Aquatic Ecosystem Management

Aquatic chemistry primer

This primer is intended to supply only the most basic information about aquatic chemistry.
For detail consult a textbook, eg.:
Stumm, W and Morgan, JJ. 1996. Aquatic Chemistry: Chemical Equilibria and Rates in Natural Waters, Third Edition. John Wiley & Sons, Inc, New York;
Hunt, DTE & Wilson, AL. 1986. The Chemical Analysis of Water: general principles and techniques. Royal Society of Chemistry, Cambridge.,
or look up the following websites:
Sextant (Dave McShaffrey, Marietta College),
Chapter3: Aquatic Chemistry (Todd Rasmussen, University of Georgia)

Sections:


Water is of course the very basis of aquatic chemistry and the structure of water dictates its behaviour.

The chemical formula of water is commonly written H2O, but it is more instructive to see it as HOH, since this how the atoms are actually arranged. In fact, the bonds linking the two hydrogens to the oxygen are not in a straight line but at an angle, and this has the consequence of making the molecule polar, meaning that there is an electrical charge difference beween one end of the molecule and the other (if the atoms were in a straight line they would cancel each other out: this is the case for CO2). Since there are 2 centres of charge it is also referred to as a dipole.


A measure of a substance’s polarity can be made by assessing its dielectric constant, which is its ability to increase the electrical storage of a capacitor: the molecules line up in an electric field (when placed between 2 charged plates) and the energy required to do this can be released again when the potential is removed from the plates.

Non-polar molecules include molecules composed of 2 atoms of the same substance (O2), ones where the structure is symmetrical so that the charges cancel out, and ionically balanced compunds such as many hydrocarbons, which are insoluble in water.

Other important properties of water are its thermal behaviour: on freezing it expands so that ice floats, and its specific heat is greater than any other substance, meaning that bodies of water show great thermal stability than terrestrial environments (specific heat is the amount of heat energy required to raise the temperature of a substance by 1°C), and its ability to dissolve other substances.

Water itself is more than just molecules of H2O: dissociation into the ions H+ and OH- occurs and in pure water there 10-7 grams per litre of H+ (and the equivalent amount of OH-). This is normally expressed as a pH value, which is the negative exponent of the hydrogen ion concentration, ie. pH 7.0. Increasing or decreasing the pH value by one unit therefore means a tenfold change in hydrogen ion concentration. Acids and bases change the pH by altering the dissociation of water, acids increase H+ (decrease pH). pH influences the solubility of many other substances in water. (see also here). The ability of a water to resist the addition of hydrogen ions (ie. acid) is measured as its alkalinity, expressed as mgl-1 of CaCO3 (even though there may be other compounds present contributing to the alkalinity, eg. carbonates, bicarbonates and hydroxides of different metals). It is determined by titration with acid to a pH of (usually) 8.3.

Some substances dissolved in water also dissociate into ions. Positively charged ions (H+, Ca++, Na+, etc.) are called cations, negatively charged are anions (OH- Cl- SO4=). The number of ions present influences the ability of the solution to conduct electricity: the conductivity is thus an index of the amount of dissolved ionisable salts. Conductivity is measured in micro-siemens (µS). Substances which do not ionise (eg. sugar) do not affect conductivity. On the other hand, both change the refractive index of water (the extent to which light is bent), and a refractometer can be used to measure the amount of dissolved material in either salt solutions eg. seawater, or sugar solutions, as in brewing. Salts are defined chemically as being the product of a reaction between a metal or base and an acid, eg, potassium chloride, ferrous sulphate, sodium acetate, ammonium chloride (most bases are hydroxides of metals).

