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Dissolved Oxygen (DO)

Handheld Dissolved Oxygen IP67 Dissolved Oxygen Meters

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Industrial Dissolved Oxygen 

Dissolved Oxygen Probes

DO Probe Conditioner


Portable pH Testers

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Replacement pH Electrodes


Free Chlorine

Total Chlorine

Cyanuric Acid

Redox (ORP)

Portable Redox Testers

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Conductivity (EC)

Portable Conductivity Testers

Handheld Conductivity Meters

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Replacement Conductivity Cells

Total Dissolved Solids (TDS)

Portable TDS Testers

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Ion Concentration

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Portable Salinity Testers

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Salinity Sensors

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Chlorine Dioxide

Paracetic Acid

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pH, Reference, Conductivity, Redox, Dissolved Oxygen, Ion Selective Electrodes, Calibration Solutions, Hollow Cathode Lamps and

Deuterium Lamps

Technical Area


Please click on one of the following links for a detailed introduction to your chosen application  :

Dissolved Oxygen  



Conductivity (EC)

Total Dissolved Solids (TDS)  

ORP (REDOX)     






Introduction to Dissolved Oxygen


Many gases mix with water like nitrogen and oxygen without chemically reacting with it. Some gases chemically react with water e.g. ammonia, CO2, and HCl. Oxygen does not react with water. Dissolved Oxygen (DO) is really a physical distribution of oxygen molecules in water. There are two main sources of DO in water: atmosphere and photosynthesis. Waves and tumbling water mix air into water where oxygen readily dissolves until saturation occurs. Oxygen is also produced by aquatic plants and algae as a by-product of photosynthesis.

Basic Principle in DO measurement (OR What is Being Measured?)

The electrochemical method of measuring DO requires a cathode, anode, electrolyte solution and a gas permeable membrane. The material of the membrane is specially selected to permit oxygen to pass through. Oxygen is consumed by the cathode which will create a partial pressure across the membrane. Oxygen will then diffuses into the electrolyte solution.

Thus, a DO meter actually measures the pressure of oxygen in water. It can be used to measure DO in any medium.

Polarographic or Clark Cell Method

Dr. Clark first discovered the cell to measure oxygen in 1956. This is basically an amperometric cell that is polarized around 800 mV. This cell, named after Dr. Clark, is built around the popular Ag/AgCl half-cell and a noble metal such as gold, platinum or palladium. Reduction of oxygen is achieved between 400 to 1200 mV, hence a need for a voltage of around 800 mV. This is provided externally by a battery source.

Electrolyte used: KCl or KBr

Cell Reaction
Anode 2Ag + 2Cl- 2AgCl + 2e-
Cathode (Platinum, gold or palladium) 2e- + O2 + H2O 2 OH-
Total Reaction 2e- + O2 + H2O + 2Ag + 2Cl- 2 OH- + 2AgCl + 2e-

From the above reaction, every time oxygen is reduced at the cathode, 4 electrons or current is generated directly proportional to the oxygen consumed (reduced) at the cathode.

There are four major problems associated with this type of DO measurement: .

Problem Description
Isolation of Anode Since the net result of the chemical reaction is AgCl, over time, a build up of AgCl will coat the anode. Once the whole anode is covered, reaction stops and the oxygen probe stops working. The probe can be reactivated by cleaning the anode to remove the AgCl deposit which can be a time-consuming procedure.
Zero shift The result of the above reaction produced more OH- ions which will move the pH value of the electrolyte. The electrolyte, which is around neutral pH value, will moves into the alkaline range. This causes a zero shift, and over time, the electrolyte will need to be changed.
Depletion of Chloride The net reaction also consumes Cl- ions. Over time, the chloride ions will be consumed and the electrolyte needs to be replaced.
Warm-up Time The major disadvantage is the need for an external power source of approximately 800 mV to be applied to the electrode. As soon as the probe is disconnected, power supply is cut off. On connecting the probe again, the user must wait for the probe to be polarized, that is, for the current loop to be stabilized. This warm-up time is approximately 10 minutes. Any measurement taken before this warm-up time period will be normally a higher value and will result in wrong readings.

Galvanic Cell Method

The galvanic probe principle was introduced by Macreth in 1964 and has gone through a few changes.

The main advantage of a galvanic probe is that is does not need an external power supply to provide polarization as required by the Clark Cell. This is achieved by using two dissimilar metals. In the presence of a electrolyte, there is an electromotive voltage produced between the two metals. At approximately 800 mV, this is large enough to reduce the oxygen at the cathode. If lead and gold or lead and silver is used, the differential voltage is approximately 800 mV.

