Testing the electrical conductivity of water provides much practical information about a solution. Not only is the conductivity measurement itself useful, but it can also be used to estimate the total dissolved solids (TDS) or salinity of water. Because conductivity measurements are simple, and fast, they are highly suitable for routine testing and long-term monitoring. Some examples of applications of conductivity measurement are described below.
Natural Waters, Aquaculture and Environmental Applications
In natural waters, conductivity is mainly used to estimate the concentrations of dissolved salts in the water, which in can provide insights into processes affecting the water. In river waters, for example, the conductivity (and TDS) of water may increase in the summer when evapotranspiration is high and decrease when the water is diluted by snowmelt or heavy rains. In coastal areas, the conductivity of water may change with mixing with salt water, and the conductivity of water may rise when it becomes contaminated with road salt in areas with cool climates.
For water resources, the conductivity may indicate whether or not the water is too saline to be drinkable or useable for irrigation or industrial use.
In places where there is potential for water to become polluted, the water may be monitored for changes in conductivity that could indicate contamination from a spill or leak. In ecosystems and aquaculture, aquatic plants and animals have certain ranges of salinity that they can tolerate. Because of this, the conductivity of water bodies such as ponds may be monitored to warn if the salinity is in danger of falling outside of the tolerable range for certain fish species, for example.
Water Treatment and Industrial Applications
Water treatment may be used to make water safe to drink or suitable for industrial use. In many industrial applications, scale (precipitation of mineral deposits) or corrosion may be a concern. Because conductivity can be used to estimate the dissolved mineral content of water, it may be used to monitor demineralization processes used to prevent scale or remineralization processes used to prevent corrosion. Conductivity may also be used to monitor the effectiveness of desalinization, which is another water treatment process that removes salts to make water drinkable or useable for industrial processes.
In other industrial applications, conductivity measurements may be used to detect leaks (such as in heat exchangers), where the leaking water may have a higher conductivity. Conductivity may also be used to monitor the effectiveness of rinsing procedures, where a low conductivity of water in contact with the rinsed object indicates an effective rinse. In special circumstances, such as in ammonia solutions, conductivity can even be used to measure pH with more precision than a typical pH meter due to the strong relationship between conductivity and pH.
Agricultural and Hydroponics Applications
For irrigation, the salinity of water is an important factor. If the salinity is too high, salts will accumulate in soil as the water evaporates, which may degrade soil quality and inhibit plant growth. Water with a conductivity of less than 700 uS/cm is acceptable for unrestricted irrigation use, and the use of water with conductivity values greater than 3000 uS/cm should be severely restricted.
Conductivity can also be used to monitor nutrient concentrations in liquid fertilizers. A quick check of the conductivity of liquid fertilizers can guard against mistakes such as improper mixing or malfunctioning injectors, protecting crops from wasteful over-fertilization or inadequate fertilizer application.
Similar to fertilizer application, conductivity is used in hydroponics to monitor the concentrations of nutrient solutions. If the conductivity gets too high, indicating a nutrient concentration at toxic levels, plants may be harmed or die. Low conductivities can indicate inadequate nutrient supply. Conductivity monitoring can be used as part of automated nutrient supply systems. In addition to monitoring nutrient supply, conductivity measurements can be used to make sure that salt concentrations are in the range tolerated by the plant.
Conductivity measurements are simple and fast, making them very practical for making routine assessments of the salt concentrations of water. Whether assessing the concentrations of salts, contaminants or nutrients, measuring conductivity can reduce the need for more expensive or time-consuming tests. There are many factors that affect conductivity, such as the concentrations and types of dissolved salts present in the water, so knowledge of the chemistry of the system in question is often necessary for interpreting conductivity measurements.
 Abrol, I. P., Yadav, J. S. P., & Massoud, F. I. (1988).Salt-affected soils and their management(No. 39). Food & Agriculture Org.
Water has the ability to conduct electricity due to the presence of charged ions in solution. Ions are atoms of molecules that have a net electrical charge, and they include cations (positively charged ions) and anions (negatively charged ions). The most abundant charged ions in natural water typically include the cations sodium (Na+), potassium (K+), calcium (Ca+2) and magnesium (Mg+2) and the anions chloride (Cl-), sulfate (SO4-2), nitrate (NO3-) and bicarbonate (HCO3-). Many other ions can be also found in water, including organic ions and other inorganic ions.