Solid substances may be readily soluble in water or only slightly, but the solubility can be increased by increasing the temperature. However, in the case of gases, the opposite is the case. As temperature rises, the amount of gas dissolved decreases, hence the formation of bubbles as water is brought to the boil. In nature, this has profound effects on organisms which require to obtain their oxygen from the water. An increase in temperature may reduce the concentration of oxygen below that required to sustain life; thus salmonid fish are restricted to relatively cold (and unpolluted) waters. Carbon dioxide is readily soluble in water, and also reacts with carbonates and bicarbonates. This forms the major buffering system in natural waters, around a neutral pH of 7.0. As acid is added to a buffered water, some of the salts present change to a different state, thus preventing a change in pH. In the case of carbonate-bicarbonate, adding acid changes the balance in favour of CO2, which is released to the atmosphere, and prevents a pH shift. Changing temperature may also drive off CO2 and thus cause a rise in pH.

CO2 + H2O <> H2CO3 carbonic acid

H2CO3 <> HCO3- (bicarbonate) + H+

Most plants take up their carbon requirement as bicarbonate, though some use dissolved CO2, so plant growth also removes CO2 from the buffering system, resulting in a rise in pH.

Carbonates in water are often associated with the metal ions Ca++ and Mg++. Such waters are described as “hard” (due to their feel when used for washing using old-fashioned soap as opposed to “soft” waters, since the salts precipitate out the soap and prevent the water feeling soapy). Hard water also produces “scale” when boiled, as the carbonates precipitate out:

Ca(HCO3)2 (soluble) <>CaCO3 (insoluble) + H2O + CO2

There are 2 types of hardness, depending on the forms of calcium and magnesium present, referred to as “temporary “ and “permanent”; temporary hardness is removed by boiling, when the CO2 is driven out of calcium bicarbonate to form calcium carbonate; permanent hardness is due to calcium sulphate, which is only sparingly soluble, but does not form a less soluble salt on boiling as in the reaction above.

Natural waters contain a whole range of different dissolved materials. Sea water tends to have a fairly constant composition, though in some areas it may be diluted by the influx of fresh water from major rivers, etc. Natural fresh waters however are much more variable as they are influenced very strongly by the geology of the land over (or under) which they have run since falling as rain. Some inland saline waters are very much more concentrated than the sea and some soda lakes have corrosively alkaline waters. However, all these waters share the fact that they contain equivalent (not equal) concentrations of anions and cations.

Some definitions:
atomic mass is the mass in AMU (atomic mass units) of an atom, relative to hydrogen (=1).
Molecular weight is the weight (mass) of a molecule, being the sum of the atomic masses of its constituents. eg. Hydrogen = 1AMU; oxygen = 16 AMU, H2O = 18 molecular weight.
1 mole = the molecular weight expressed in grams (= the weight containing Avogadro’s number of molecules).
1 Molar = a solution containing 1 mole per litre of the solute.
Valency = the number of bonds an atom (or ion) can form with others when forming a compound, eg. in water hydrogen has a valency of 1 while oxygen has a valency of 2, hence the molecular structure of water. This is the same as the charge on the ion: compare O= and H+ . Some ions are trivalent, eg. PO43-

In a solution, the ionic charges must balance, so all the anions must be added together in a way which takes account of their charges; this is done by considering their equivalence. Thus a molar solution of NaCl contains 1 mole (= 1equivalent) of sodium and 1 mole of chloride;
1 molar sodium carbonate (Na2CO3) contains 2 moles (2 equivalents) of sodium and 1 mole (2 equivalents) of carbonate;
1 molar orthophosphoric acid (H3PO4) contains 3 moles of hydrogen (3 equivalents) and 1 mole of phosphate (3 equivalents).

The solubility of substances in water is affected by temperature, as already mentioned, and in the case of gases, by pressure (more gas can be dissolved under pressure: think of carbonated drinks). Solubility is also influenced by the presence of other dissolved substances, and this has an especially important biological consequence in the case of oxygen. Solubility of oxygen in fresh water is about 10mg/l (at 15°C), but in sea water is about 1 mg/l lower. This means it is easier for sea water to to become deoxygenated.

Solubility of solids, particularly of metals, may be increased by the presence of certain other substances called chelating agents. These are molecules which can form double-bonded complexes with metals, particularly those with a valency greater than1 (eg. Mg2+, Fe3+). There are synthetic chelating agents such as EDTA, and natural ones such as citrate and humus. The complexes may be much more soluble than the original metal salt, allowing toxic levels to accumulate in the water.
On the other hand, some complexes may precipitate out, leading to a loss from solution of an element, eg. the precipitation of aluminium by complexation with humus.

