Buffers for duffers …

In Ecology in the hard rock café I wrote about the challenges of living in an aquatic world where carbon – one of the raw materials for photosynthesis – was in short supply.   What I did not write about in that post is that this carbon also gives freshwater some useful additional properties.   In brief, rainwater is not pure water, but absorbs carbon dioxide from the atmosphere.  This, in turn, makes rainwater slightly acidic and, when it falls onto rocks, this weak acid dissolves the minerals from which the rock is made.  This adds two other forms of carbon to the water – bicarbonate and carbonate (the latter, particularly, from limestone).

Each of these three types of carbon in freshwater can convert to either of the other two types, with the speed of the reaction depending on the balance between the forms (the “law of mass actions”).  In essence, the reactions proceed until equilibrium is obtained, and this equilibrium, in turn, depends upon the pH of the solution.  These processes are summarised in the diagram below.

Relationship between pH and the proportion of inorganic carbon as free carbon dioxide (or carbonic acid, H2CO3 – orange line), bicarbonate (HCO3 – green line) and carbonate (CO32- – blue line).

The chemistry behind this is not easy to explain but a consequence is that any attempt to shift the pH (e.g. by adding acid) causes an automatic adjustment in the balance between the different forms of carbon.  Some of the hydrogen ions that could make the water acid are, instead , bound up as bicarbonate, and the pH, as a result, does not change.  The greater the quantity of inorganic carbon in the sample, in other words, the greater the capacity of the water to resist changes in pH.   The carbonate, bicarbonate and free carbon dioxide together act as a “buffer”, a chemical shock absorber.   Think of it as equivalent to the responsible use of a credit card or savings account to defer the cost of an unexpected bill (a car repair, for example) so that your current account does not go overdrawn.

Because life largely evolved in well-buffered marine systems, the enzymes that run our cells generally work best within a narrow range of pH (approximately 6-9).   Cells – unicellular life forms in particular – get stressed if pH strays outside this range, so the greater the buffering capacity, the easier it is for cells (life at high pH can bring additional complications, but we don’t have time to go into those here).  “Alkalinity”, as I mentioned in the earlier post, is the measure that ecologists use to assess the strength of the buffer system in a lake or river.  The principle of the measurement is straightforward: we add a dilute acid very slowly and watch what happens to the pH.   At first, nothing happens but, as soon as the water’s natural buffering capacity has been exceeded, pH drops rapidly.

I have a small portable alkalinity titration kit which involves adding drops of bromophenol blue indicator to a sample of stream or lake water.  This gives the water a blue colour when the pH is greater than 4.6.  As the pH falls, the solution becomes colourless and, eventually, turns yellow.   If you look at the graph above you will see that, at pH 4.6 most of the bicarbonate (HCO3) has been converted to carbon dioxide so the buffering capacity is pretty much non-existent.  This means that I can use the quantity of acid that is needed to make the bromophenol blue change colour as a measure of the buffering capacity of the water.

Alkalinity titrations beside Ennerdale Water (see top photograph) using a Hanna HI 3811 alkalinity test kit.  The right hand image shows acid being added to the water sample with a 1 ml pipette.  The blue colour shows that pH has not yet dropped below 4.6.

All this talk of chemical equilibria seems to be a long way from the natural history that is the core business of this blog.  Yet, at the same time, these reactions describe natural phenomena every bit as real as the plants and animals that attract the interest of naturalists.   Geology and chemistry ultimately create the context within which biology flourishes, but it is rare to meet a chemist who can talk with a naturalist’s passion.  I think that this is partly because chemistry tends not to describe tangible features of the landscape but, instead, quickly gets lost in abstract equations.  However, it is also a matter of culture: chemists need clinical separation from the mud and filth to maximise precision, whilst ecologists feel the lure of the field.  There is, nonetheless, a very basic and necessary link between the chemistry and ecology of aquatic systems.   Geology may shape a landscape but chemistry is one of the key mediators that determines the types of plants that cloak the hills and vales.  We ignore it at our peril.

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