Hence, a galvanic probe is really a self-polarizing amperometric cell. The single biggest advantage is the fact that the cell is now always ready and there is no warm up time.

Electrolyte used: KCl or KBr

Cell Reaction
Anode (Zinc or Lead) Zn Zn2+ + 2e-
Cathode 2e- + O2 + H2O 2OH-
Total Reaction Zn + 2e- + O2 + H2O Zn2+ + 2e- + 2OH-

Zn + O2 + H2O Zn (OH)2

ZnO (white precipitate) + H2O or:

Zn + O2 ZnO

Hence one molecule of oxygen produces 4 electrons and there is a direct relationship between the oxygen consumed at the cathode and the current produced by the cell.

The net result of the chemical reaction is simply ZnO which is reasonably stable and does not coat the anode. Water is recreated and the electrolyte is not consumed. Theoretically, the electrolyte will go on forever without replenishment.

Hence, because of the advantages galvanic probe has, Eutech Instruments has chosen this superior technology to make it easier for users.

DO Measurement Considerations

The amount of DO that can be held by water depends on 3 factors: water temperature, salinity, and atmospheric pressure; Amount of DO increases with decreasing temperature (colder water holds more oxygen).;Amount of DO increases with decreasing salinity (freshwater holds more oxygen than saltwater does);Amount of DO decreases with decreasing atmospheric pressure (amount of DO absorbed in water decreases as altitude increases).

The chart below shows the solubility of DO in mg/l in water at various temperature.

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Introduction to Turbidity Measurement


Turbidity is a cloudiness or haziness in water (or other liquid) caused by individual particles that are too small to be seen without magnification, thus being much like smoke in air. Liquids can contain suspended solid matter consisting of particles of many different sizes. While some suspended material will be large enough and heavy enough to settle rapidly to the bottom of a container if a liquid sample is left to stand (the settleable solids), very small particles will settle only very slowly or not at all if the sample is regularly agitated or the particles are colloidal. These small solid particles cause the liquid to appear turbid.

There are several practical ways of quantifying cloudiness in water, the most direct being some measure of attenuation (that is, reduction in strength) of light as it passes through a sample column of water. The now little-used Jackson Candle method (units: Jackson Turbidity Unit or JTU) is essentially the inverse measure of the length of a column of water needed to completely obscure a candle flame viewed through it. The more water needed (the longer the water column), the clearer the water. Of course water alone produces some attenuation, and any substances dissolved in the water that produce colour can attenuate some wavelengths. Modern instruments do not use candles, but this approach of attenuation of a light beam through a column of water should be calibrated and reported in JTUs.

A property of the particles — that they will scatter a light beam focused on them — is considered a more meaningful measure of turbidity in water. Turbidity measured this way uses an instrument called a nephelometer with the detector setup to the side of the light beam. More light reaches the detector if there are lots of small particles reflecting the source beam than if there are few. The units of turbidity from a calibrated nephelometer are called Nephelometric Turbidity Units (NTU). To some extent, how much light reflects for a given amount of particulates is dependent upon properties of the particles like their shape, colour, and reflectivity. For this reason (and the reason that heavier particles settle quickly and do not contribute to a turbidity reading), a correlation between turbidity and TSS is somewhat unique for each location or situation.

Turbidity in lakes, reservoirs, and the ocean can be measured using a Secchi disk. This black and white disk is lowered into the water until it can no longer be seen; the depth (secchi depth)is then recorded as a measure of the transparency of the water (inversely related to turbidity). The Secchi disk has the advantages of integrating turbidity over depth (where variable turbidity layers are present), being quick and easy to use, and inexpensive. It can provide a rough indication of the depth of the euphotic zone with a 3-fold multiplication of the secchi depth. However, this cannot be used in shallow waters where the disk can still be seen on the bottom.

There are frequently standards on the allowable turbidity in drinking water. In the United States (as of 2003) the allowable standard is 1 NTU, with many drinking water utilities striving to achieve levels as low as 0.1 NTU

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Introduction to pH and pH Measurement



The pH of a solution measures the degree of acidity or alkalinity relative to the ionization of water sample. Pure water dissociates to yield 10-7 M of [H+] and [OH-] at 25 oC; thus, the pH of water is neutral i.e. 7. 

pHwater = - log [H+] = - log 10-7 = 7

Most pH readings range from 0 to 14. Solutions with a higher [H+] than water (pH less than 7) are acidic; solutions with a lower [H+] than water (pH greater than 7) are basic or alkaline.

pH Measurement

Measuring pH involves comparing the potential of solutions with unknown [H+] to a known reference potential. pH meters convert the voltage ratio between a reference half-cell and a sensing half-cell to pH values.