The ability of a solution to donate or receive electrons (ie to oxidise or reduce) is measured by its Redox potential. Its redox state influences the speciation of metals (which can affect their solubility and thus the toxicity of the water) and determine eg. whether sulphur is present as sulphate ion , or as H2S.
(for Redox see also here).



Analytical Techniques in water quality assessment


(see: Hunt, DTE & Wilson, AL. 1986. The Chemical Analysis of Water: general principles and techniques. The Royal Society of Chemistry, Cambridge;
APHA. 1996. Standard Methods for the Examination of Water and Wastewater, 20th edition, eds Eaton, AD, Clesceri, LS and Greenberg, AE. American Public Health Association. )

Substances and conditions which are important constituents or pollutants in water fall into the following categories:

1: naturally occurring substances/conditions which may be essential (in the right concentrations)- Oxygen, carbon dioxide, carbonate/ bicarbonate, pH, alkalinity, BOD, COD, organic carbon, (ortho)phosphate, nitrate, nitrite, ammonium, silicate, metals (Al, Ca, Mg, Mn, Fe, Bo, Na, K, Cu – as ions), anions (Cl-, SO4=, etc.)

2: substances which ought not to be present (though may occur naturally in some places):
heavy metals (Pb, As, Hg, Cd, Zn, Sn) , hydrocarbons both straight-chain (aliphatic) and ring-containing (aromatic), both of which may be halogenated (which makes them more persistent), other organic compounds of human origin such as organophosphorus pesticides, polycyclic aromatic hydrocarbons and radioactive isotopes.

A range of analytical techniques are employed to test waters for the presence and abundance of these.

Methods include:

Titration (volumetric analysis), in which a reagent of known concentration (expressed in Normality: 1N = 1M/valency) is added gradually to a known volume of the determinant in solution, until an end-point is reached when the determinant is fully reacted. The volume of reagent added can then be used to calculate the concentration of determinant. End-point is measured with an indicator (eg. pH indicator) or using some form of instrumentation (eg. pH meter). A simple example is the titration of a base with an acid; the end point is neutrality.

Electrochemical measurement using probes or sensors, eg. pH, dissolved oxygen, conductivity. Some of these sensors may also be employed along with other techniques such as titration or ion chromatography.

Voltammetry. This and the related technique of polarography measure ions by their electrochemical properties. They use a dropping mercury electrode, or mercury film, in a cell containing the sample. A potential is applied to the electrode and gradually changed, while the current passing is measured. The current changes discontinuously with voltage applied, with plateaus corresponding to the depolarisation of the electrode by a particular ion, and the current change following each plateau relates to the concentration of that ion. It has limited use in water analysis.

Spectrophotometry (/colorimetry): The absorbance of light (visible or UV) by material in solution is proportional (between limits) to its concentration. Colorimetry is relatively crude, using filters to produce light which spans quite a broad range of wavelengths, whereas spectophotometry typically uses a waveband of 4nm or less, making it much more selective. Some determinands are assayed directly in the sample (eg. high concentrations of nitrate) whereas the majority must be mixed with reagents which produce a coloured product (in proportion to the amount of determinand present). Spectrophotometric detectors are also used in High Performance Liquid Chromatography (HPLC)

Flame Emission Spectrometry. Note that spectrometry is different from spectrophotometry, meaning the measurement of spectral lines. When an element is heated to incandescence, it emits light at unique wavelengths. If a solution is introduced as fine droplets into a flame, and the light emitted is passed through filters or a monochromator, the intensity of emission at the wavelength(s) characteristic of the element being investigated can be measured, and related to the concentration of the element in solution. It is most widely used for the measurement of Na and K, plus some other, but not all, metals.

