besides determining the alkalinity and graphically determining the buffering capacity of each. The student will also identify the weak or strong buffers.

In the second part of the experiment (Part B) the student will prepare solutions with different proportions of carbonate and phosphate salts, and will titrate them to evaluate their buffering limits and the proportion that will yield the best buffering capacity. The student will observe the effect of varying the proportions of the acid to the conjugate base as well as of altering the concentration of the buffering constituents present. The differences among the conjugate systems used will also be observed.

In Part C the student will use the Henderson-Hasselbalch equation to determine the conjugate base/weak acid ratio to buffer a desired pH and prepare it experimentally. The initial pH and buffering capacity are then determined by titration, and they are monitored through the pH changes that result from adding specific amounts of standardized acid or base.

Safety Measures

The titrants should not come in contact with the skin or eyes because they may be corrosive. The student must consider all the safety measures normally taken when handling this kind of reactants. In case of spillage of an acid solution or of skin contact, wipe clean with a clean cloth and wash thoroughly and abundantly with water (sprinkle the table or surface with sodium bicarbonate). All of the residues generated in this experiment can be disposed of down the drain once they have been neutralized.

Experimental Sequence Estimated time required: 5 min per sample Part A. Measurement of alkalinity or ANC Samples and procedure:

1. Collect samples (or prepare them synthetically) of: (1) river or lake water, (2) ocean water, (3) mineral water (or if available, use groundwater), (4) tap water, and (5) D.I. water. Use a calibrated pH meter throughout the entire experiment and wait until stable readings are obtained.

2. Measure the pH and temperature of each sample. If the pH is above 8.3, determine the P-alkalinity. If it is not, only the M-alkalinity (i.e., methyl orange or bromocresol green alkalinity at pH = 4.3) can be measured.

3. If pH < 4.3, measure the acidity of the sample.

4. To prevent masking of the endpoint when using colored indicators, make sure the sample is colorless and free of turbidity. To have this condition, filter the sample prior to titration and—if the color were a problem—add a small amount of activated carbon (prior to filtration). Filtration can be done by means of a syringe equipped with a filter holder and a fine-pore filter. To eliminate any free chlorine that might interfere with the titration, add a drop or two of 0.01 M sodium thiosulfate.

5. Place a microburet in a stand, rinse it with a small amount of the 0.01 M acid titrant, and fill it to the desired mark.

6. With a volumetric pipet, take a 10-mL portion of the water sample and place it in a 25- or 50-mL beaker. Place a micro magnetic stir bar inside the beaker and put the beaker on a stirring plate. Immerse the bulb of the pH electrode in the liquid sample, without touching the stir bar. Measure the pH.

7. Add 2 to 3 drops of the phenolphthalein indicator (e.g., with a Beral pipet). Observe if any color appears.

8. If upon adding the P indicator there is no color, then the P-alkalinity is zero. In this case, immediately add 1 to 2 drops of the methyl orange (or of the bromocresol green) indicator and start titrating to the endpoint.

9. If a pink color appears upon adding the phenolphthalein indicator, titrate drop wise (in 0.1 or 0.2 mL increments), stir gently, and record the resulting pH and volume of titrant added. Note the pH reading at the point when the color of the indicator disappears (it must be close to 8.3). Then, immediately add 2 to 3 drops of the methyl orange (or the bromocresol green) indicator, and continue the titration until the exact 4.5 endpoint is reached (the solution turns salmon in the case of methyl orange or yellowish in color, in the case of the bromocresol green indicator).

10. Repeat the same technique for each water sample.

B. Buffering Capacity

Estimated time required: 15 min per sample


B.l Buffering capacity of natural samples. The measurements carried out in part A to determine the alkalinity of several water samples will also serve to determine the buffering capacity of each sample.

B.2 Factors that affect the pH and buffering capacity

B.2.1 Prepare a Na2C03 + NaHC03 solution by introducing with a volumetric pipet 10 mL of freshly-prepared 0.1 M solutions of each salt into a 25-mL volumetric flask. Add D.I. water up to the mark. Mix perfectly and pour the resulting solution into a 50-mL beaker.

(a) Based on the known pKa values and the Henderson-Hasselbalch equation, calculate the pH of this solution.

(b) Immediately after preparing the solution, measure its pH by placing the pH electrode inside the solution and reading the stabilized value. Remember that if the solution is allowed to rest for a long time, atmospheric CO2 will dissolve in it and the pH values will be altered.

(c) Rinse a 25-mL buret with 0.25 M HC1 and fill it with the same acid solution. Start titrating the solution with the pH probe inside, gently mixing with the magnetic stirrer. Add small increments of acid. Record the volume of acid added and the resulting pH; make sure that by adding the acid there is a change of more than one pH unit in the resulting solution. Preferably in the last part of the titration, switch from the 25-mL buret to a microburet in order to add small increments (e.g., 0.1 mL) up to the lowest constant pH value attainable. Record all your data.

B.2.2 Proceed exactly as in point B.2.1, but use 5 mL of each carbonate and bicarbonate solution

Table C.l Ka and pKa Values for the Acid and Conjugate Base of the Phosphoric Acid System


Conjugate base





7.62 x 10"3




6.23 x 1(T8




2.2 x 10"13


(instead of 10 mL) and follow the same technique as established in B.2.1 (a) through (c).

B.2.3 Proceed exactly as in point B.2.1, but use 10 mL of the Na2C03 solution and 1 mL of the NaHC03 solution.

B.2.4 Proceed exactly as in point B.2.1, but use 10 mL of the NaHC03 solution and 1 mL of the Na2C03 solution.

B.2.5 Proceed exactly as in all the previous steps, but use 5 mL of a KH2PO4 solution and 5 mL of a K2HPO4 solution.

C. Calculation, preparation and evaluation of a specified pH buffer solution

Estimated time required: 20 min per sample


C.l Based on the Ka values for the phosphoric system (see Table C.l) and the Henderson-Hasselbalch equation, the student groups will prepare 50-mL of a phosphate buffer with a requested pH. (The proportions used in experiment B.2.5 are not permitted here). For example, the groups can prepare buffers with pH values of 4, 5, 6, 7, 8, or higher.

Each group will then take its corresponding buffer and measure the acid buffering capacity and base buffering capacity, using the following method. Use a calibrated pH meter throughout the entire experiment and wait until stable readings are obtained.

C.2.a Immediately after preparing the phosphate buffer solution, place 25 mL in a 50-mL beaker and measure its pH.

C.2.c Repeat the process with another 25 mL aliquot of buffer solution, but this time titrate with 0.25 M NaOH. Record all your titration volumes.



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  • john williams
    How to measure buffering capacity useing alkalinity?
    2 years ago
  • kathrin beich
    How to calculate buffer capacity for environmental water samples (mmol hcl/l solution)?
    2 years ago
  • lucas
    Does the alkalinity suggest your groundwater can buffer acidity well?
    1 year ago
  • janina
    How to find the alkalinity of water when using buffer capacity?
    11 months ago
  • ethan
    How to break down the buffering capacity of water?
    11 months ago
  • Aaron
    Does lower alkalinity have better buffer capacity?
    6 months ago
  • eleleta
    What impact does total alkilintiy have on buffering capacity?
    4 months ago
  • Sheryl Haas
    Does silica add alkalinity and buffering capacity?
    2 months ago

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