## 3.1.3. Hydrosphere

Authors: John Parsons

Reviewers: Steven Droge, Sean Comber

Learning objectives:

You should be able to:

- describe the most important chemical components and their sources
- describe the most important chemical processes in fresh and marine water.
- be familiar with the processes regulating the pH of surface water.

Keywords: Hydrogen bonding, carbonates, dissolved salts

The properties and importance of water

Water covers 71% of the earth’s surface and this water, together with the smaller amounts present as gas in the atmosphere, as groundwater and as ice is referred to collectively as the hydrosphere. The bulk of this water is salt water in the oceans and seas with only a minor part of freshwater being present as lakes and rivers (Figure 1).

Figure 1. Global hydrological cycle and water balance (arrows are fluxes of water per year). Adapted from Kayane (1992) and Peixoto (1994) by Steven Droge 2019.

Water is essential for life and also plays a key role in many other chemical and physical processes, such as the weathering of minerals and soil formation and in regulating the Earth’s climate. These important roles of water derive from its structure as a small but very polar molecule arising from the polarised hydrogen-oxygen bonds (Figure 2). As a consequence, water molecules are strongly attracted by hydrogen bonding, giving it relatively high melting and boiling points, heat capacity, surface tension, etc. The polarity of the water molecule also makes water an excellent solvent for a wide variety of ionic and polar chemicals but a poor solvent for large nonpolar molecules.

Figure 2. Hydrogen bonding between water molecules

The freshwater environment

As mentioned above, freshwater is only very small proportion of total amount of water on the planet and most of this is present as ice. Since this water is in contact with the atmosphere and the soils and bedrock of the Earth’s crust, it dissolves both atmospheric gases such as oxygen and carbon dioxide and salts and organic chemicals from the crust. If we compare the relative compositions of cations in the Earth’s crust and the major dissolved species (Table 1) it is clear that these are very different. This difference reflects the importance of the solubility of these components. For ionic chemicals, this depends on both their charge and their size (expressed as z/r2, where z is the charge and r the radius of an ion). As well as reflecting the properties of the local crust, the composition of salts is also influenced by precipitation and evaporation and the deposition of sea salt in coastal regions.

Table 1. Comparison of the major cation composition of average upper continental crust and average river water. (*except aluminum and iron from Broecker and Peng (1982))

| Upper continental crust (mg/kg) (Wedepohl 1995*) | River water (mg/kg) (Berner & Berner 1987*)
Al | 77.4 | 0.05
Fe | 30.9 | 0.04
Ca | 29.4 | 13.4
Na | 25.7 | 5.2
K | 28.6 | 1.3
Mg | 13.5 | 3.4

The pH of surface water is determined by both the dissolution of carbonate minerals and carbon dioxide from the atmosphere. These components are part of the set of equilibrium reactions known as the carbonate system (Figure 3).

Figure 3. The carbonate system of equilibria regulating the pH of surface water. Source: http://butane.chem.uiuc.edu/pshapley/GenChem1/L26/3.html

At equilibrium with the current atmospheric CO2 concentration and solid calcium carbonate, the pH of surface water is between 7 and 9 but this may reach more acidic values where soils are calcium carbonate (limestone) poor. This is illustrated by the pH values measured in a river in Northern England, where acidic, organic carbon-rich water at the source is gradually neutralised once the river encounters limestone rich bedrock (Figure 4).

Figure 4. Water chemistry in the Malham Tarn area of northern England, showing the relationship between pH, alkalinity and dissolved calcium as water flows from bog on siliceous mudrock to limestone where the pH is buffered around 8 due to once limestone weathering. Redrawn from Fig. 5.6 in Andrews et al. (2004) by Wilma Ijzerman.

As well as these natural processes, there are human influences on the pH of surface water including acidic precipitation resulting from fossil fuel combustion and acidic effluents from mining activities caused by oxidation and dissolution of mineral sulphides. Regions such as Southern Scandinavia with carbonate-poor soils are particularly vulnerable to acidification due to these influences and this is reflected in for example, reduced fish populations in these vulnerable regions (see Figure 5). More recently, reduced coal burning and the decline in heavy industry is resulting in the recovery of pH values in upland areas across Europe.

Figure 5. Average catch size of salmon in seven rivers in southern Norway receiving acid precipitation and 68 other rivers that do not receive acid precipitation. Redrawn with data from Henriksen et al. (1989) by Wilma Ijzerman.

Dissolved oxygen is of course essential to aquatic life and concentrations are in general adequate in well mixed water bodies. Oxygen can become limiting in deep lakes where thermal stratification restricts the transport of oxygen to deeper layers, or in water bodies with high rates of organic matter decomposition. This may result in anoxic conditions with significant ecological impacts and on the behaviour of chemical contaminants.

