Why talk about ph? If plants did not care about soil ph, we would not either. (See 12 th ed., Fig. 9.19; 13 th, 14 th, 15 th ed. 9.

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1 Soil Acidity 12th-14th ed., Chap. 9 Basics of ph (review). Neutral ph = 7 Solutions with lower ph are acidic. Solutions with ph greater than 7 are called either basic or alkaline. Real-world solution ph values (13 th, 14 th ed., Fig. 9.2; 12 th ed. Fig. 9.1). Note that ph of unpolluted rainwater and distilled water is usually around 5.6 to 5.8 depending on how completely those solutions have equilibrated with CO 2 in air. Upper and lower limits of ph Often represented as 0 and 14. Wrong! About the lowest and highest ph values you can achieve in a lab using the strongest possible acids and bases are -2 and 16. Soils in humid temperate climates rarely have ph < 3.5 or ph > 8.2 Soils in dry climates: ph < 8.5 if calcium rich; ph < 11 if sodium rich. Why talk about ph? If plants did not care about soil ph, we would not either. (See 12 th ed., Fig. 9.19; 13 th, 14 th, 15 th ed. 9.23) Classification of acidity (Note: This scheme is based on how well the acidity is hidden.) Active acidity H in soil solution; not hidden at all. Directly sensed by: ph electrodes, ph-indicating dyes, plant roots, soil microbes and invertebrates. Magnitude. In a soil with ph = 4 and 20% moisture, 2 lb of lime, in theory, could neutralize all the active acidity in an acre furrow slice. Acidity due to aluminum (Though mentioned, not explicity considered as a category for soil acidity in Brady.) Al 3 present particularly in soil solutions with ph < 4.5. Invisible to ph electrodes and ph-indicating dyes. Toxic to most plants (at higher concentrations) and to many microbes. Magnitude: Usually small. Negligible near ph 7, much greater at low ph. Digression: Why is Al 3 an acid? Add to distilled water ph change HCl Lower CaCl 2 None AlCl 3 Lower But, unlike classical acids like HCl, HNO 3, or H 2 SO 4, there are no obvious H s in the molecule that are just waiting to fall off as H ions. So how can Al 3 (which is what you get when you add AlCl 3 to water) be an acid. Answer: Al 3 is a Lewis acid that creates H ions by attacking water molecules. Many metallic ions, like Fe 3, also act as Lewis acids.

2 Aluminum toxicity. As soil ph falls, the concentration of Al 3 in solution and on cation exchange sites increases (12 th ed., Fig. 9.2; 14 th ed., Fig. 9.19). Organic matter can bind aluminum in non-exchangeable forms that pose no danger of toxicity. The figure at left shows how higher Al 3 levels produce shorter cotton roots. The figure on the right shows how raising ph can decrease toxicity from Mn (another Lewis acid), thereby boosting yields. (14 th ed., 9.20; cf 12 th ed., Fig ; 13th ed., Fig. 9.20) Some crops have been genetically engineered to produce elevated levels of enzymes that contribute to resistance to aluminum toxicity (12 th ed., Fig. 9.28; 13 th ed., Fig. 9.30; 14 th ed., Fig. 9.29). Salt-replaceable acidity H and Al 3 ions occupying cation exchange sites by simple ionic attraction. Hidden from direct measurement with probes or dyes Can be displaced into solution (where the H can be detected) by cations from salts Magnitude. Depends on texture. Maybe 100 more than active acidity. Residual acidity H, Al 3, AlOH 2, and Al(OH) 2 bound to clays through covalent (or partially covalent) bonds. These cations occupy exchange sites in ways that resist displacement by simple cations from salts. For example, AlOH 2, and Al(OH) 2, hold onto cation exchange sites far more tightly than Al 3 does, so most of the divalent and monovalent aluminum hydroxides contribute to residual acidity. Hidden from measurement by probes and dyes, and undetectable by life forms. Magnitude. The largest category of acidity. Increases both with increasing cation exchange capacity and with decreasing ph. Requires 1000 to 100,000 more lime to neutralize than the active acidity in the same soil. Classification of Cations Genuinely acidic cations H, Al 3, (and Fe 3 at very low ph) Reason for classification: ph /!log 10 {H } Al 3 and Fe 3 are Lewis acids that will create H in solution when bases are added to try to raise ph. All other cations Examples: Ca 2, Mg 2, K, Na, NH 4 Common names for this category: Basic cations (old, out of favor, but still common) Non-acid cations Not actual bases. Add CaCl 2 or KCl to water and ph stays the same.