In acidic or alkaline solutions, the voltage on the outer membrane surface changes proportionally to changes in [H+]. The pH meter detects the change in potential and determines [H+] of the unknown by the Nernst equation:

E = Eo + (2.3RT)/nF log {unknown [H+]/internal [H+]}  

where: E = total potential difference (measured in mV); Eo = reference potential; R = gas constant; T = temperature in Kelvin; n = number of electrons; F = Faraday's constant; [H+] = hydrogen ion concentration.

pH Temperature Compensation

The pH of any solution is a function of its temperature. Voltage output from the electrode changes linearly in relationship to changes in pH, and the temperature of the solution determines the slope of the graph. One pH unit corresponds to 59.16 mV at 25 C, the standard voltage and temperature to which all calibrations are referenced. The electrode voltage decreases to 54.20 mV/pH unit at 0.0 C and increases to 74.04 mV/pH unit at 100.0 C.

Since pH values are temperature dependent, pH applications require some form of temperature compensation to ensure standardized pH values. Meters and controllers with automatic temperature compensation (ATC) receive a continuous signal from a temperature element and automatically correct the pH value based on the temperature of the solution. Manual temperature compensation requires the user to enter the temperature of the solution in order to correct pH readings for temperature. ATC is considered to be more practical for most pH applications.

pH System

A successful pH reading is dependent upon all components of the system being operational. Problems with any one of the three: electrode, meter or buffer will yield poor readings.

pH Electrodes : 

A pH electrode consists of two half-cells; an indicating electrode and a reference electrode. Most applications today use a combination electrode with both half cells in one body. Over 90% of pH measurement problems are related to the improper use, storage or selection of electrodes.

pH Meters : 

A pH meter is a sophisticated volt meter capable of reading small millivolt changes from the pH electrode system. The meter is seldom the source of problems for pH measurements. Today pH meters have temperature compensation (either automatic or manual) to correct for variations in slope caused by changes in temperature. Microprocessor technology has created many new convenience features for pH measurement; auto-buffer recognition, calculated slope and % efficiency, log tables for concentration of ions and more. 

pH Buffers : 

These solutions of known pH value allow the user to adjust the system to read accurate measurements. For best accuracy:

  • Standardization should be performed with fresh buffer solutions.
  • Buffer used should frame the range of pH for the samples being tested.
  • Buffers should be at the same temperature as the samples. (For example: if all your samples are at 50 C, warm your buffers to 50 C using a beaker in a warm bath.)

Buffer values are dependent upon temperature.

pH Electrode Training Guide

Section 1: Electrode Construction

The pH Sensitive Membrane

The most common type of sensitive membrane used on a pH electrode is a blown glass bulb or rod. The glass used on Russell Mainstream electrodes is suitable for most applications. If, however, the application involves the constant monitoring of high temperature liquids or high pH values (above pH 13), then an alternative glass type can be specified. A bulb configuration will provide a fast response and accurate results when used in a sample of low ionic strength whereas a rod or bullet shaped membrane is very rugged and will be more resistant to breakage.

The Reference Cell

Housed within the outer chamber of the pH electrode is a reference system which is designed to provide a stable reference voltage for the sensor. This reference 'half-cell' will maintain a constant output in all liquids. Reference cells consist of an internal element (usually a Ag/AgCl wire), an electrolyte (usually KCl solution) and a liquid junction. The liquid junction provides a leak path for the internal electrolyte to 'weep' into the sample chamber and provide an electrical contact with the liquid to be measured. If the liquid junction is not efficient then measurements will be inaccurate.

The Cap/Cable/Connector

Electrodes used in laboratories are usually fitted with 16mm diameter caps to suit cantilever arm electrode stands. The cable used is a high grade, screened coaxial type with low noise characteristics. Because of the high impedance of pH electrodes, typically 100 megohms, connectors should always be kept clean and dry. Detachable cable electrodes should not be used in very humid environments.

Section 2: How to Specify an Electrode

The following check list, when used with the pH electrode selection chart, will help to identify a suitable electrode for any given application.

