What is pH?
pH is a measure of acidity or alkalinity of water on a log scale from 0 (extremely acidic) through 7 (neutral) to 14 (extremely alkaline). It is the negative base-10 log of the hydrogen ion (H+) activity in moles per litre. Hydrogen ions predominate in waters with pH less than 7; hydroxyl ions (OH-) predominate in waters with pH greater than 7. The pH of marine waters is close to 8.2, whereas most natural freshwaters have pH values in the range from 6.5 to 8.0. Most waters have some capacity to resist pH change through the effects of the carbonate-buffer system. In this system, hydroxyl ions produced during the hydrolysis of bicarbonate neutralise H+ ions, and maintain pH at a near constant level. Bicarbonate ions (HCO3-) are acquired from the weathering of silicate or carbonate minerals as rainwater passes through the soil zone.
Figure 1. The pH scale and H+ concentrations (modified after Miller1 for coastal situations). A difference of one unit corresponds to a tenfold change in pH.
What causes the pH of coastal waterways to change?
- The pH of coastal waters responds to changes in: (i) dissolved carbon dioxide concentrations; (ii) alkalinity; (iii) hydrogen ion concentrations; and (iv) in a small way to temperature. The concentration of carbon dioxide in solution is proportional to the partial pressure of the gas above the solution (i.e. Henry’s Law). Therefore an increase in CO2 in the atmosphere (due to the greenhouse effect) directly leads to an increase in the amount of CO2 absorbed by oceans. This is called ocean acidification, and is a very topical issue.
- The magnitude of pH change (due to the above factors) varies with salinity because various ions are involved in acid-base reactions, and because the concentration of salt influences various equilibrium constants. In natural waters, pH increases with salinity until calcium carbonate (CaCO3) saturation is reached (see Figure 2). When CaCO3 (calcite or aragonite) precipitates, the carbonate-alkalinity of water decreases. This causes a reduction in the buffering capacity of water and a decrease in pH. Calcium carbonate starts to precipitate when salinities exceed ~1 ppt in most Australian waters2.
- When seawater (pH 8.2; Figure 1 and 2) mixes with typical river water (pH 7-7.5; Figure 1 and 2), pH tends to decrease (Figure 2). However, river water has a much higher pH than seawater if it evaporates to the same salinity. Therefore, in evaporative coastal settings, mixtures of seawater and river water with relatively more river water may have higher pH levels than mixtures with relatively less river water, at the same salinity3 (Figure 2). It follows then, that freshwater deprivation due to the extraction and diversion of riverine water can alter natural pH ranges and gradients in coastal waterways.
Figure 2. Modelled change in pH with increasing salinity for evaporated coastal rainwater (meteoric), a basaltic groundwater, and mixtures between meteoric water and seawater (1 and 50%)3.
- Photosynthetic consumption of carbon dioxide (especially in algal blooms) can drive pH to high levels (see Figure 3). This is because there is less carbonic acid formation when carbon dioxide is consumed, and therefore less dissociation of carbonic acid into hydrogen ions.
- The decomposition of organic matter in the presence of dissolved oxygen (e.g. by oxygen reduction) increases the carbon dioxide content of water, and the lowers the pH. The first step in the nitrification reaction generates H+ ions and can also cause a local lowering of pH.
- Hydrogen ions are consumed in the processes that decompose organic matter in the absence of dissolved oxygen (e.g. nitrate reduction (a.k.a. denitrification), iron or manganese reduction and sulfate reduction). These processes cause pH to increase.
- The disturbance of tropical coastal (acid sulfate) soils, and the reclamation of coastal wetlands (including salt marshes and mangroves), can cause the oxidation of iron-sulfides (usually pyrite) stored in the soil. This can produce acid drainage, which has a typical pH range of 4 to ~2 (Figure 1)4. Mine drainage can have a similarly low pH.
- Acidic nitric and sulfur oxides derived from coal-fired power stations, some industrial operations, vehicle exhaust and emissions from thermal power stations give rise to atmospherically derived acids and potential acid deposition/rain. With the exception of industrial centres in the La Trobe Valley and Metropolitan Melbourne, acid deposition is relatively minor in Australia5.
- Humic acid waters are known in the Australian coastal zone6. The pH of these waters can be as low as 4.5, and is influenced by the organic matter content7.
- Chemical spills or the dumping of cleaners into stormwater drains can affect the pH of receiving waters.
Significance of pH
Most aquatic organisms and some bacterial processes require that pH be in a specified range. For example, the activity of nitrifying bacteria is optimal over a narrow pH range from 7 to 8.58. If pH changes above or below the preferred range of an organism (including microbes), physiological processes may be adversely affected7. This is especially true for most organisms if the ambient pH drops to below ~7 or rises to above 9. Physical damage to the gills, skin and eyes can also occur when pH is sub-optimal for fish, and skin damage increases susceptibility to fungal infections such as red spot disease. pH is driven more frequently to greater extremes under eutrophic conditions, allowing algal species with tolerance to extreme pH levels to grow and dominate communities, and to potentially form algal blooms.