3 Percent Non-acid Saturation, by definition: cmol of charge in non acid cations %NAS cmol of charge on exchange sites 100 The above equation assumes that the number of cations held on exchange sites hugely exceeds the number of cations free in the bulk soil solution: usually a safe assumption in moist climates. As formulated above, %NAS quantifies that percentage of all the positive charges held on cation exchange sites that happen to be part of non-acid cations, rather than acidic cations. In a soil where essentially all the cations consist of H, Al 3, Ca 2, Mg 2, and K, percent nonacid saturation would be calculated as follows: 2 2 2[ Ca ] ex 2[ Mg ] ex [ K ] ex % NAS = [ H ] 3[ Al ] 2[ Ca ] 2[ Mg ] [ K ] ex ex ex ex ex Where square brackets, [ ] ex, mean number of cmols of the enclosed ion per kg of soil. When an individual cation has a charge greater than 1, the concentration of the ion is multiplied in the above formula by the charge per ion to produce the total number of positive charges associated with that ion. The numerator in the formula is identical to the denominator with all the concentrations of the acidic cations deleted. %NAS is powerfully related to soil ph, because... [ H ] ex [ Nonacid ] ex [ H ] aq [ Nonacid ] aq as the ratio of the concentrations of acidic to non-acid cations held on exchange sites increases, the same ratio for the concentrations [ ] aq of cations free in the bulk solution also increases. The two ratios are not equal, but they are in equilibrium with one another. If you pour a strong solution of CaCl 2 on a soil sample, you will be adding a lot of Ca 2 ions (as well as a lot of Cl! ions) to the soil solution. This increases the denominator of the fraction on the right of the above equation, which decreases the size of that fraction. You will have decreased the ratio of acidic to non-acidic cations in the bulk soil solution. In response, a net movement of Ca 2 ions from the solution onto cation exchange sites will occur, and this will displace H ions from exchange sites they previously occupied. You added something to the solution that caused the ratio on the right to drop, and the ratio on the left also dropped in response. Although they are formulated a little differently mathematically, the ratio on the left increases when %NAS increases. Bottom line: When H ions are abundant in the soil solution (and ph is therefore low), acidic cations are also abundant in cation exchange sites (and %NAS is also low).

4 The text has a pair of figures illustrating the way in which exchangeable non-acid (a.k.a baseforming ) cations and acidic cations change their concentrations as a function of soil ph. (12 th ed., Fig. 9.4; 13 th and 14 th eds., Fig. 9.5) Buffering means the ability to resist a change in some variable, like ph. The vast pools of acidic and non-acid cations stored on exchange sites buffer soil ph. Pour 1 ml of sulfuric acid in 1 kg of water, and the ph of the water will drop enormously. Mix 1 ml of sulfuric acid into 1 kg of moist topsoil, and the ph of the soil will not drop nearly as much as the ph of the water did. The soil buffered ph. Pure water with just enough HCl or NaOH added to adjust ph across a range of 2 to 10 lacks buffering capacity. If such a solution is added to a dry soil sample, the stored cations on exchange sites very quickly adjust the ph of the water (the new soil solution) to the value the soil thinks is correct, i.e., the value that is in equilibrium with %NAS of the soil. %NAS controls soil ph (12 th ed., Fig. 9.5; Dropped from later editions). Note that the total amount of acidity stored on exchange sites does not determine soil ph; the fraction of the exchange sites covered by acidic charges does. A large coffee urn with a glass-tube level indicator provides a useful analogy for the relationship between reserve acidity and active acidity. (12 th ed., Fig. 9.7; 13 th ed., Fig. 9.9, 14 th ed., Fig. 9.10) Old clays like acid. Take samples of an old, hydrous oxide-type clay, an intermediately old 1:1 clay like kaolinite, and a relatively young 2:1 clay and adjust all their ph values to 7.0, and the old clay will have the highest percent acid saturation (and, therefore, the lowest percent nonacid saturation). Soils scientists actually use percent non-acid saturation at ph 7 as a quick-tomeasure way to figure out how old the clays in a particular soil are. Practical ph Measurement Colorimetric method (ph-indicating dyes) Advantages Easily transported to field No calibration required Quick to get first reading Cheap if used rarely Disadvantages Not terribly accurate Difficult to bulk multiple samples to get average Can turn pockets funny colors Bottom line Good for one-time quick and dirty measurement (e.g., enough to tell you that it s time to send in a sample to a soil testing lab).

