By David J Strumfels on
http://AMedelyofPotpourri.Blogspot.com/
(Wikipedia, 2/28/14): "Between 1751 and
1994 surface ocean pH is estimated to have decreased from
approximately 8.25 to 8.14 (how measured in 1751?),[5] representing
an increase of almost 30% in H+ ion concentration in the
world's oceans.[6][7] Available Earth System Models project that
within the last decade ocean pH exceeded historical analogs [8] and
in combination with other ocean biogeochemical changes could
undermine the functioning of marine ecosystems and many ocean goods
and services.[9]"
8.25 to 8.14 -- what does that mean? What it
does not mean is that the oceans are acidic, for their pH would have
to be below 7.0 to say that. In fact, we would call it alkaline
(the opposite of acidity). But acidification indicates a
direction, not a specific place on the pH scale (some climate
warming skeptics seem not to understand the difference) .
Chemically, pH is defined as the "negative log (hydrogen
ion concentration)". But what does THAT mean?
The hydrogen ion, or H+ (a hydrogen
atom without its electron, or a bare proton), is the main acidic
species in water. If the concentration of H+ is
8.00E-8 = 0.00000001 Molar (moles per liter, or just M), then the log
of that number is -8.00, its negative log is 8.00, and so the pH is
8.00. As to the specific cases here: pH 8.25 =>
-8.25; antilog (-8.25), or 10 to the power of the number)
yields 5.62E-9 H+ M, while 8.16 => -8.16 =>
7.14E-9 H+ M. That's a difference of 1.52E-9 H+
M concentration, or 0.00000000152M. Yes, it is a 29% increase
in acidity, mathematically. Of course, 29% of practically
nothing is even closer to actually nothing. But
chemically, especially biochemically, it can be very important.
For example, your blood pH has to be kept within a
narrow range of 8.25 and 8.35, or illness, even death, can
result. I do not know all the reasons for this, but I can tell
you that many biomolecules have both basic (opposite of acidic) and
acidic forms, and they are chemically different. They may have
to be kept within a very narrow equilibrium of basic and acidic
forms, each running a different reaction. Or some are all
necessarily basic at this pH, while others all acidic. At pHs
near 7.0 (neutrality), these equilibria are extremely
sensitive to these tiny changes I outlined above.
So it is not difficult to see how a drop in 0.09
pH units (combined with a degree or two warming) could wreak havoc on
numerous sea organisms, plant, animal, protozoan, or bacterial. On
the other hand, the immensity of the oceans assures that there will be
pH (and temperature) variations, in time and location, so sea life
should be expected to be a little more hardy. But there are still
fairly strict limits. The 21'st century will surely test those
limits.
A little more chemistry now. We blithely speak of
CO2 increasing water acidity, but how?
First, CO2 is mildly soluble in
water, the lower the water temperature, or the higher the water
pressure, the more soluble (for thermodynamic reasons we need go into
here) it is. That isn't enough for acidity, however. There has to
be a chemical reaction between CO2 and
H2O
first: CO2 +
H2O
<=> HCO3-
and H+.
The <=> sign means the reaction goes in both directions; which
set of reactants/products is favored depends on
various conditions. Cold leans toward the first
two, pressure the opposite way. This again is thermodynamics, with
enthalpy and entropy competing against each other. At the low
concentrations of CO2,
in general the latter is favored. But it's still a tiny contribution
of H+, enough, as
you've seen, to lower ocean water an average of 0.1 pH units over the
last two hundred years (and I would not be surprised if 0.5 – 0.7
units of that is within the last 30-40 years, and
it ~doubles by 2100 – this is serious).