These ions carry electrical charge and can move through water, which allows water to conduct an electrical current. The measure of the ability of water to carry electrical current is called its electrical conductivity. Higher concentrations of ions in water increase its ability to conduct electricity and thus its conductivity. Distilled water, on the other hand, has a very low concentration of ions and a low conductivity.
Technical note: Sometimes electrical conductivity is referred to as specific conductance.
The opposite of conductivity is resistivity. Resistivity is the ability of a material (such as water) to resist the flow of electricity. Resistivity is the reciprocal of conductivity, such that
Resistivity = 1/Conductivity
From this relationship, we can see that water with a high conductivity has a low resistivity, and vice versa. For example, distilled water will have a high resistivity and a low conductivity.
The typical unit for reporting conductivity is microsiemens per cm (µS/cm). This unit is also sometimes written as micromhos per cm (µmho/cm), where 1 µS/cm equals 1 µmho/cm. Potable water typically has conductivity values ranging from 50 to 1500 µmho/cm. At higher conductivities, the water starts to become too salty to drink.
Technical note: Notice that “mho” is the reverse spelling of “ohm,” the common unit for electrical resistance.
Because conductivity varies slightly with temperature, conductivity values are usually reported as temperature-compensated values that represent what the conductivity would be at 25°C. This makes it easier to compare conductivity values for samples with different temperatures.
How is conductivity related to total dissolved solids (TDS)?
Total dissolved solids (TDS) refers to the total amount of dissolved material present in water. TDS is usually reported in milligrams per liter (mg/L) or ppm (parts per million). This means that, if one liter of water with a TDS of 500 mg/L was completely evaporated, 500 mg of solid residue would be left behind. Usually, the dissolved solids include mostly dissolved mineral ions such as sodium, chloride, and the other ions mentioned above. TDS can also include other inorganic ions, dissolved organic material, and non-ionic matter such as dissolved silica. Although a relatively small amount of the TDS includes non-ionic matter that does not carry electrical charge, waters with higher values of TDS generally have higher values of conductivity.
Because of this, a measurement of conductivity (which is quick and easy) can be used to estimate TDS (which is more expensive and time-consuming to measure directly). However, the relationship between conductivity and TDS varies with the chemistry of the water because ions differ in their ability to transmit electrical charge through water. Some ions carry electrical charges faster than others because of factors such as the size and mass of the ions and how they interact with water molecules.
The general equation for estimating TDS from conductivity is as follows:
TDS (mg/L) =k· EC (µS/cm)
where EC is electrical conductivity, andkis the conversion factor, which is related to the chemical composition of the water.
For typical natural waters such as stream and lake water, the value of the conversion factor is usually between 0.6 and 0.7, and a value of 0.64 is considered to be typical. For a solution containing mostly sodium and chloride ions, values of 0.49 to 0.56 are typical, depending on the concentration of salt.
For a precise estimate of TDS from conductivity, the chemistry of the solution should be considered in the selection of the conversion factor. If the composition of the solution is known, then the true TDS of a representative sample of water can be calculated by taking the sum of the measured concentrations. Alternatively, the true TDS value of a representative sample can be directly measured. The correct value of the conversion factor can then be calculated based on the true TDS and the measured conductivity.
If the correct value of the conversion factor cannot be calculated, then a typical or default value of the conversion factor (such as 0.64) will result in a TDS estimate that is at least in the right ballpark.
How is conductivity related to salinity?
Salinity refers to the salt content of water. Because most dissolved solids typically consist of inorganic ions, which are the components of salts, the concepts of salinity and TDS are very similar. In fact, the two concepts are sometimes considered to be synonymous. However, salinity is often expressed in terms of mass of salt per mass of water. For example, ocean water typically has about 35 grams of salt in one kilogram of water, so its salinity can be expressed as 35/1000 or 0.035. This can also be expressed as 3.5% or 35 parts per thousand (ppt).
Salinity is often used to describe seawater and brackish water, but it can also be used to describe fresh water and brines. Because the proportions of the most important ions in seawater are nearly constant, oceanographers can use very precise formulas to estimate salinity from electrical conductivity and temperature.
In cases where salinity is measured in mg/L (for example, for lake water, swimming pools, or irrigation water), salinity can be estimated from electrical conductivity using the same formula presented for TDS in the previous section.
 American Public Health Association (APHA) (2005) Standard methods for examination of water and wastewater, 21st edn. APHA, AWWA, WPCF, Washington.