5 Electrometric method (ph-electrode and meter) Advantages Once calibrated, subsequent measurements take little effort Capable of great accuracy Cheap if used frequently Disadvantages Calibration process takes time Accuracy no better than dyes without proper calibration Time required per sample can be a minute or more due to slow stabilization of reading, esp. in acidic soils. Initial acquisition costs > dyes for a good system. (You get what you pay for.) Slow to revive after sitting around idle for weeks Miserable, stinking electrodes break all the time! Well suited for lab operations requiring many accurate readings. The need for ph adjustment We live in a leaching environment. Lots of rainwater and snow melt percolates downward through our soil to the water table. An area that receives only enough precipitation to support grasslands (but not forests) is a non-leaching environment, as is a desert. In a leaching environment, farmed soils slowly become acidified, because non-acid cations are constantly leaching (with the passing water) downward to groundwater. Those cations are replaced on exchange sites by acidic H ions from the carbonic acid rainwater. (Acid rain just speeds up the process.) Eventually, the acidity must be neutralized to adjust ph to suit plant needs. Even in an leaching environment, nobody needed to lime a natural broad-leafed treedominated forest (before the invention of acid rain) because the tree roots intercept almost all the non-acid cations before the can be leached out of the soil. These cations are recycled to the surface of the soil when trees shed their leaves. Ecologists call natural forests tight with respect to nutrients (like non-acid cations). They call farmed soils leaky. For most of the growing season, much root-free soil exists between crop rows through which non-acid cations can freely leach. Liming agents Lime. A calcium-containing material that may be applied to soil to neutralize soil acidity. Agricultural limes. Calcitic limestone. Liming material consisting of ground calcite (primarily CaCO 3 ). Dolomitic limestone. Liming material consisting of ground dolomite (a mixture of CaCO 3 and MgCO 3 ). Reacts more slowly than calcite of same particle size. Note: Mg is an important plant nutrient. Non-agricultural limes. Quick lime (a.k.a., burned lime). Primarily calcium oxide, CaO produced by heating calicite in a furnace: CaCO 3 6 CaO CO 2. Hydrated lime (= slaked lime). Produced by adding one H 2 O molecule per O

6 atom in a quick lime to create a hydroxide. For example, a hydrated calcitic lime is produced as follows: CaO H 2 O 6 Ca(OH) 2 Speed of action Particle size. Finer particles react faster than coarser particles of the same chemical composition. Chemistry. Particles of the same size react in the following order of decreasing speeds: Oxides. Hydroxides >> CaCO 3 > CaMg(CO 3 ) 2 Extent of incorporation. Without mixing (e.g., in turf, effects of surface-applied lime may move downward at only 1 cm per year through a fine- textured soil. Amount of lime to add. You need two numbers: Lime requirement or recommended liming rate. The amount of pure CaCO 3 that would be required to raise ph to the desired value. Units could be lbs per100 ft 2 or tons per acre. Commercial soil tests always list the lime requirement this way. Neutralizing value. The strength of your commercial liming agent relative to the strength of pure CaCO 3. Units might be (lbacaco 3 )/(lbalime). Quick and slaked limes generally have neutralizing values > 1 (or > 100%), but calcite containing many impurities is likely to produce (when ground) a lime with a neutralizing value < 1. Bags of lime should have this number on the label. You now can calculate actual liming requirement: (Lime requirement lb CaCO 3) Actual lime requirement = lb CaCO (Neutralizing value ) lb lime 3 Two lime-related terms Caustic. Capable of burning or eating away organic tissue through chemical action. The oxides and hydroxides (burned and slaked) are caustic. The carbonates (the ground rocks) are not. First symptom: fingers feel slippery. Hygroscopic. Capable of absorbing water from the vapor in air. Causes powdered materials to cake and become difficult to spread. The oxides and hydroxides are hygroscopic. The carbonates are not. Hygroscopic limes should be sold in bags with waterproof liners.