Atomic Absorption Spectrometry (AAS). This differs from FES in that the sample is nebulised in a flame or heated furnace but not to the point of emitting light. Instead, usually, light with spectral lines corresponding to a particular element (emitted from a special lamp with a hollow cathode constructed of that element) is passed through it and the absorption of the light measured by a photomultiplier. Variations exist for certain elements. It is most widely used in the determination of metals.

Chromatography. This technique got its name from early methods used to separate coloured compounds, but it really has nothing to do with colour. It relies on separating molecules from each other on the basis of their ability to move through a matrix (the stationary phase) when a mobile phase is allowed to pass through it.

The stationary phase may be solid or liquid and the mobile phase liquid or gas. The simplest chromatographic technique is paper or thin-layer chromatography in which molecules may be separated in 2 dimensions by using 2 mobile phases in turn, ar right angles to each other. However, the techniques most commonly used in water analysis are High Performance Liquid Chromatogtaphy (HPLC) in which the moblile phase is liquid, and Gas Chromatography (GC), in which the mobile phase is gas. Both have stationary phases mounted in columns, which may be relatively wide (~5mm or more) packed with solid material, or capillary (1mm-0.25mm), with the stationary phase distributed around the walls. These latter techniques lend themselves to automated operation, and are principally used for assay of organic compounds.

GC is used for the analysis of gases or liquids which can be vaporised at relatively low temperatures. With liquid samples, the sample itself is diluted in a solvent with (usually) a much lower boiling point, beteen room temperature and about 100°C – the lower the better. It is injected into a heated (eg 250°C) tube at the end of the column (at slightly above ambient) which has an inert carrier gas passing through it (eg. H2, He or N2). Molecules in the sample enter the column and attach themselves to the stationary phase. The solvent and other low-boiling point components start to migrate along the column (which in a capillary column may be 10-60metres long, made of fused silica tubing) and the solvent would normally emerge first after 1-3 minutes. After a short time the column temperature is gradually increased at a controlled rate up to (a maximum of) about 300°C, which forces higher BP components into the mobile phase, and separates them from each other, largely on the basis of molecular weight. As they emerge from the end of the column, they pass through a detector which determines the amount of material in each peak. There are several different types of detector; Flame Ionisation Detector (FID), Flame Photometric Detector (FPD), Thermionic Detector (TID), Electron Capture Detector (ECD), and Mass Selective Detector (MSD). All but the last determine only the amount of material in each peak, so that identification of the material depends on comparing its retention time with a known standard. Each has its particular uses, eg. the ECD for halogenated compounds, the FID as an inexpensive, fairly sensitive general purpose detector. Typical uses for GC in water analysis are hydrocarbons, organophosphorus pesticides, organochlorines.

The MSD carries the analysis a stage further by ionising the materials emerging from the column and then separating them on the basis of molecular (ionic) mass, producing a “mass spectrum”. Each molecule breaks down in a characteristic way into component ions, and the spectrum can be used to identify them. When used like this the MSD is not as sensitive as the FID for many analytes, but it can also be used to monitor for specific ions, if the nature of the analyte is known or suspected, and in this mode (Selective Ion Monitoring) is one of the most sensitive detector techniques (it is used eg. for the detection of narcotics).

HPLC separates material in a way analogous to that described for GC, but isothermically, using a liquid mobile phase at high pressure. It is used for organic material which is not volatile, such as pesticides, herbicides and phenols. Again, there is a range of detectors available: UV absorption, fluorometric, refractometric, and recently MSD (HPLC-MS).

Ion Chromatography (IC) is related to HPLC, but specifically designed for (low molecular weight) inorganic ions. They are separated on ion-exchange resin columns, and different columns are required for anions and cations. Detection is by conductivity, after passing through a second column which suppresses the conductivity due to the eluent (“mobile phase”). IC is used for eg. NO3-, NO2-, PO4=, SO4=, I-, etc. Its advantage is the simultaneous quantification of all such ions in one run, compared with running separate assays by spectrophotometric methods, and has similar sensitivity in many cases. The disadvantage is that it cannot be used with saline waters as the high Cl- levels swamp the detector, and organics must be removed by pre-treatment.

 

March 2004


MSc Aquatic Ecosystem Management

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