What exactly is soil acidification, anyway? As the name might indicate, soil acidification refers to an increase in the concentration of hydrogen ions within soils, and a subsequent drop in pH (which is simply a measure of the concentration of hydrogen). Soil pH generally ranges from around 4.5 to 9, with those values being the extremes. A very strongly acid soil is only about as acidic as a tomato, while a very strongly basic soil is only about as basic as milk of magnesia. Regardless of the initial pH, any soil can be acidified.
Why do soils acidify? Short version - hydrogen ions accumulate in the soils. Long version - hydrogen ions are added to the soil at a rate or in a quantity that overwhelms the availability of the soil to neutralize them. Let’s take a voyage into the exciting world of soil chemistry.
We’ve actually already met most of the players in this story. Remember the cations from Group I and Group II of the periodic table? Elements like sodium and potassium shed their outermost electrons to become positively charged cations. Well, in soils, you have a class of cations called base cations - calcium (Ca), magnesium (Mg), potassium (K) and sodium (Na). They’re called base cations because they are cations that counteract acids, which in chemical nomenclature makes them bases. We’ll do straight-up acid-base chemistry one of these days. Anyway, we have the base cations in soils, and some other cations - hydrogen (H), and metals like aluminum (Al) and iron (Fe). We also have actual soil particles, called colloids. Colloid surfaces tend to be negatively charged, which attracts the various cations. In a soil, all of the soil colloids have a number of associated cations - the charge attraction here is called adsorption.
Now, when a hydrogen ion enters the soil, it first enters the soil solution - water in the soil. If the hydrogen ion stays in solution, the pH of the solution will drop. Often, the H will switch places with an adsorbed base cation. Now, instead of, for example, an Mg ion on the soil colloid and an H ion in solution, there is an H ion on the soil colloid and an Mg ion in the soil solution. The soil solution maintains a constant pH, and the Mg ion is either taken up by a plant (Mg is a plant nutrient), or else leaches to a stream (most soil solutions eventually leach to surface water). All of this works very nicely when the rate at which base cations adsorb to the soil colloids is the same at which base cations switch places with H ions and desorb from the soil colloids. The trouble starts when the rate of H input exceeds the rate of base cation input - eventually, the soil solution will lose its ability to neutralize acidity, and the pH will start to decline.
This process happens over a number of time scales. Rainwater naturally has a pH of around 5.5, thanks to the carbon dioxide dissolved within (even before global warming, there was carbon dioxide in the atmosphere). So, over time, soils can acidify simply through rainwater input. This takes a while - ultisols, the ancient, weathered soils of the tropics are typically low in pH, but are also typically hundreds of thousands of years old. Soils forming under conifer canopies are often acidic as well, as a consequence of organic acids leaching out of leaf litter of the course of several thousand years. And then we have acid rain - rainwater with elevated acid concentrations stemming from industrial emissions. Acid rain can acidify a soil in as little as forty years.
Soil acidification, particularly rapid soil acidification, can prove to be bad on many levels. First, there is the potential for acid damage to plants. Most of the negative effects of soil acidification, however, are a bit more indirect than that. Take a look at this chart.
Soil pH affects many things, including how available nutrients are within a soil. A number of important plant nutrients, including nitrogen (N) and phosphorus (P) are available in several different forms within soils, dependent on pH. As it happens, the forms of N and P that are easiest for plants to take up are predominantly found at higher pH - as soils become more acid, the actual amount of N and P does not necessarily change, but the ease with which plants take them up from the soil does. The base cations tend to be less available at extremely low pH, and in lower concentrations - after all, acid soils are often a consequence of low base cation concentrations. Micronutrients like zinc (Zn) and manganese (Mn) become more available at lower pH, but there is a reason these elements are micronutrients. Low pH soils can increase the risk of Mn and Zn toxicity.
This isn’t even scratching the surface of aluminum toxicity - Mn and Zn, aluminum is more available at lower pH, and can be quite a potent toxin to plant and animal life.
Soil acidification has all kinds of other effects, many of them quite complex and indirect. That’s the insidious thing about it - the popular image of acid rain is marble statues melting, but the real damage is a subtle, long-term alteration of natural soil chemistry, in a way that is seldom good for plant life. Not as scary as a melting statue, perhaps, but frightening enough in its own right. Fortunately, much has been done over the last few decades to combat the problem, and as the last few years of my life show (I hope), much more is in the works.
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