By now students on GE3246 are probably either bored or confused by my seemingly constant references to pH, acid deposition, acidification and their role in environmental pollution and its effects. The truth is, pH, and more particularly variations in pH, has a major part to play in discussions concerning environmental pollution; changes in pH can be an effect (a result) of pollution. Moreover, pH variations can influence the amount of human and ecosystem harm caused by other chemicals in the environment, such as heavy metals, but also impact the availability of nutrients essential to the fertility of soil and hence to plant growth and crop productivity. Already in GE3246 we have discussed acid deposition – its sources and effects. A major source is air pollution in the form of carbon, sulfur and nitrogen oxides [CO2, SO2 & NO] released from power stations etc. Ammonia (NH3), associated with traffic pollution and the misuse of fertliizer, is also a source of acidification pressure. These gases combine with moisture in the atmosphere to form acid and when deposited go on to cause acidification of lakes and rivers and the ocean. In the lecture this coming week we will add the acidification of soils and its effects to this list ….
Environmental acidification occurs naturally, of course, particularly where the underlying geology is base-poor. Humans have through their activities made the situation far worse, however, by causing anthropogenic (or cultural) acidification.
So what is pH and how is it linked to environmental acidification and anthropogenic acidification in particular?
The Danish scientist S. P .L. Sørensen first introduced the term pH (though he denoted it as pH) to mean the amount (or power) of hydrogen (H+) ions (protons) in a solution. If a solution (e.g. lake or ocean water, or even the water between individual particles of soil, known as soil pore water) becomes more or less acidic, the concentration of H+ ions becomes stronger or weaker. Acidity (or its converse alkalinity) is the amount of hydrogen and hydroxide (OH–) ions present in solution, or relative concentrations of these. As acidity increases, the amount of H+ ions increases while levels of OH– ions decline. The table below shows the reciprocal relationship between H+ and OH– ions.
Note that the pH scale ranges from 0 to 14. Solutions with a pH less than 7 are generally referred to as acidic, while those with a pH greater than 7 are basic or alkaline. A pH of 7.0 is taken to be neutral, because the H+ and OH– ions balance one another. It is perhaps worth pointing out that temperature directly effects pH. An increase in temperature causes ionisation to proceed at a higher rate than at lower temperature, causing more H ions to be freed from water molecules ( H2O ⇌ 2H+ + OH– ). It is also worth pointing out that pH is the negative log of the molar hydrogen ion concentration (-log10[H+], or [H+] = 10-pH). What that means in English is that as pH falls, acidity increases, and a pH fall of 1.0 unit represents a 10x increase in acidity (or 10 x more H+ ions)!
So how is pH linked to anthropogenic acidification and its environmental effects, such as coral and forest death?
Hydrogen ions are highly positively charged (the + indicates that). Hence acids can be highly reactive and corrosive (ever spilt battery acid on your clothes?). OH– is also highly reactive, hence strong alkaline solutions, such as bleach, are highly caustic (common bleach is Sodium hypochlorite, NaClO, composed of one sodium (Na) atom, one chlorine (Cl) atom and one oxygen (O) atom.). Heavy metals are also positively charged (positively charged atoms are known as cations – e.g. copper (Cu2+), magnesium (Mg2+), manganese (Mn2+), lead (Pb2+), zinc (Zn2+)), in fact most inorganic contaminants are positively charged (or cations). Many nutrients that plants required (e.g. Potassium, K+, Calcium, Ca2+) are also cations. This sounds confusing, but some metals, including some heavy metals (copper, iron, zinc etc), are classed as nutrients at low concentrations (i.e. they are micro-nutrients) but are toxic at higher concentrations. As mentioned in an earlier lecture, the dose makes the poison (although hopefully everyone realises by now that this is a gross over-simplification).
Acidification has direct effects on biota. Generally however its effects are less direct, and take place through the effects on concentrations of heavy metals, water transparency, availability of nutrients, including calcium carbonate (CaCO3), the material for coral exoskeletons. The infografic below shows how key aquatic organisms may be lost as the pH of their habitat becomes more acidic.
Acidification impacts soils through changing the availability of chemical elements, some of which can be harmful to biota and have negative impacts on agricultural productivity. Soil particles, organic matter etc, tend to be negatively charged – and therefore they can hold (adsorb) cations on their surface (when something is adsorbed on a surface it is basically adhered, or stuck/glued, to the surface). Soils differ in this ability – known as Cation Exchange Capacity (CEC), depending on their texture (e.g. proportion of clay, silt and sand sized particles, amount of organic matter). Generally organic-rich, finely textured soils have have high CEC – they are sticky as far as nutrients and other cations go – and are therefore often relatively “fertile”, and good for agriculture. Coarse-grained (with a high proportion of sand-sized particles) soils with low organic matter content tend to have low CEC, and any nutrients present may be easily lost to leaching and flooding of the soil. The schema below shows how the finest plant roots (root hairs) interact with surrounding soil particles and soil pore water, and the cations attached to negatively-charged soil particles.
Note that in order to obtain essential nutrients (cations) from soil a plant must pump out H+ ions. These H+ ions are swapped for the cations (such as K+ and Ca2+). For an excellent, short video on the Cation Exchange process between plants and soil/soil pore water (and the influence of soil texture) – see here. More on the influence of soil texture is available via the short video linked here.
Interestingly, the video explaining Cation Exchange linked above states that the vast majority of cultivated soils are negatively charged but some soils, mostly in the tropics, are positively charged. We’re in the tropics – so which soils are positively charged? The answer is those soils that are ancient and that have been subjected to intense and deep weathering over a long period of time. As a result of prolonged, intense weathering clay minerals disintegrate, losing their silicon (Si) in the process. As a result, the weathered soil has a lower negative charge (and may even have a positive charge), and thus a lower CEC. This reduced CEC is one of the reasons why lateritic soils in the tropics are less fertile than their younger, less intensively weathered counterparts in more temperate latitudes.
Increased acidity (increased concentration of H+ ions) results in many more H+ ions in the soil pore water competing with dissolved cations for attachment sites on soil particles. Because the H+ ions are generally more reactive they tend to attach to vacant attachment sites on soil particles and displace already attached cations of, for example, heavy metals. The latter are then dissolved in the soil pore water where they are much more biologically available, and can be taken up by plants through their roots. In solution, heavy metals such as Aluminium (Al) bind with phosphorus fertilizers (generally in the form of Phosphate, PO4), forming – for example, AlPO4, which is a form of phosphate that plants cannot use (cannot take up). Hence fertilizer applications to acid soils are often ineffective, and farmers must first find a way of neutralising the acidity (generally by liming, the application of quicklime, CaO).
The pH of solution also influences solubility of heavy metals. Generally heavy metals are in solution in acidic water. At higher pHs (strongly alkaline), the OH– ions bind with the heavy metal cations forming solid compounds, which precipitate-out (they form “precipitates”). Varying the pH of wastewater, reservoir water etc entering water treatment plants is the main way of removing heavy metals from large volumes of water that are then used for human consumption, irrigation etc., as mentioned in an earlier lecture.