  • Sample Type
  • Temperature
  • Pressure
  • Expected pH Range
  • Viscosity of Sample
  • Sample Volume
  • Make and Model of pH Meter (to determine type of connector)
  • Cable Length Required
  • Preferred Body Material (Glass or Plastic)

Section 3: Calibration of pH Meter and Electrode

To achieve accurate, reproducible results a great deal of attention needs to be paid to the calibration method. A decision should be made on the accuracy required for the measurement. This will enable the user to choose the type of calibration required and the appropriate type of equipment to be used. The following recommendations will ensure the best levels of accuracy.

  • All solutions should be stirred gently to ensure the sensor is measuring a true representation of the beaker contents.
  • Calibration buffers should be chosen which have pH values either side of the expected sample value, i.e, for a sample which has an expected pH of 5, pH buffers with a value of pH 7 and pH 4 should be used.
  • Always use a 'control' buffer to keep a check on the drift of the electrode. A method commonly used is to place the electrode into a buffer, which has a value close to the sample pH, between measurements.
  • Fresh buffer solutions should be used. Changing all solutions daily is a good practice.
  • All solutions should be maintained at an equal temperature.
  • Rinse the electrode thoroughly in deionised water between measurements.
  • When calibrating the electrodes, allow sufficient time to elapse for the reading to stabilise before adjusting the meter. At least one minute, preferably longer.

Section 4: Procedure for Calibrating the pH Meter

  • 1 x high quality pH/mV meter
  • 1 x 100ml pH 7.00 buffer solution
  • 1 x 100ml pH 4.00 buffer solution
  • 1 x 100ml pH 5.00 buffer solution
  • 1 x calibrated glass thermometer
  • 1 x temperature controlled water bath (required if the sample value is different to ambient)
  • 1 x combination pH electrode
  • 4 x 200ml beakers
  • 3 x Teflon stirrer paddles
  • 1 x magnetic stirrer
  • 1 x cantilever arm electrode stand
  • 1 x fast flow wash bottle containing deionised water


  • a) Assemble all equipment.
  • b) Lower fill hole sleeve on the electrode (if fitted) and thoroughly rinse the electrode tip.
  • c) Lower electrode into gently stirred pH 7.00 buffer and allow to stabilise.
  • d) Check the temperature of the calibration solutions and adjust the default reading on the pH meter, if applicable.
  • e) After 1 - 2 minutes adjust the calibration control on the pH meter to the appropriate pH value.
  • f) Raise electrode from beaker and thoroughly rinse with deionised water.
  • g) Lower electrode into gently stirred pH 4.00 buffer and allow to stabilise.
  • h) After 1 - 2 minutes adjust the slope control on the pH meter to the appropriate temperature corrected value.
    NOTE: Many modern microprocessor controlled pH meters have automatic buffer recognition.
    Please consult the instrument manual for specific adjustment information.
  • i) Rinse the electrode and repeat stages c) to h) to confirm calibration.
  • j) Rinse the electrode and lower into pH 5.00 buffer.
  • k) After stabilising, record the reading in pH 5.00 buffer.
  • l) Between measurements in the sample, rinse and lower the electrode into the control buffer (pH 5.00) for comparison with the recorded reading (remember to check temperature pH versus pH values).

Section 5: Care and Maintenance of Electrodes

By following this advice, it is possible to significantly increase the expected life of an electrode and also to improve the quality of measurement results.

  • pH electrodes must always be stored wet. There are many opinions on which storage solution is the best. Russell Mainstream electrodes are all supplied soaked in a saturated KCl solution with the exception of double junction electrodes which are stored in the appropriate refill electrolyte for their application.
  • For short term storage, soak the electrode in KCl.
  • For long term storage, fill the soaking boot, fit over the end of the electrode and seal with parafilm.
  • Electrodes should never be stored in any of the following liquids:
    Deionised water, sample solutions, solvents, hydrochloric acid, pH buffers containing mercury based preservatives.
  • Sensing tips should always be rinsed after use.
  • Reference cells should always be kept regularly topped up with electrolyte.
  • Connectors must be kept clean and dry.
  • If the electrode needs to be cleaned physically, always use soft tissue soaked in a mild detergent or propanol.
  • Regularly inspect the glass pH sensitive membrane for cracks or chips.
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Introduction to Conductivity


Conductivity is the ability of a material to conduct electric current. The principle by which instruments measure conductivity is simple - two plates are placed in the sample, a potential is applied across the plates (normally a sine wave voltage), and the current is measured. Conductivity (G), the inverse of resistivity (R) is determined from the voltage and current values according to Ohm's law. 