Test strips or sticks used to measure pH are inexpensive and easy to use. So why would you want to use apH meter, which is more expensive and takes more time to use?
Resolution refers to the smallest measureable change in the quantity being measured. For pH strips or digital pH meters, this refers to the intervals over which the reading is given. For most pH meters, the resolution is at least 0.1 pH units, and resolutions of 0.01 and 0.001 units are common.
For test strips, the resolutions are much lower. Typical pH strips that cover a wide pH range have colors that are spaced 0.5 or 1.0 pH unit apart. Specialty pH strips that cover a narrow range can have color graduations as low as 0.2-0.4 pH units apart. Interpolation between colors is possible, but even the lowest resolution pH meters generally provide better resolution than the highest resolution pH strips.
2. Precision and Accuracy
Precision refers to the reproducibility of a measurement. That is, a more precise measurement technique will yield measured values that are closer to each other as the measurement is repeated. Because the resolution forms the upper limit of precision, low-resolution measurement techniques such as test strips have less potential to be precise compared with higher resolution techniques such as pH meters. Although higher resolution does not necessarily mean higher precision, pH meters generally have higher precision than pH strips.
Accuracy refers to the amount of uncertainty in a measurement, such that a technique with greater accuracy will yield measurements with less uncertainty. An accurate measurement technique will produce measurements that are close to the true value. As with precision, test strips have less potential to be accurate compared with pH meters because of their relatively low resolution.
In summary, pH meters are usually much more precise and accurate than test strips.
2. Objectivity and Consistency
When a test strip is used, it changes color and is compared with a color chart. A problem with this is that different people see colors differently and may tend to interpret the color comparison in different ways. Because of this, two different people may report two different pH values when shown the same test strip. In addition to this, colors appear differently under different lighting conditions. Unless lighting conditions are consistent, this may add to the uncertainty in test strip measurements.
Additionally, interpretation of the test strip requires a certain amount of judgement, especially when interpolation between graduations is involved. For example, if it is not clear which color on the chart is the best match for the test strip, it is a matter of judgement to choose one color or the other, or to report an intermediate value. Because of this, a single operator may not be entirely consistent in making these judgement calls from measurement to measurement. Even the same operator may report different pH values at different times, even if the color does not change.
Test strip results are also subject to variation due to the way the strips are handled. For example, the length of time that the strip is exposed to the solution, the manner in which the strip is handled after it is removed from solution, and the length of time between removing the strip and comparing the colors may all affect the final result. If all of these conditions are not kept consistent, then some additional variation in the results will likely occur.
The subjectivity of test strips can be partly removed by scanning the strip and having the color analyzed by computer, but care must be taken so that user handling and scanning/lighting conditions are consistent. Digital pH meters, on the other hand, produce unambiguous digital output, and measurement protocols using pH meters are relatively easy to keep consistent.
3. Shelf life
Test strips contain reagents that will tend to deteriorate over time. If properly stored under ideal conditions, a set of strips will stay good for approximately one year. However, exposure to moisture, sunlight, and high or low temperatures will tend to degrade the reagents in the strips, which will affect their performance. Test strips that have become degraded may become unacceptably inaccurate.
That said, pH meters also require care. The pH electrode must be properly cared for and stored, and the calibration solutions themselves have a limited shelf life.
4. Continuous Measurement
In some applications, such as acid-base titrations, it is desirable to monitor pH over time in order to detect or record changes. Aside from their other limitations, test strips are not as suitable for this as pH meters because each strip can only be used for one measurement. Digital pH meters, on the other hand, report pH values continuously, and automated recording of measurements is even possible.
In applications where relatively low levels of resolution and accuracy are acceptable, pH test strips may be a good choice because of their low cost and ease of use, as long as relatively fresh strips are used that have been stored under optimal conditions.
Where moderate to high levels of resolution and accuracy are needed, pH meters are most likely the best choice, even though they require a greater up-front investment and more skill to use. When making the choice of measurement technique, one should consider the possible hidden costs of inaccurate measurements, such as addition of excessive material to correct pH, sub-optimal growing conditions, or degradation of the quality of the final product.
pH is an important quantity that reflects the chemical conditions of a solution. The pH can control the availability of nutrients, biological functions, microbial activity, and the behavior of chemicals. Because of this, monitoring or controlling the pH of soil, water, and food or beverage products is important for a wide variety of applications.
