133
Article
pH, the CO2 system and freshwater science
J.F. Talling
Hawthorn View, The Pines, Bongate, Appleby, Cumbria CA16 6HR, UK. Email: [email protected]
Received 28 January 2010; accepted 21 April 2010; published 3 August 2010
Abstract
Appreciation of pH as an ecological index has varied considerably over its past history,
influenced by perceptions of chemical rigour, ease or difficulty of measurement, and multiple
chemical and biological correlations. These factors, and especially the last, are considered in
relation to the extensive range of pH in inland waters. Emphasis is placed upon the role of the
CO2 system, the components of which are subject to biological metabolism (photosynthesis,
respiration) and are extensively determined by products from rocks (e.g. limestone) and soils.
Titration alkalinity, or acid neutralising capacity, is a most valuable summarising and reference
measure. For this, and CO2 variables, potentiometric Gran titration opens new possibilities – including
the definition of negative alkalinity (acidity). The relationship of pH and titration alkalinity is close
and semi-logarithmic for waters in equilibrium with atmospheric CO2. Very high pH, above 10, can
develop from the photosynthetic depletion of CO2 and by the evaporative concentration of bicarbonatecarbonate waters in closed basins. Very low pH, below 4.5, results from the introduction of strong
acids by volcanic emissions, pyrite oxidation, ‘acid rain’ and cation exchange; here the CO2 system
lacks influence, biological diversity is reduced, and ionic aluminium often exerts toxic biological effects.
Situations of pH excursion are discussed and illustrated; they operate over day–night, seasonal and longterm time-scales. A summer rise of pH is widespread in productive near-surface waters. There is also a
seasonal pH rise in the anoxic deep water of many lakes, as a consequence of the interaction of acid–base
and oxidation-reduction systems. These can be regarded as two ‘master systems’ of environmental
chemistry and – dating from pioneer studies of wetland soils and waters – of much freshwater ecology.
Keywords: acidifuge; alkalinity; acid neutralising capacity; bicarbonate; buffering; carbonate; carbon dioxide;
hydrogen ion; pH; potentiometric titration; redox.
DOI: 10.1608/FRJ-3.2.156
Freshwater Reviews (2010) 3, pp. 133-146
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134
Talling, J.F.
Introduction
basis introduces a small temperature-dependency via water
density, and a larger one comes from the temperature-
No chemical concept, dependent upon a mathematical
sensitivity of the dissociation constant of water (Kw). The
twist, has been received more enthusiastically by
latter, like [H+] concentration, can be represented by its
ecologists than pH. Few chemists are aware that its value
negative logarithm pKw. Consequently, as Hutchinson
as a convenient variable was originally promoted by a
(1957) explains, the pH of pure water is estimated as 7.00
botanist, the Dane Sørensen. It is generally perceived as
at 25 °C, but 6.77 at 40 °C. In natural waters other pK
the exponent expressing the logarithm (base 10) of the
factors also determine a variable temperature-dependence
hydrogen ion concentration in moles per litre, with its sign
of pH, generally of the order of -0.1 unit per 10 °C rise.
reversed. If heavy isotopes are discounted, the hydrogen
Here emphasis is given to the interrelationship of many
ion is an elementary particle, the proton. The chemical
components, illustrated mainly by examples from my own
species involved in its exchange, as donors or acceptors of
experience.
protons, are protolytes.
This
simplicity
neglects
two
considerations:
quantitatively, the chemically effective quantity is related to
Ecological appreciation and
prejudice
activity rather than concentration; qualitatively, the hydrogen
ion H+ in aqueous solution is largely associated with a water
Botanical influence behind pH perhaps presaged a
molecule H2O as the hydroxonium ion, H3O+. The volumetric
welcome among biologists and ecologists of a quantity
regarded
as
and
rigorous
fundamentally
entrenched in chemistry.
It was early perceived as
having wide ecological
associations
and
this
generated a less than
critical enthusiasm. An
anecdote of this era – that
of Shelford- influenced
ecology – is recounted
by Allee et al. (1949). It
centred on a supposed
useful supremacy of pHmeasuring
equipment
among the apparatus of
ecologists on a small boat
greeted in an American
harbour. It will later be
shown
that
although
associations are indeed
Fig. 1. Examples of the depth-distribution of pH measured in two English lakes, dissimilar in depth
and planktonic production, illustrating the range encountered seasonally under varied conditions
of stratification and algal abundance. Small arrows indicate the approximate pH of surface water at
atmospheric equilibrium. For one depth-profile, estimates are also given of total CO2 (Ct) and its relation to
alkalinity (A). From Talling (2006 Fig. 6).
