Soil pH

Introduction

Perhaps the most important property of soil as related to plant nutrition is its hydrogen ion activity, or pH (the term "reaction" is also used, especially in older literature). Soil reaction is intimately associated with most soil-plant relations. Consequently, the determination of pH has become almost a routine matter in soil studies relating directly or indirectly to plant nutrition. Knowledge of soil acidity is useful in evaluating soils because pH exerts a very strong effect on the solubility and availability of many nutrient elements. It influences nutrient uptake and root growth, and it controls the presence or activity of many micro-organisms.

 

The pH scale is based on the ion product of pure water. Water dissociates very slightly:

H2O H+ + OH-

Kw = [H+] * [OH-] = 10-14 at 23C

where Kw is the ion product for water and [] indicates the activity of each component in moles per liter of solution. Since [H+] = [OH-] in pure water at 23C, each is equal to (10-14)1/2 = 10-7.

 

The pH of a solution is defined as the negative log (base 10) of the H ion activity or the log of the reciprocal of [H+]:

 

pH = -log10 [H+] = log10 (1/[H+])

 

For example, a hydrogen ion activity of 1/10,000 (or 10-4) mol/L would equal pH 4. Water with equal numbers of H+ and OH- (hydroxyl) is neutral at pH 7 at 23C . pH values below 7 are increasingly acid with excess H+ or hydrogen ions. At 100C the pH of pure water is 6.0 and at 0C is 7.5 (i.e., temperature affects pH).

 

Carbon dioxide dissolves in water to form carbonic acid. Otherwise-pure water in equilibrium with CO2 at its standard atmospheric concentration of 0.033% (330 ppmv) will have a pH of 5.72. CO2 concentration may be as high as 10% in poorly aerated soil pores; water in equilibrium with this air would have a pH of 4.45, although other components of soil solution can raise or lower it.

 

Soil pH

Three soil pH ranges are particularly informative: a pH <4 indicates the presence of free acids, generally from oxidation of sulfides; a pH <5.5 suggests the likely occurrence of exchangeable Al; and a pH from 7.8-8.2 indicates the presence of CaCO3 (Thomas 1967).

 

The fundamental property of any acid in general (and therefore of a soil acid) is that of supplying protons, and therefore the H+ ion activity of a system is fundamentally its proton supplying power. In an analogous fashion, the redox potential (Eh) of a system is its electron supplying power.

 

Hydrogen ions in solution are in equilibrium with those held on soil particle surfaces (i.e., on exchange sites) The soil pH as actually measured represents the active (in solution) hydrogen ion concentration. The total acidity of the soil includes both active and "reserve" (or exchangeable) acidity. Thus, two soils with the same pH may have much different amounts of reserve acidity and one may be more difficult to neutralize than another.

 

Exchangeable aluminum also contributes to soil acidity. When an Al3+ ion is displaced from an exchange site into the soil solution, it hydrolyzes, splitting water and releasing a hydrogen ion to solution:

Al3+ + H2O = AlOH2+ + H+

 

Lime requirement is the amount of a base (in practice, lime or calcium carbonate) needed to neutralize enough of the exchangeable acidity to raise soil pH to a desired value that is more suitable for crop growth.

 

In most soils it has been noticed that pH tends to increase with depth. This is because the upper horizons receive maximum leaching by rainfall, and by dissolved carbonic acid and organic acids which remove metal cations (eg., Ca++, K+, Mg++) and replace them with H+ ions. Lower horizons are not so strongly leached and, in fact, in dryer areas may accumulate calcium and other materials removed from the upper soil.

 

Measuring pH

There are many factors that affect soil reaction as measured in the laboratory. The pH of many soils tends to increase as the sample is diluted with water. Such pH changes may be caused by variables such as carbon dioxide partial pressure, salt concentration, hydrolysis, and solubility of soil constituents. Various soil:water ratios have been proposed for pH determinations. These range from very dilute suspensions (1:10 soil:solution ratio) to soil pastes. The general effect sees the pH of most soils increasing with dilution, and becoming constant at about a soil:water ratio of 1:5.