The marine environment

Freshwater eventually moves into seas and oceans where the concentrations of dissolved species are much higher than in the freshwater environment. This is partly due to the effects of evaporation of water from the oceans but is also be due to specific marine sources of some dissolved components. Estuaries are the transition zones where freshwater and seawater mix. These are highly productive environments where increasing salinity has a major impact on the behaviour of many chemicals, for example on the speciation of metals and the aggregation of colloids as a result of cations shielding the negative surface change of colloidal particles (Figure 6). Increasing salinity also affects organic chemicals, with ionic chemicals forming ion pairs, and even reducing the solubility of neutral organics (the so-called salting-out effect). As well as these chemical effects due to increasing salinity, the lowering of flow rates in estuaries leads to the deposition of suspended particles.

Figure 6. The Electrical Double Layer (EDL), comprising a fixed layer of negative charge on a clay particle (due to isomorphic substitutions and surface acids) and a mobile ionic layer in solution. The latter is caused because positive ions are attracted to the particle surface. Note that with increasing distance from the particle surface the solution approaches electrical neutrality. (Source Steven Droge 2019)

Since the concentrations of pollutants are in general lower in the marine environment than in the freshwater environment, concentrations in estuaries decrease as freshwater is diluted with seawater. Measuring salinity at different locations in estuaries is a convenient way to determine the extent of this dilution. Components that are present in higher concentrations in seawater will of course show an increase with salinity unless. Plotting salinity against the concentrations of chemicals at different locations can yield information on whether they behave conservatively (i.e. only undergoing mixing) or are removed by processes such as degradation or partitioning into the atmosphere or sediments. Figure 7 shows examples of plots expected for conservative chemicals and those that are either removed in the estuary or have local sources there. Models describing the behaviour of chemicals in estuaries can be used with these data to derive the rates of removal or addition of the chemical in the system.

Figure 7. Idealized plots of estuarine mixing illustrating conservative and non-conservative mixing. CR and CS are the concentrations of the ions in river and seawater respectively. Redrawn from Figure 6.3 in Andrews et al. (2004) by Wilma Ijzerman.

The open ocean is sufficiently mixed for the composition of major dissolved constituents to be fairly constant, except in local situations as a result of upwelling of deep nutrient-rich waters or the biological uptake of nutrients. In coastal regions the concentrations of chemicals and other components originating from terrestrial sources may also be locally higher. The major components in seawater are listed in Table 2 with their typical concentrations.

Table 2. Major ion composition of freshwater and seawater.

| Seawater (mmol/L) (Broecker and Peng, 1982) | River water (mmol/L) (Berner and Berner, 1987)
Na+ | 470 | 0.23
Mg2+ | 53 | 0.14
K+ | 10 | 0.03
Ca2+ | 10 | 0.33
HCO3- | 2 | 0.85
SO42- | 28 | 0.09
Cl- | 550 | 0.16
Si | 0.1 | 0.16

These concentrations may be higher in waterbodies that are partly or wholly isolated from the oceans and are impacted by evaporative losses of water (e.g. Mediterranean, Baltic, Black Sea). In extreme case, concentrations of salts may exceed their solubility product, resulting in precipitation of salts in evaporate deposits.

As is the case in freshwater, carbonates play an important role in regulating the ocean pH. The fact that the oceans are supersaturated in calcium carbonate makes it possible for a variety of organisms to have calcium carbonate shells and other structures. The important processes and equilibria involved are illustrated in Figure 8. There is concern that one of the most important effects of increasing atmospheric carbon dioxide will be lowering of ocean pH to values that will result in destabilisation of these carbonate structures.

Figure 8. (a) Schematic diagram to illustrate the buffering effect of CaCO3 particles (suspended in the water column) and bottom sediments on surface water HCO3- concentrations (after Baird and Cann 2012). (b) A sample of the seawater in (a) will have a pH very close to 8 because of the relative proportions of CO2, HCO3- and CO32-, which in seawater is dominated by the HCO3- species. Increased CO2 concentrations in the atmosphere from anthropogenic sources could induce greater dissolution of CaCO3 sediments including coral reefs. Redrawn from Figure 6.8 in Andrews et al. (2004) by Wilma Ijzerman.

References

Andrews, J.E., Brimblecombe, P., Jicketts, T.D., Liss, P.S., Reid, B.J. (2004). An Introduction To Environmental Chemistry, Blackwell Publishers, ISBN 0-632-05905-2.

Baird, C., Cann, M. (2012). Environmental Chemistry, Fifth Edition, W.H. Freeman and Company, ISBN 978-1429277044.

Berner, E.K., Berner, R.A. (1987). Global water cycle: geochemistry and environment, Prentice-Hall.

Broecker, W.S., Peng, T.S. (1982). Tracers in the Sea, Lamont-Doherty Geol. Obs. Publ.

Henriksen, A., Lien, L., Rosseland, B.O., Traaen, T.S., Sevaldrud, I.S. (1989). Lake Acidification in Norway: Present and Predicted Fish Status. Ambio 18, 314-321

Wedepohl, K.H. (1995). The composition of the continental crust, Geochimica Cosmochimica Acta 59, 1217-1232.