© Freshwater Biological Association 2010
numerous, they are often
the result of indirect and
non-rigorous correlation.
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pH, the CO2 system and freshwater science
A later swing to a more critical – at times over-critical
circum-neutral salts (Sillén, 1961), the free forms of strong
– approach prevailed in the 1950s and 1960s. This centred
acids and strong bases are insignificant. Instead, pH is
on imperfections of measuring methods and equipment –
usually determined by the ionised forms and derivatives
including colorimetric standards unsupported by standard
of weak acids and weak bases. Here the concentrations,
buffers, and ‘pH papers’ with appreciable intrinsic
and so activities, of components of the CO2 -system related
buffering. The widely used measurement via electrical
to H2CO3 generally predominate:
potential, with a reference electrode, suitable offsets and
scaling, usually depended on a glass electrode whose high
CO2 + H2O = H2CO3
resistance introduced problems of electrical measurement
(1)
– especially in more dilute natural waters. Examples,
H2CO3 = H+ + HCO3-
instanced by Schindler, occurred in early work on the
(2)
Experimental Lakes on the Precambrian Shield of Canada.
HCO3- = H+ + CO32-
There was similar experience in the English Lake District;
this led to certain astonishingly high records of pH obtained
(3)
with the dissociation of water
by a bacteriologist being discounted by some chemical
and botanical colleagues. They were later found to have a
H2O = H+ + OH-
basis in reality (Talling, 1976, 1985; Maberly, 1996; examples
(4)
in Fig. 1) among some of the most remarkable chemical
and sometimes gain or loss of CO32- and HCO3- to or from
behaviour of the English Lakes! Problems of potentiometric
calcium carbonate of limited solubility.
measurement in this phase were reduced by improvement
in measuring equipment. Nevertheless, a really critical
The importance of this system is contributed by
1.
recognition of potential errors in pH in electrometric
measurement, as by Covington et al. (1985a, b) and
a metabolic input and product
2.
Davison (1990), was – and is – often not applied in practice.
The chemical significance of measurements to hundredths
the abundance of sedimentary carbonates – potentially
soluble in water – in the earth’s crust
3.
of a pH unit – electrically feasible – may be uncertain to
tenths of a unit. Nor is it generally appreciated that pH
the major involvement of CO2 – a pH depressant – as
the existence of a major atmospheric reservoir of
gaseous CO2
4.
the magnitude of relevant dissociation constants (K),
is a temperature-dependent quantity, as is implied by the
yielding pK values and hence large changes in the
temperature-dependence of relevant dissociation constants,
chemical species within the common pH range of 4
and that it can be altered by exposure of a water sample
to 10. In ascending this span, the predominant form
to atmospheric CO2. Precautions when sampling may
of dissolved inorganic carbon changes from CO2 to
therefore be advisable (Mackereth et al., 1989). Continuous
HCO3– to CO32-, as shown in the relative Buch diagram
recording in situ is also possible (e.g. Maberly, 1996).
reproduced in many textbooks.
Outside the CO2-system, there are in natural waters
Some relevant chemical
background
other weak inorganic acids – notably boric acid and
silicic acid – whose molecules and their ionic derivatives
act as proton-donors and acceptors and so influence pH.
In ascending the pH scale from 0 to 14, one moves from
Concentrations of borate are significant in sea water (e.g.
an acidic region of predominant proton (H ) donors to one
Hill, 1963). Those of silicic acid – Si(OH)4 , for which the
of predominant proton acceptors. In most natural waters,
substitution of fictitious SiO2 is unfortunately common – are
and notably present-day sea water after its remarkably
appreciable in most fresh waters and exceed 0.5 mmol L-1
near-balanced evolutionary interaction or ‘titration’ to
in many tropical examples. However, its ionisation – first
+
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Freshwater Reviews (2010) 3, pp. 133-146
136
Talling, J.F.
to SiO(OH)3- – is appreciable only above pH 9. It has been
titration (Gran titration: Gran, 1952), originally applied to
usual to neglect an influence of this, and of organic acids at
sea water (Dyrssen, 1965; Dyrssen & Sillén, 1967; Edmond,
lower pH values, although detectable effects can exist (e.g.
1970) and then introduced to fresh waters (Talling, 1973).