 

There is no standard procedure for measuring soil pH. Some of the details that vary from one laboratory to the next are: soil:solution ratio, use of a salt solution (e.g., 0.01 M CaCl2) rather than water, method of mixing, time of standing before reading, etc. Soil may be weighed, or measured as a volume (McLean 1982). Therefore, when reporting sol pH, it is essential to include at least a brief summary of the procedure followed.

 

The exact placement of the pH electrode in the sample may be important. When placed in the settled sediment of a suspension of soil of appreciable cation exchange capacity (CEC), a lower pH is generally measured compared to the measurement obtained in the supernatant solution (called the suspension effect). However, the sediment pH can be lower than, equal to, or higher than that of the supernatant depending on the soil and existing conditions. For example, if the soil has a net positive charge and more OH than H ions are dissociated from the soil, the sediment may have a higher pH than the supernatant (Coleman and Thomas 1967).

 

Soil factors in the field that influence soil reaction include degree of base saturation, type of colloid, carbon dioxide partial pressure, oxidation potential, soluble salts, and so on. In addition to these factors the measured pH may vary because of the manner in which the sample is handled in the laboratory before and during the determination. Acquaintance with these variables is necessary for intelligent measurement and interpretation of soil reaction.

 

pH can be determined using either colorimetric or electrometric methods. The choice of method depends upon the accuracy required, the equipment available, or convenience. Many organic dyes are sensitive to pH, the color of the dye changing more or less sharply over a narrow range of H-ion activity. These methods tends to be slower, less precise, and obscured from view by soil particles and organic matter. Hence, they are used mostly in the field where pH is to be approximated.

 

The electrometric method involves a glass electrode that is sensitive to H+: there is an exchange of ions between solution (H+) and glass (Na+) (Westcott 1978). A reference electrode that produces a constant voltage is also required. The electrode pair produces an electromotive force (emf or voltage) that is measured by a millivoltmeter. The relation between emf and pH is governed by the Nernst equation:

where E = emf produced by electrode system

Eo = a constant dependent on the electrodes used

R = gas constant

T = absolute temperature

n = number of electrons involved in equilibrium (1 in this case)

F = Faraday constant

 

(Willard et al., 1988, p. 675).

 

Note that temperature is a factor in the equation. At 25C this equation simplifies to

E = E + 0.0591 pH

which means a change of 1 pH unit produces a change in emf of 59.1 mV, at 25C. This temperature-dependence of pH is important to remember when calibrating a pH meter.

 

 

 

Soil reaction classes

The following descriptive terms are used for specified ranges of soil pH (Soil Survey Division Staff, 1993):

Ultra acid <3.5

Extremely acid 3.5-4.4

Very strongly acid 4.5-5.0

Strongly acid 5.1-5.5

Moderately acid 5.6-6.0

Slightly acid 6.1-6.5

Neutral 6.6-7.3

Slightly alkaline 7.4-7.8

Moderately alkaline 7.9-8.4

Strongly alkaline 8.5-9.0

Very strongly alkaline >9.0

 

References

Cole, CV. 1957. Hydrogen and calcium relationships of calcareous soils. Soil Sci. 83:141-150.

Coleman, NT and GW Thomas. 1967. The basic chemistry of soil acidity. In RW Pearson and F Adams (eds.) Soil acidity and liming. Agronomy 12:1-41. Am. Soc. of Agron., Inc., Madison, Wis.

McLean, EO. 1982. Soil pH and lime requirement. In AL Page (ed) Methods of Soil Analysis, Part 2: Chemical and Microbiological Properties. 2ed. Agronomy #9. pp. 199-224. American Society of Agronomy Inc, and Soil Science Society of America Inc. Madison, Wisconsin.

Soil Survey Division Staff. 1993. Soil Survey Manual. United States Department of Agriculture Handbook N0. 18. Washington, DC. 437 pp.

Sposito, G. 1989. The Chemistry of Soils. New York: Oxford University Press. 277 pp.

Thomas, GW. 1967. Problems encountered in soil testing methods. p. 37-54. In Soil testing and plant analysis, Part 1. Soil Sci. Soc. of Am. Spec. Pub. no. 2, Madison, Wis.

