Biological and Environmental pH

Human blood must stay within a strict pH range of 7.35 to 7.45. To achieve this, the body uses the bicarbonate buffer system. This system relies on an equilibrium between dissolved carbon dioxide ($CO_2$), carbonic acid ($H_2CO_3$), and bicarbonate ions ($HCO_3^-$). When excess acid ($H^+$) enters the bloodstream, it reacts with bicarbonate to form carbonic acid, which then decomposes into $CO_2$ and water to be exhaled by the lungs. Conversely, if the blood becomes too basic, carbonic acid dissociates to release $H^+$ ions. This is a classic application of Le Chatelier's Principle: the system shifts to counteract the change in concentration.

Environmental pH is equally delicate. Acid rain forms when sulfur dioxide ($SO_2$) and nitrogen oxides ($NO_x$) from industrial combustion react with water vapor to form strong acids like $H_2SO_4$ and $HNO_3$. While the rain itself harms leaves, the most devastating impact is underground. As soil pH drops, $H^+$ ions displace essential nutrients like $Ca^{2+}$ and $Mg^{2+}$. More dangerously, low pH causes aluminum leaching. Aluminum ions ($Al^{3+}$) are normally trapped in minerals, but in acidic conditions, they dissolve into groundwater, where they are highly toxic to fish by clogging their gills and disrupting salt balance.

The ocean absorbs about 30% of human-emitted $CO_2$. While this slows global warming, it triggers ocean acidification. When $CO_2$ dissolves in seawater, it forms carbonic acid, which releases $H^+$ ions. These $H^+$ ions then react with carbonate ions ($CO_3^{2-}$) to form bicarbonate ($HCO_3^-$). This is a major problem for 'calcifying' organisms like corals, oysters, and plankton. They need $CO_3^{2-}$ to build their calcium carbonate ($CaCO_3$) shells. As $H^+$ increases, it 'steals' the carbonate ions, making it harder for these organisms to grow and even causing existing shells to dissolve.