Talling, 1973) and need more study. At the other end of the
It involves an antilogarithmic conversion of pH values
pH scale, below pH 4.5, there is a potential mobilisation of
met during the titration.
ionic species of aluminium that influence pH and exert a
– one in a higher band of pH – can in a closed system
biological toxicity (see, e.g., contributions in Mason, 1990).
yield a measure of total (i.e. free + ionic) CO2 between
Besides the CO2 system, one should recognise the
the endpoints which they indicate (Edmond, 1970).
Two successive conversions
following chemical agencies as influencing pH in natural
Several chemical species are potentially involved in
waters. First, there is the introduction of strong acids,
the acid neutralising capacity. Normally the bicarbonate
most universally in atmospheric precipitation bearing
and carbonate ions are far predominant, but formally
products of combustion (especially compounds of
allowance is required for presence in the samples of
sulphur and nitrogen). In this way the pH of rainwater
pH-dependent OH- and H+ ions and possibly anions of
is often reduced from c. 5.0 to c. 4.3. Also widespread is
other weak acids (A-) such as SiO(OH)3- from silicic acid:
a mineral origin of sulphuric acid, as by the oxidation of
pyrite FeS2. Examples are described in Geller et al. (1998).
alkalinity (orANC) = [HCO3-] + 2 [CO32-] + [OH-] + [A-] - [H+]
Second, there are the less readily resolved effects of weak
(5)
organic acids (discussed in Perdue & Gjessing, 1990 and by
It is consequently a mixed sum of concentrations that is best
Tipping, 2002) that often originate in the catchment. Third,
expressed (as above) in units of milli-equivalents or micro-
there is ion-exchange involving base cations. Fourth, there
equivalents L-1 (displaced in much modern chemistry)
can be influence from co-existing oxidation-reduction (redox)
rather than the corresponding molar units (represented by
reactions in the aquatic medium. In the wider sense, these
terms in brackets above) – although components other than
can include aspects of carbon and nitrogen metabolism.
HCO3- are typically negligible over the common pH range
Broad quantitative discussion of factor interaction, as
of 6.0 to 8.5. All these component quantities, like those of
affecting pH in lakes, is given by Psenner & Catalan (1994).
‘free CO2’ that conventionally deals with the aggregate CO2
A measure of capacity in proton exchange has long been
+ H2CO3 (denoted H2 CO3*), tend to change logarithmically
used in limnology. It was originally known as one form of
over the pH scale whether expressed in equivalents or
titration alkalinity, and latterly also (and more expressively)
molar units. The latter are shown in Fig. 2 for a water
as acid neutralisation capacity (ANC) or the German Säure-
of alkalinity 0.40 meq L-1, similar to that of the English
bindungvermögen (SBV).
lake Esthwaite Water, with pH varied by the addition or
It represents the capacity to
neutralise strong acid added to the bicarbonate endpoint,
depletion of CO2.
and can be expressed in milli- or micro-equivalents (meq
Alkalinity also sets an upper (and temperature-sensi-
or µeq) per L. Strictly it is a difference, not a concentration,
tive) limit to pH-rise induced by the removal of CO2. If that
and has negative values in acidic solutions. A former
is complete, and the quantity A- is small or negligible as
and misleading expression used mg L-1 units of fictitious
often assumed, then alkalinity is largely made up of OH-
CaCO3 (1 meq L = 50.05 mg L CaCO3). This alkalinity
and approximates its concentration, that in turn – by the
is not to be confused with an alkaline reaction as high
known dissociation constant of water Kw – determines pH.
pH. The endpoint was classically approximated using
At this point all CO2- related quantities – free CO2, HCO3-,
methyl orange indicator, and later mixed indicators, with
CO32- – are extinguished. This situation is shown in Fig. 2
endpoint at pH 5 approximately; a small correction of c.
and was approached in some experimental exposures,
0.01 meq L was often applicable (e.g. Sutcliffe et al., 1982).
‘pH-drifts’, with the cyanoprokaryote (cyanobacterium,
More accurate determination is possible by potentiometric
cyanophyte, ‘blue-green’) Microcystis aeruginosa as an
-1
-1
-1
© Freshwater Biological Association 2010
DOI: 10.1608/FRJ-3.2.156
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pH, the CO2 system and freshwater science
effective photosynthetic remover of CO2 (Talling, 1976).