G = I/R = I (amps) / E (volts)

Since the charge on ions in solution facilities the conductance of electrical current, the conductivity of a solution is proportional to its ion concentration. In some situations, however, conductivity may not correlate directly to concentration. The graphs below illustrate the relationship between conductivity and ion concentration for two common solutions. Notice that the graph is linear for sodium chloride solution, but not for highly concentrated sulphuric acid. Ionic interactions can alter the linear relationship between conductivity and concentration in some highly concentrated solutions.

Units of Measurement

The basic unit of conductivity is the siemens (S), formerly called the mho. Since cell geometry affects conductivity values, standardized measurements are expressed in specific conductivity units (S/cm) to compensate for variations in electrode dimensions. Specific conductivity (C) is simply the product of measured conductivity (G) and the electrode cell constant (L/A), where L is the length of the column of liquid between the electrode and A is the area of the electrodes (see Figure 1).  

C = G x (L/A)

If the cell constant is 1 cm-1, the specific conductivity is the same as the measured conductivity of the solution. Although electrode shape varies, an electrode can always be represented by an equivalent theoretical cell.

The following table shows optimum conductivity ranges for cells of three different constants:

Cell constant Optimum Conductivity Range (S/cm)
0.1 0.5 to 400
1.0 10 to 2000
10.0 1000 to 200,000

Conductivity Temperature Compensation

Conductivity measurements are temperature dependent. The degree to whcih temperature affects conductivity varies from solution to solution and can be calculated using the following formula:

Gt = Gtcal {1 + a(T-Tcal)}

where: Gt = conductivity at any temperature T in C, Gtcal = conductivity at calibration temperature Tcal in C, a = temperature coefficient of solution at Tcal in C.

Substance at 25C Concentration Alpha (a)
HCl 10 wt% 1.56
KCl 10 wt% 1.88
H2SO4 50 wt% 1.93
NaCl 10 wt% 2.14
HF 1.5 wt% 7.20
HNO3 31 wt% 31.0

Common alphas (a) are listed in the table above. To determine that a of other solutions, simply measure conductivity at a range of temperatures and graph the change in conductivity versus the change in temperature. Divide the slope of the graph by Gtcal to get a.

All meters have either fixed or adjustable automatic temperature compensation referenced to a standard temperature - usually 25C. Most meters with fixed temperature compensation use a a of 2%/C (the approximate a of NaCl solutions at 25C). Meters with adjustable temperature compensation let you to adjust the a to more closely match the a of your measured solution.

Conductivity Meter Calibration and Cell Maintenance

Conductivity meters and cells should be calibrated to a standard solution before using. When selecting a standard, choose one that has the approximate conductivity of the solution to be measured. The conductivity of some common solutions is shown in the table below.

Solution Conductivity
Absolute pure water 0.055 S/cm
Power plant boiler water 1.0 S/cm
Good city water 50 S/cm
Ocean water 53 mS/cm

A polarized or fouled electrode must be cleaned to renew the active surface of the cell. In most situations, hot water with a mild liquid detergent is an effective cleanser. Acetone easily cleans most organic matter, and chlorous solutions will remove algae, bacteria or molds. To prevent cell damage, abrasives or sharp objects should not be used to clean an electrode. A cotton bud works well for cleaning but care must be taken not to widen the distance of cell. 


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Introduction to Total Dissolved Solids (TDS)


Total Dissolved Solids (TDS) are solids in water that can pass through a filter. TDS is a measure of the amount of material dissolved in water. This material can include carbonate, bicarbonate, chloride, sulphate, phosphate, nitrate, calcium, magnesium, sodium, organic ions, and other ions. A certain level of these ions in water is necessary for aquatic life. Changes in TDS concentrations can be harmful because the density of the water determines the flow of water into and out of an organism's cells (Mitchell and Stapp, 1992). However, if TDS concentrations are too high or too low, the growth of many aquatic life can be limited, and death may occur.

Similar to TSS, high concentrations of TDS may also reduce water clarity, contribute to a decrease in photosynthesis, combine with toxic compounds and heavy metals, and lead to an increase in water temperature. TDS is used to estimate the quality of drinking water, because it represents the amount of ions in the water. Water with high TDS often has a bad taste and/or high water hardness, and could result in a laxative effect.