Changes in pH can also have indirect impacts on aquatic organisms. For example, changes in pH can alter the biological availability of metals, the speciation of nutrients and the toxicities of ammonium, aluminium and cyanide7. Increases in pH can also cause the electrostatic forces that bind viruses to particles to be overcome, thus facilitating their release to the water column9. pH is important in calcium carbonate solubility (calcite or aragonite), which is important for shell-forming organisms. Shell growth (i.e. calcification) is inhibited if water becomes too acidic.
Considerations for measurement and interpretation
It is good practice to take pH measurements with all physical, chemical and biological samples. pH of water is best measured in situ using a meter equipped with a pH electrode. A high degree of precision can be expected from the method if careful attention is paid to the calibration and to the maintenance of electrodes and buffer solutions. Values are reported in standard pH units and usually to one or two decimal places. Repeated measures of pH should be reported as medians and ranges of measured values.
As with dissolved oxygen, pH measurements are most informative when the full diurnal range is known (see Figure 3). In a diurnal cycle, the lowest pH is expected at dawn because CO2 produced by decomposition and aerobic respiration would have accumulated since the previous dusk. Conversely highest pH is expected during the daylight hours, because pH rises at the rate at which carbon dioxide is fixed by plants. Diurnal changes in pH can be tracked using moored continuously recording pH sensors. At the very least, measurements should be taken at both dawn and midday. As a general rule, pH values in coastal waters that are higher than 9 and lower than 7 should be investigated.
Figure 3. Diurnal variations in surface pH in Lake Wollumboola (NSW) during December 2000 (from Haines et al., 200110 ). Peak pH concentrations coincided with peak dissolved oxygen concentrations in mid-afternoon at the height of algal photosynthesis10. Minimum algal productivity occurred at pre-dawn10. The large diurnal pH range indicates that the lake is highly productive.
Existing information and data
Large amounts of pH data exist for estuaries and coastal waters from around Australia. This data is held by the collecting agencies (State, local Government, community groups and environmental consultants). The ANZECC/ARMCANZ guidelines7 list default trigger values for pH for different bioregions, for comparison with pH medians, but recommend developing local objectives. The ANZECC/ARMCANZ guidelines7 also outline protocols to better define trigger values for pH.
More information on pH (changed from natural).
Lynda Radke, Geoscience Australia
- Miller, G.T. 2000. Living in the Environment, Brooks/Cole Publishing Company, Pacific Grove. Haines, Skyring, Stephens, Papworth (2001) Managing Lake Wollumboola’s Odour Problem. Proc. 11th NSW Coastal Conference, Newcastle 13-16 November 2001 ↩
- Radke, L.C., Howard, K.W.F., and Gell, P.A. 2002. Chemical diversity in southeastern Australian saline lakes I. Geochemical causes. Marine and Freshwater Research 53, 1-19. ↩
- Radke, L.C. 2002. Water allocation and critical flows: potential ionic impacts on estuarine organisms. Proceedings of Coast to Coast 2002 – “Source to Sea”, Tweed Heads, pp. 367-370. ↩ ↩
- Sammut, J., Melville, M.D., Callinan, R.B. and Fraser, G.C. 1995. Estuarine acidification: Impacts on aquatic biota of draining acid sulphate soils. Australian Geogrpahical Studies 33(1), 89-100. ↩
- Ayers, G.P. and Manton, M.J. 1991. Rainwater composition at two BAPMoN regional stations in SE Australia, Tellus 43B, 379-389 ↩
- Hawkins, P.R., Taplin, L.E., Duivenvoorden, L.J. and Scott, F. 1988. Limnology of oligotrophic dune lakes at Cape Flattery, North Queensland, Marine and Freshwater Research 39, 535-553. ↩
- ANZECC/ARMCANZ (October 2000) Australian and New Zealand Guidelines for Fresh and Marine Water Quality. ↩ ↩ ↩ ↩ ↩
- Henriksen, K, and Kemp, W.M. 1988. Nitrification in Estuarine and Coastal Marine Sediments, pp. 207-249 in T.H. Blackburn and J. Sorensen (eds.), Nitrification in Estuarine and Coastal Marine Sediments. Nitrogen Cycling in Coastal Marine Environments, John Wiley and Sons Ltd. ↩
- Miller, B.M. Issues for the modelling of fate and transport of viruses in estuarine environments, 15th Australasian Coastal and Ocean Engineering Conference, September 2001, Gold Coast. ↩
- Hinga, K.R. 2002. Effects of pH on coastal marine phytoplankton, Marine Ecology Progress Series 238, 281-300. ↩ ↩ ↩