It was found to be low in Aulacoseira subarctica Haworth
Upper limiting values of pH, so engendered, could exceed
(formerly Melosira italica
11.0, but – as also found by Maberly et al. (2009) for less
(Talling, 1976), and may contribute to growth restriction of
effective CO2 removers limited by H2CO3* depletion –
the related A. granulata at elevated pH (9.0+) generated by
varied with the initial alkalinity and so concentrations of
growth of phytoplankton in the Nile (Talling et al., 2009).
exchangeable CO2. They are also markedly temperature-
In my opinion, alkalinity is the best single measure
dependent. This has implications for an assessment of
for interpreting the wider ‘ecological’ meaning of a pH
accentuated limiting pH in warm tropical waters, for which
measurement on a water sample. In part this is because
a high pH implies more CO2-depletion than in colder waters.
alkalinity is an unaltered or ‘conservative’ property
The capacity for removal, and hence CO2 depletion with
regarding the addition or the removal of CO2, which
raised pH, varies widely in different groups and species
respectively depress or raise pH in natural waters subject
of algae and macrophytes. Differences were formerly
to biological activity. It could, for example, resolve the
ascribed to the presence or absence of an ability to directly
meaning of a high pH such as 10, which might result from
utilise bicarbonate as C-source (e.g. Ruttner, 1947), later
CO2 depletion in a productive water of low alkalinity or
linked more widely to the possession of carbon-concen-
from high alkalinity in a saline soda lake. It is temperature-
trating mechanisms (e.g. Maberly & Madsen, 2002). There
independent (cf. Dyrssen & Sillén, 1967). It also represents
is limited capacity in chrysophytes, related to their absence
ions, and especially bicarbonate, that make up most of the
from the more alkaline waters (Saxby-Rouen et al., 1999)
total anion concentration of many inland waters. These
and the probable lack of a carbon-concentrating mechanism
and other correlations, soon to be discussed, once led an
(Maberly et al., 2009). Diatoms show very variable capacity.
experienced limnologist (C.H. Mortimer) to remark in my
subsp. subarctica O. Müll.)
hearing that if he had to choose a
single chemical measure of a fresh
water he would take alkalinity.
Buffer intensity (β) is another
chemical property of note.
It
indicates the resistance to pH
change induced by unit addition of
CO2, acid, or alkali. The first is most
prevalent in ecological situations,
where buffer intensity with respect
to CO2 change can be calculated
(e.g. in mmol CO2 per pH unit) from
parameters of the CO2 system (e.g.
Talling, 1973). In this system, buffer
intensity varies very non-linearly
with pH, and is low in a pH region
– slightly temperature-dependent
around pH 8.1 to pH 8.3 – where
induced molar changes in HCO3Fig. 2. Logarithmic and temperature-sensitive distributions with pH, varied by CO2 change, of
chemical components active in operation of the CO2 system. This system is modelled for a fresh
water of low ionic strength (0.001), alkalinity (0.40 meq L-1) and non-carbonate alkalinity (A-),
showing extinction of CO2-system components (free CO2 = H2CO3*, HCO3-, CO32-) at high pH
with OH- domination of alkalinity. Modified from Talling (1985 Fig. 1).
DOI: 10.1608/FRJ-3.2.156
and CO32- concentrations over the
pH scale are minimal. At much
lower and higher pH it is raised
Freshwater Reviews (2010) 3, pp. 133-146
138
Talling, J.F.
by exponential increase in H+ and OH- concentrations
assessing aquatic photosynthesis and production (Verduin,
and possibly by other chemical species, such as ionic
1956; Byers & Odum, 1959; Byers, 1970). Photosynthesis
forms of aluminium below pH 5. Saline inland waters of
follows the day–night (diel) cycle of solar radiation, and
low alkalinity exist but are not necessarily more strongly
is often very localised over depth. Diel cycles of pH,
buffered than much more dilute fresh waters. Waters of
elevated by day, are consequently near-universal in waters
moderate alkalinity (e.g. 1 – 2 meq L -1) tend to have pH
productive of aquatic macrophytes or phytoplankton.
values near the point of minimum buffer intensity, and are
The latter situation is illustrated in Fig. 3, which shows the
in consequence prone to upward shift of pH when CO2
distribution of temperature, dissolved oxygen and pH in
is consumed in photosynthesis, and fall of pH when CO2
a Nile reservoir that is seasonally rich in phytoplankton.
accumulation results from decomposition and respiration.