Westcott, CC. 1978. pH Measurements. San Diego: Academic Press, Inc. 172 pp.

Willard, H.H., L.L. Merritt, J.A.Dean, and F.A. Settle. 1988. Instrumental Methods of Analysis. 7th ed. Wadsworth Publishing C. Belmont, CA.


 

Procedure: Soil pH (Electrode Method)

Materials

1.         Sieved soil

2.         Plastic beakers (50 or 100 ml).

3.         Glass stirring rod.

4.         pH meter with combination electrode.

Reagents

1.         2 M CaCl2.

2.         pH buffer solutions -- 7, 4, 10 as needed.

Method

 

Note: The actual soil:water ratio used, and the choice of water or salt solution, are up to the discretion of the analyst, and should be reported with your results. Determining pH in both water and 0.01M CaCl2 gives you useful information with only a little more work.

 

Another note: Be aware that a small amount of fill solution from the reference electrode may leak into your sample. This solution may contain K+, Cl-, or Ag+, as well as mercury, a common preservative in commercial buffer solutions.

 

1.         Weigh 10.0 g soil (air-dry or moist) into a plastic beaker. Add 50.0 ml DDW. Stir with glass rod to mix; let stand 30 minutes, stirring occasionally.

2.         Calibrate pH meter at pH 7 and 4 (or 7 and 10 for alkaline soils).

3.         Swirl the suspension, then carefully insert the combination pH electrode into it. Record pH of supernatant after 15 seconds of settling (or when the reading settles down). The observed pH will vary with where and when in the suspension you take the reading, so be consistent in your method.

4.         Rinse electrode with DW after each measurement; check calibration periodically.

5.         Add 0.25 ml of 2 M CaCl2 to the 1:5 suspension, to make a 0.01 M solution. Mix, let stand, and measure pH as above.

6.         Optional: Saturated paste -- Weigh 20.0 g soil into a 50-ml beaker. Add small increments of water and mix thouroughly until a saturated paste is formed. This stage occurs when there is a smooth, shiny surface to the mixture, but no free water on top. Let stand 30 minutes and measure pH as above.


The Care and Feeding of pH Electrodes

Keep bulb in water/solution as much as possible.

Don't push bulb into bottom of container, or scratch it.

Rinse electrode(s) thoroughly between sample/buffer measurements. After rinsing, blot drydont wipe, which will cause static charges to build up in the electrode.

Check level of fill solution (saturated KCl, or 4M KCl/AgCl, depending on the type of electrode) in reference electrode; it should be well above the level of the solution you are measuring, to provide sufficient hydrostatic head for a steady flow. Refill if neccessary.

There should be free KCl crystals in bottom of reference electrode (that's how you know it's a saturated solution). However, be sure crystals are not plugging up ceramic junction.

Open fill hole when using, to allow free flow of ions during measuring; close fill hole when done, to minimize wasting of fill solution.

Store combination electrode in 10:1 pH 4 buffer solution:saturated KCl (i.e., dilute the KCl solution with buffer.

If using separate electrodes: store glass (pH) electrode in pH 4 buffer; store reference electrode in fill solution diluted by 10.

Because fill solution is flowing out of the reference electrode, you contaminate a sample whenever you place a pH electrode in it, making it unusable for other measurements. In addition to KCl, the fill solution may contain slight amounts of buffer solution that diffuses into it through the porous junction. Commercial buffer solutions often contain mercury as a preservative.

 

Calibration

If all of your sample pH's are clustered closely around a single value, you can use a one-point calibration. However, in most cases you should do a two-point calibration, which will allow you to measure a range of pH.. Buffers and samples must be at the same (preferably room) temperature. For most precise calibration, use fresh buffer solution. Rinse electrode thoroughly between each step.

 

1.         Place electrode in pH 7 buffer. Set to 7.00 with CALIBRATE control. This control sets the intercept of the pH vs voltage regression line.

2.         Place electrode in pH 4 (OR 10) buffer. Set to 4.00 (OR 10.00) with TEMPERATURE control. This action sets the slope of the regression line.

3.         Repeat steps 1 and 2 until both readings are accurate without changing the controls.