The TDS concentration of a water sample can be estimated from specific conductance if a linear correlation between the two parameters is first established. Depending on the chemistry of the water, TDS (in mg/l) can be estimated by multiplying specific conductance (in micromhos/cm) by a factor between 0.55 and 0.75.

TDS can also be determined by measuring individual ions and adding them up.


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Introduction to ORP (REDOX)

Oxidation-Reduction Potential (ORP) or Redox potential measurements are used to monitor chemical reactions, to quantify ion activity, or to determine the oxidizing or reducing properties of a solution. ORP is a measurement of the electrical potential of a redox reaction and serves as a yardstick to judge how much oxidation or reduction takes place under existing conditions. 

ORP electrodes measure the voltage across a circuit formed by the measuring metal half cell and the reference half cell. When the ORP electrode is placed in the presence of oxidizing or reducing agents, electrons are constantly transferred back and forth on its measuring surface, generating a tiny voltage. The ORP measurement can be made using the millivolt mode of a pH meter.

ORP measurement may be utilized very successfully in many commercial and industrial applications. These include:

  • Cyanide Oxidation

  • Aquarium Monitoring

  • Chromate Reduction

  • Drinking Water

  • Swimming Pool Water

  • Pulp Bleaching

  • Cooling Tower

  • Ozone Monitoring

  • Water Pollution Monitoring

ORP technology has been gaining recognition worldwide and is found to be a reliable indicator of bacteriological water quality for sanitation - determine free chlorine parameter. In swimming pool application, the ideal ORP value is approximately 700 mV where the Kill Time of E.Coli bacteria is the fastest to ensure good water quality. However ORP value also depends on the pH of pool water, which is typically between 7.2 and 7.6 pH. 

The pH of pool water has to be maintained at optimum level by dosing appropriate chemicals. If the pH of swimming pool is acceptable and ORP value is below 700 mV, then hypochlorite or other oxidizing chemicals need to be added.

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Introduction to Salinity

Salinity refers to the concept of saltiness of a body of water. Water is defined as saline if it contains 3 to 5% salt by volume. The ocean is naturally saline at ~3.5% salt (see sea water). Some inland salt lakes (seas) are even saltier. The technical term for ocean saltiness is halinity, from the fact that halides (chloride, specifically) are the most abundant anion in the mix of dissolved solids. In oceanography, it has been traditional to express salinity as concentration in parts per thousand (ppt or o/oo), which is grams salt per litre of water.

The salt content of most lakes, rivers, and streams is so small that these waters are termed fresh or even sweet water. Salt is difficult to remove from water, and salt content is a factor in water potability. Salinity is an ecological factor of considerable import, influencing the types of organisms that live in a body of water. As well, salinity influences the kinds of plants that will grow either in a water body, or on land fed by a water (or by a groundwater). A plant adapted to a saline conditions is called a halophyte (for salt loving). See also biosalinity. Animals and bacteria that can live in very salty conditions are classified as extremophiles.

Systems of Classification of water bodies based upon Salinity

Marine waters are those of the ocean, another term for which is euhaline seas. The salinity range for euhaline seas is 30 to 35 o/oo. brackish seas or waters have salinity in the range of 0.5 to 29 o/oo; and metahaline seas from 36 to 40 o/oo. These waters are all grouped as homoiohaline because their salinity is derived from the ocean (thalassic) and essentially invariant, in contrast to poikilohaline environments in which the salinity variation is biologically significant (Dahl, 1956).

Poikilohaline waters may range anywhere from 0.5 o/oo to greater than 300 o/oo. The important characteristic is that these waters tend to vary in salinity over some biologically meaningful range seasonally or on some other roughly comparable time scale. Put simply, these are bodies of water with variable salinity. The following table, modified from Por (1972) follows the "Venice system" (1959):

>300 o/oo --------------------
60 - 80 o/oo --------------------
40 o/oo --------------------
30 o/oo --------------------
18 o/oo --------------------
5 o/oo --------------------
0.5 o/oo --------------------


Highly saline water is referred to as brine.


  • Dahl, E. 1956. Ecological salinity boundaries in poikilohaline waters. Oikos, 7(I): 1–21.
  • Por, F. D. 1972. Hydrobiological notes on the high-salinity waters of the Sinai Peninsula. Mar. Biol., 14(2): 111–119.
  • Venice system. 1959. Final resolution of the symposium on the classification of brackish waters. Archo Oceanogr. Limnol., 11 (suppl): 243–248.

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