Here the density stratification imposed by daytime rise of
Incidentally, animal ecologists have only rarely taken
water temperature stamps similar depth-time patterns on
episodes of the former rise into account (as by Dunn, 1967).
the chemical variables. The daytime or diurnal rise of pH
Salinity has but little direct influence on pH insofar
is reversed by a largely nocturnal re-distribution of CO2
as it is mainly contributed by neutral salts of strong acids
from depth and net production in respiration, and by its
and bases. Sea water, with high concentrations of Na and
slow re-entry from the atmospheric reservoir. In general,
+
Cl-, is the commonest example. Here, as in
most fresh waters, the magnitude of titration
alkalinity (circa 2.3 meq L-1 in oceanic water)
is of most significance.
However, salinity
has a small systematic effect by altering ionic
strength, a characteristic that widens the gap
between concentration and chemical activity
of dissolved constituents; it thereby alters the
values of effective or apparent dissociation
constants (K’), reducing their pK’ (negative
logarithmic) values.
Means of calculating
this effect on CO2-related quantities in
fresh waters are set out by Mackereth et al.
(1989). In general, critical treatments of the
variable CO2 system in fresh waters (e.g.
Howard et al., 1984) have been fewer than
the corresponding studies on sea water.
Excursions of pH over time
The susceptibility of pH to free CO2 content
is obviously liable to promote its changes
in time, as natural waters are subject to
biological influence from CO2 consumption
(photosynthesis)
and
CO2
production
Fig. 3. Day-night cycles of stratification during a phase of phytoplankton abundance
(respiration, decomposition). In fact, changes in a Nile reservoir, 6–8 October 1955, showing (a) incident photosynthetically active
in pH were formerly used as a method for solar radiation Io (W m-2) and the depth-distribution of (b) temperature, (c) dissolved
oxygen, (d) pH. Adapted from Talling (1957 Fig. 4).
© Freshwater Biological Association 2010
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pH, the CO2 system and freshwater science
the rate of such re-entry is limited by the small atmospheric
accumulation of CO2. Even wider changes (e.g. with
proportion, and hence partial pressure, of CO2, although
pH range 10.35 to 6.47 in 1971–72) occur in the more
there is stimulation by chemical reactivity to waters above
productive Esthwaite Water, but here another chemical
pH 9 (Emerson, 1975). For these situations of diurnal
system – of oxidation-reduction or redox – also intervenes
excursion involving pH, the limitations of atmospheric
to modify pH in deep water. There the deep pH decline
replenishment are clear. There are relatively few examples
is reversed to pH rise in late summer, and is accompanied
in which the changing content of total CO2 has been
– and partly induced – by a major increase in alkalinity.
measured by direct chemical methods (e.g. Schindler &
These extra changes are a consequence of final de-
Fee, 1973).
oxygenation at depth, that induces several chemical
The annual cycle in most fresh waters involves
reductions (e.g. Fe3+ to Fe2+) in which electron transfer also
changes in the same variables. Those of pH are again most
leads to proton transfer (Heaney et al., 1986; Talling, 2006),
generally linked with net consumption and production
and previously free CO2 is incorporated in bicarbonate.
of CO2; elevated pH may be correlated with a lower rate
Bacterial reductions are implicated (Kelly et al., 1982).
of photosynthesis per unit biomass (e.g. in Loch Leven:
Interaction between the two systems also influences
Bindloss, 1974). There are often depth-differences of pH
reaction kinetics, as is manifest in strong pH-sensitivity
associated with temperature-density stratification and
of the rate of Fe2+ oxidation (Davison & Seed, 1983).
the distribution of biological activity. Figure 4 illustrates
Some other pH trends, often longer-term, do not
such patterns of seasonal change in two English lakes.
depend on changes in CO2. Some aquatic macrophytes,
Windermere South Basin shows, in summer, elevation
notably species of the bog-moss Sphagnum, induce
of pH in the upper layers and its depression in the lower
acidification by processes of ion-exchange (e.g. Clymo,
– which respectively correspond to net depletion and
1963). Inputs of strong acids – including H2SO4 – from
‘acid rain’ influenced by industry, are another
now-familiar source of long-term acidification
(e.g. Mason, 1990).
Prospects of freshwater
susceptibility, and recovery (e.g. Tipping et
al., 1998), depend critically upon sources of
alkalinity to a water-body from the catchment
and in the water-body itself. The latter may
exceed the former (Psenner & Catalan, 1994).
The internally derived alkalinity may modify
biological production, and be itself augmented
by experimental stimulation of that production
(e.g. Kelly et al., 1982; George & Davison, 1998).
Ecological and environmental
correlations
Acidophile and acidophobe species have long
been distinguished by freshwater and terrestrial
Fig. 4. Seasonal changes during 1967 in the magnitude and depth-distribution
of temperature, dissolved oxygen and pH in two English lake basins, illustrating
upward shifts of pH due to photosynthetic CO2 depletion above and interaction
with redox shifts induced by anoxia below. Adapted from Talling (2006 Fig. 4).
