Exercise 4: Determination of Exchangeable Acidity

Soils in humid regions normally accumulate increasing amounts of exchangeable acidity as they weather. The hydrogen and other acidity cations (Al³⁺, Fe³⁺, and Mn⁴⁺) that come from several sources can consume or adsorb OH^– and therefore release H⁺. Carbon dioxide is either given off during respiration in the soil or carried into the soil by rain water. It is converted to carbonic acid when combined with water. Other inorganic acids such as nitric and sulfuric acids are produced in the soil through microbial processes. Nitric acid production is especially large where ammonium or urea forms of nitrogen fertilizer or manure are used. Also, many different organic acids are produced during organic matter decomposition and these, too, tend to acidify the soil.

Other reactions occur when soils become strongly acid (usually below pH 5.5 or so). Aluminum, iron, and manganese ions from within the structures of soil minerals replace some of the hydrogen on exchange sites. These are considered acidic cations because they release hydrogen ions from water molecules during hydrolysis. These hydrogen ions contribute to acidity and soils may develop toxicity that can damage or kill many plants–especially some legumes. Aluminum toxicity is most common, but iron and manganese toxicities also have been reported. The acidic cations (hydrogen, aluminum, iron, and manganese) are so closely allied that they are grouped together and contribute to exchangeable acidity.

Whatever the source, the acidic ions in the soil solution can displace base cations from the cation exchange sites of the soil. These base cations are gradually removed from the soil by either plant utilization or leaching. The cations utilized by growing plants may be recycled and returned to the soil unless they are removed by animals or in harvested crops. The acidic ions replacing the base cations make the soil more acid and less fertile.

A clear distinction should be made between the solution hydrogen to be measured in the exercise on soil pH and the exchangeable acidity to be measured in this exercise. The hydrogen ions in solution are sometimes referred to as active acidity. The acidic ions on exchange sites are sometimes referred to as potential or reserve acidity. The amount of reserve acidity in an acid soil is likely much, much greater than the amount of active acidity. Because the two types of acidity are in equilibrium, it is impossible to neutralize the active acidity without also neutralizing most of the reserve acidity. Unfortunately, the equilibrium governing the relationship between active and reserve acidity is not the same for all soils. Both types of acidity must be measured for a full evaluation of soil acidity. For example, a decision to lime or not lime a soil should be based on the active acidity (soil pH), but the quantity of lime needed to raise the pH by a certain amount depends on the reserve or exchangeable acidity.

The exchangeable acidity will be measured in this exercise by leaching the acidity from the soil with a barium acetate solution, and the resulting leachate will be titrated with dilute base. A dilute sodium hydroxide solution is used as the base because the amount of exchangeable acidity to be measured is usually quite small. Barium acetate is used as the extracting agent because barium ions have a strong replacing power, enabling them to displace hydrogen, aluminum, and other ions from the exchange sites. Recall the position of the acidic cations in the lyotropic series. It takes a strong replacing ion to displace the acidic ions on the cation exchange sites.

The barium acetate solution is adjusted to pH 8 to produce 100% base saturation in the soil. Some alkaline soils contain such low concentrations of hydrogen ions that they absorb H⁺ from a neutral solution and therefore indicate a negative amount of exchangeable hydrogen. This problem should not occur with the solution at pH 8.

Another procedure to measure soil acidity uses a pH-buffered solution and an extracting cation such as Ca²⁺. We will use this method in determining the lime requirement of soil and will be able to compare that value to the value obtained in this exercise.

Procedure

1. Weigh 3.00 ± 0.10 g (record exact weight) of soil (<2 mm) into a 50-mL polyethylene centrifuge tube in duplicate.

The barium ions will replace hydrogen and other cations from the cation exchange sites on the colloids. Barium is a stronger replacing ion than ammonium, potassium, and sodium ions and is therefore preferred for replacing the strongly held Al³⁺ ions. Also, barium tends to flocculate the soil and this helps in the centrifugation procedure.

2. Using a dispenser, add 15.0 ± 0.1 mL of 0.25 M (0.5 N) barium chloride (BaCl₂, pH = 8.0) solution to each tube. Briefly shake and then set on the platform of a shaker horizontally for a 30-min, low-speed equilibration.

The shaking is needed to assure good contact between the solution and the soil for the replacement reaction to occur.

3. When equilibrated after 30 min, allow the soil to settle for 1 min and then centrifuge @ 4000 rpm for 10 min.
4. Filter the supernatant using a Whatman No. 5 filter paper and collect the filtrate in a 250-mL Erlenmeyer flask.

This step separates the soil particles from the solution.

5. Prepare a blank by filtering 15.0 ± 0.1 mL of BaCl₂ solution into a 250-mL flask.

The barium chloride solution at pH 8 absorbs some carbon dioxide from the air and thereby accumulates some hydrogen ions that do not come from the soil. The blank measures these ions so a correction can be made.

BaCl₂ + H₂O + CO₂ ó BaCO₃ + 2HCl

6. Add 5 drops of phenolphthalein indicator to each flask.

Phenolphthalein will serve as the indicator to identify the endpoint. It changes from colorless in acid or neutral solutions to pink in alkaline solutions. The slightly alkaline endpoint given by phenolphthalein is necessary because the solution contains stronger base cations (calcium, magnesium, potassium, sodium, etc.) than acids (acetic). Thymol blue can be used as an alternate indicator in place of phenolphthalein. Its color changes from yellow to blue.

7. Titrate with standardized 0.02 N NaOH to a permanent faint pink color.

A dilute base is used to accurately titrate the small amount of acid contained in the soil extracts and the blank (step 7). The endpoint is near when the solution turns pink around a fresh drop of NaOH but the pink color disappears with mixing. Unlike when a strong base is used in titration (Exercise 1), this end point is a faint pink color that stays in the entire volume of the solution. Your laboratory instructor has standardized the NaOH against potassium hydrogen phthalate (KHP) using the procedure outlined in Exercise 1.

8. When finished, place the titrated solution in the waste container provided.

Barium compounds are hazardous wastes and must be disposed of properly by ISU Environmental Health and Safety.

9. Calculate the exchangeable acidity in milliequivalents per 100 g of soil.

Subtract the amount of H⁺ contained in the barium chloride “blank” from the amount of exchangeable acidity in the soil extract. Then, adjust the results mathematically to a sample size of 100 g.

Remember

Normality × mL = milliequivalents (meq), so

Net meq of NaOH consumed = (V_replicate-V_blank) × N_NaOH, V refers to volume in mL

meq/100 g soil = (net meq of NaOH consumed) × (100/wt of soil analyzed in grams)

Name____________________

Date_____________________

Section__________________

Soil number ________

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