DOI: 10.1608/FRJ-3.2.156
ecologists, from floristic and faunistic surveys.
In fresh waters the species itself rarely generates
the associated condition, although abundant
Freshwater Reviews (2010) 3, pp. 133-146
140
Talling, J.F.
Fig. 5. The relation with alkalinity of pH (measured at 15–25 °C) in surface samples of lake waters subjected to air-equilibrium. Superimposed
are natural ranges of pH recorded in some of the lakes. Adapted from Talling (1985 Fig. 4).
phytoplankters and acidophile species of Sphagnum
(already mentioned) are conspicuous exceptions. More
pH equilibrium = 11.3 + log [HCO3-]
= 11.3 + log [alkalinity]
usually the pH condition arises passively, and often
(6)
from other chemical correlations of pH. One is that with
for the pH range 6 to 9 and alkalinity (measured in eq L-1)
calcicole and calcifuge species influenced by exposure
below 10-2 eq L-1 (i.e. 10 meq L-1). Temperature-influence
to calcium carbonate, abundant as marine residues in
exists but is minor; it has effects on the solubility of
limestone and chalk and subject to weathering to yield
gaseous CO2 (Bohr coefficient) and on relevant dissociation
waters with elevated levels of both calcium and alkalinity.
constants (K values). A recent long-term trend to higher
These waters are generally of moderately high pH as pH
concentrations of atmospheric CO2 is not without a small
rises near-linearly with the logarithm of alkalinity when
but significant effect on the pH of oceanic waters (Raven,
the concentration of free CO2 is near that at equilibrium
2005). Altitude is influential via atmospheric pressure and
with that of the atmospheric reservoir. An approximation
hence the partial pressure of atmospheric CO2.
under current conditions, excluding saline waters and
Figure 5 illustrates the semi-logarithmic relationship
those at high altitude, is derived by Psenner & Catalan
for a wide range of lake waters, with some large individual
(1994) as
deviations from linearity mainly associated with extreme
© Freshwater Biological Association 2010
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pH, the CO2 system and freshwater science
natural excursions of free CO2 concentration. Remarkably
pH. However, high pH induced by CO2 uptake in low
Windermere, a lake of low alkalinity, has been found to
alkalinity waters is not excluded; evidence includes that
have occasionally a pH of more than 10 in surface water.
of Reynolds (1986) for Staurastrum pingue above pH 10.0 in
Although pH elevation by CO2 depletion is here more
an experimental enclosure and a maximum of S. cingulum
extensive than pH depression by CO2 accumulation, the
at pH 10.2 in Windermere South Basin (Maberly, personal
latter condition is more predominant year-round in this
communication). Physiological work (Spijkerman et al.,
and many other lakes, and most running waters (e.g.
2005) has shown that planktonic desmids include species
Rebsdorf et al., 1991), with implications for the overall
with diverse capacity for CO2 uptake, not all restricted
balance of metabolism affected by external inputs of
to free CO2 as the source. Examples of alkalinity as a
sources of CO2 (e.g. Cole et al., 1994; Maberly, 1996). Higher
prime factor in the distribution of diatoms have been
alkalinity, shown in Fig. 5 from selected African lakes, is
demonstrated by Hustedt (1949), Richardson (1968),
there closely correlated with salinity as the HCO3- + CO32-
Gasse et al. (1983) and Wood & Talling (1988) in African
+
pair generally predominates among anions, as does Na
lakes, and by Knudson (1954) in the English Lakes. Still
among cations. Consequently evaporation in closed basins
more widely studied have been the pH-preferences
results in a range of saline ‘soda lakes’, of high pH and low
of diatoms.
content of Ca2+ and Mg2+ (Talling & Talling, 1965; Wood &
recoverable indicator species and associations, widely
Talling, 1988). Such development to high alkalinity, and
used to trace past conditions on the pH scale (e.g. Battarbee
correlatively high pH, would be impossible from CaCO3
et al., 1990) rather than the correlated scale of alkalinity.
-based marl lakes – water-bodies recently reviewed by
Exceptionally, both scales have been used (Catalan et al.,
Pentecost (2009). In these the low solubility of CaCO3 is
2009). On the short time-scale of a few hours, population
combined with its frequent deposition, in part biogenic.
activity and growth can be extinguished by CO2 depletion
One such lake, the upland Malham Tarn in Yorkshire,
rather than correlated increase of pH – as was evident in
UK, had in 1985–87 a seasonal range of alkalinity of 1.0 to
pH-drift experiments involving pH shifts recorded in
2.4 meq L-1and, in antiphase, a pH range in surface water
illuminated algal suspensions with variable alkalinity
Diatoms have been a popular source of
of 7.7 to 8.7 (Talling & Parker, 2003). Much of the
variation of alkalinity was caused by the seasonal
abstraction of CaCO3 by a stonewort (Chara sp.;
Pentecost, 1984). Here and beside Sunbiggin Tarn,
another upland water in Cumbria (Holdgate,
1955), there is wide local variation of calcicole and
calcifuge plants in peripheral waters and soils,
influenced by adjacent limestone and upward
building of acid raised bog.
These influences
apart, soil-water typically has a high content of
free CO2, which can fall by an order of magnitude
upon exit to surface waters with opportunity for
atmospheric equilibration and associated rise in pH.
Both alkalinity and pH are characteristics that
often have close correlations with the distribution
of species of algae, most notably desmids and Fig. 6. Continuous record of a ‘pH drift’ over time during photosynthetic
diatoms. Desmids often reach remarkably high
diversities in waters of low alkalinity and low
DOI: 10.1608/FRJ-3.2.156
uptake of CO2 by a dense suspension of the dinoflagellate Ceratium hirundinella,
in an initial lake-water medium with alkalinity values (inserted above, in µeq
L-1) successively altered by additions of HCl and Na2CO3. Adapted from
Talling (1976 Fig. 17c).
Freshwater Reviews (2010) 3, pp. 133-146
142
Talling, J.F.
northern Europe is divisible into
two groups, separable by their
response to prolonged exposure to
pH less than 5.7’. This is broadly
borne out by the natural distribution
with pH of numerous species
surveyed by Psenner & Catalan
(1994), who suggest that there is a
corresponding minimum HCO3requirement of about 20 µeq L-1.
Additionally, tolerance of low pH
appears to be reduced in waters of
low salt content via adverse balance
between ion (e.g. Na+, K+) loss
and uptake. Examples appear in
experimental work on crustaceans
(Potts & Fryer, 1979) and in patterns
of their geographical distribution in
Yorkshire. In this topographically
varied county there is a strongly
bimodal frequency distribution
of pH in standing fresh waters,
with maxima near pH 7.6 (mostly
lowland) and 4.0 (mostly upland),
the latter marked by a reduced
occurrence
Fig. 7. The pH scale, showing distribution of components and situations discussed in the text.
of
many
species.
Nevertheless, there is an appreciable
occurrence of small crustaceans
(Fig. 6), performed by Talling (1976) and Maberly et al.
– including the near-omnipresent Chydorus sphaericus
(2009). On the longer time-scale of population dynamics in
– in very acid waters of pH 3 to 4 (Fryer, 1980, 1993a, b).
nature, the eligibility of pH as a density-dependent factor
Another factor behind the elimination of acidifuges is
or factor-associate has scarcely been examined rigorously.
susceptibility to toxicity of ionic aluminium, especially Al3+,
Possible evidence is reviewed by Shapiro (1990),
abruptly mobilised from other insoluble forms at pH below
including the frequent association of bloom-forming
5 with zero to negative alkalinity. This additional toxicity
cyanoprokaryotes (‘blue-greens’) and elevated pH.
factor at low pH, ‘beyond the CO2 system’, has recently
Both these components are favoured – independently
engaged much attention in relation to episodic and long-term
of physiology – by shallow upper mixed layers.
acidification of surface waters (examples in Mason, 1990).
Turning to freshwater invertebrates and fishes, and
Figure 7 summarises the incidence of these and other
passing down the pH scale, there are many species
transitions on the overall pH scale. Looking at the pH 0–14
recognised as acidifuge which cease to occur before the
scale as a whole, there is some symmetry of opposites in
point of zero alkalinity is reached. Thus Sutcliffe (1978)
the associated rise of [H+] or [OH-] concentrations at the
commented that ‘in general the freshwater fauna of
two ends – exceeding 1 mmol L-1 at approximately pH 3
© Freshwater Biological Association 2010
DOI: 10.1608/FRJ-3.2.156
143
pH, the CO2 system and freshwater science
in base-poor (Mortimer, 1941–42; Drake & Heaney, 1987)
and base-rich (Golterman, 1984) situations.
It seems to me that two developments in the 1930s
were especially notable. First, there was the systematic
study by Jenkin (1936) of a series of Kenyan rift lakes of
high-ascending alkalinity and pH. This benefitted from the
influence of earlier basic work at Cambridge by Saunders
(1926) on the pH–CO2 system of relationships and of his
and Jenkin’s work on the ciliate Spirostomum (Jenkin, 1927).
Fig. 8. W.H. Pearsall making electrometric measurements during
the late-1930s beside the reedswamp fringing an English lake.
Photo. C.H. Mortimer.
These investigations once led G.E. Hutchinson (formerly a
pupil of Saunders) to cite this ciliate as providing perhaps
the best established instance of a pH relationship in aquatic
and pH 11 respectively. Asymmetry in natural waters
ecology. Second, co-action between the acid–base and
appears in the main influence of labile CO2 exchange over
oxidation-reduction (redox) systems were investigated
a largely alkaline to slightly acidic band, and of cation
by Pearsall (1938) in relation to soil, fen and bog water
exchange over a decidedly acid band. Other than in these
(see Fig. 8) and extended to lake sediments by him and
processes, input of strong acid dominates the lower end,
Mortimer (Pearsall & Mortimer, 1939; Mortimer, 1941–42).
and evaporative concentration of pre-existing carbonate
Later Sillén (1967) and others considered the two realms
the upper, with typically long-persistent consequences.
centring on proton and electron transfers, here portrayed
General comments
The historical after-view would suggest
that an original hope, that pH might
represent a uniquely rewarding factor
embracing many ecological relationships,
has been fulfilled only partially. One reason
(both for and against) is the multiplicity
of variably cross-correlated properties
(e.g. pH, alkalinity, Ca2+); another is the
influence,
and
frequent
dominance,
of the relatively labile CO2 system of
relationships. In the words of Fryer (1993a),
‘just as pH is no philosopher’s stone, even
a complex of several physico-chemical
factors may not fulfil this function either’.
One relation to habitat productivity might
be induced by photosynthesis; another by
ion-speciation – e.g. of H2S–HS-–S2-, and
of H2PO4-–HPO42-–PO43- – sensitive to pH
and influencing the mobility of phosphate
DOI: 10.1608/FRJ-3.2.156
Fig. 9. Diagrammatic representation of potential chemical transfers associated with
the acid–base and oxidation-reduction (redox) systems, their interaction and ultimate
atmospheric regulation, in a stratified lake subject to deep anoxia. Adapted from Talling
(2006 Fig. 10).
Freshwater Reviews (2010) 3, pp. 133-146
144
Talling, J.F.
within a lake setting in Fig. 9, as major ‘master systems’ (i.e.
strength solutions including freshwater. The Analyst 108, 1528-
primary systems with multiple consequences approachable
1532.
by quantitative deductions; e.g. Stumm & Morgan, 1970).
Covington, A.K., Whalley, P.D. & Davison, W. (1985b).
These master systems not only underlie aquatic chemistry
Recommendations for the determination of pH in low ionic
but also – one might add – much aquatic ecology.
strength fresh waters. Pure and Applied Chemistry 57, 877-886.
Davison, W. (1990). A practical guide to pH measurement in
Acknowledgments
freshwaters. Trends in Analytical Chemistry 9, 80-83.
Davison, W. & Seed, G. (1983). The kinetics of the oxidation of
This review has benefited by helpful comment and
ferrous iron in synthetic and natural waters. Geochimica et
literature from Geoffrey Fryer FRS, John Raven FRS and
Cosmochimica Acta 47, 67-79.
Stephen Maberly, and the thoughtful comments of two
Drake, J.C. & Heaney, S.I. (1987). Occurrence of phosphorus and its
anonymous reviewers. Figures 1, 4 and 9 are reproduced
potential remobilization in the littoral sediments of a productive
by kind permission of Springer Science & Business Media,
Dordrecht.
English lake. Freshwater Biology 17, 513-523.
Dunn, I.G. (1967).
Diurnal fluctuations of physicochemical
conditions in a shallow tropical pond.
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Author Profile
Jack Talling has been engaged, since 1950, in studies
on the physiological ecology of freshwater algae
and the associated physical and chemical behaviour
of lakes and rivers. A special interest in tropical
limnology developed from early experience in Africa
at Khartoum and in East Africa. All these research
topics were maintained during work over more than
50 years with the Freshwater Biological Association
at Windermere, including a late period there as an
Honorary Research Fellow.
DOI: 10.1608/FRJ-3.2.156