Agriculture and gardening
Soil is a complex system that involves many different factors that are affected by soil pH, such as microbial activity, fungal growth, availability of nutrients, and root growth.
Under acidic conditions, many minerals in soil become soluble, releasing toxic metals such as aluminum. Some nutrients, such as phosphorus and molybdenum, become less available at lower pH values. Under alkaline (basic) conditions, the soil can become deficient in nutrients such as zinc, copper, iron, manganese, boron and phosphorus.
Most plants tend to do best in the pH range of approximately 6.0 to 7.0, which is the range over which the most nutrients are available. However, some plants prefer more acidic or basic conditions, such as blueberries (4.0-6.0) or hyacinth (6.5-7.5).
When soil pH is outside of the desired range, the pH can be altered by adding acidic (e.g., native sulfur) or basic (e.g., lime) material to the soil. To correct the pH of acidic soil by liming, an exchangeable acidity analysis must be conducted so that the required amount of lime can be calculated.
Aquaculture and aquatic ecosystems
Water that has a pH that is too low or too high can be harmful to fish and other aquatic life. At low pH, toxic metals such as aluminum can enter the water in greater concentrations, some nitrogen-bearing chemicals become more toxic, and the metabolic processes of fish can become less efficient. Water with pH below 5 can inhibit reproduction or lead to death, and young fish and other aquatic organisms are especially susceptible. Water with a pH below 6.5 can inhibit growth.
At high pH values (such as >9), most ammonium ions are converted to ammonia, which is toxic to fish. This problem gets worse with higher temperatures. Water with pH between 9 and 10 will tend to inhibit growth, and water with pH of 11 or higher will kill fish.
The pH range of 6.5-9 is acceptable for most fish. In aquariums and other closed aquatic systems, it is important that the water be sufficiently buffered (usually with bicarbonate and carbonate ions) to prevent wild swings in pH.
Whether in treating drinking water or waste water, pH is important. The pH of drinking water should be between 6.5 and 8.5. Low-pH drinking water can degrade pipes, causing toxic metals such as copper and lead to leach into the water supply. Water with a pH that is too high has an unpleasant taste, and the effectiveness of disinfectants such as chlorine is decreased.
In wastewater treatment (e.g., sewage or industrial waste), pH is controlled so that desired chemical or microbial reactions will proceed as efficiently as possible. Operators carefully monitor and adjust pH to respond to changing chemical or microbiological conditions.
Swimming pool maintenance
Swimming pools typically have pH values in the range of 7.2 to 7.8. If the pH is too high, the effectiveness of the chlorine disinfectant becomes too low, making the pool becomes susceptible to algal growth and preventing it from effectively killing viruses and bacteria. If the pH is too low, the water becomes irritating to the eyes and nose, and it may corrode plaster or metal surfaces.
In the food industry, pH is measured to test for quality, to control microbial activity, to control the taste and other properties, and to prolong the shelf life of food. In milk, pH is tested to check for impurities or infection. The pH is also affected by the souring of milk and maturation of cream, and the pH determines whether cheese will be soft or hard. The pH of cream also determines whether butter will be sour or sweet. For production of yogurt, the pH of cultured milk is kept low to maintain a desirable environment for appropriate microbial activity.
The pH of food is also used to monitor its quality. For example, a pH that is too high can indicate degraded meat.
For many foods, the pH must be kept within a narrow range so that the food can be conserved for a longer period. For example, batter for baking bread is acidified to extend the shelf life of the bread, as are sauces such as mayonnaise. When canning low-acid foods (with pH < 4.6), extra care must be taken to kill bacterial spores because they can grow when pH is greater than 4.6, potentially causing botulism.
Brewing and winemaking
Similar to other processes that involve microbial activity, pH affects many different aspects of the beer brewing process. In particular, the mash pH controls the behaviors of several enzymes used in brewing, and it should be between 5.3 and 5.8 for most mashes.
The pH of wine must be kept at a low level to prevent bacteria from degrading the wine. Lower pH wines will tend to mature more slowly and will be less susceptible to spoilage. The pH of wine also affects its taste, as more acidic wines tend to be dry. The pH values for wines usually fall within 3.0 to 4.0, and white wines tend to have lower pH values than red wines.
 Brady, Nyle C., and Ray R. Weil. 2002.The nature and properties of soils. Upper Saddle River, N.J.: Prentice Hall.
Many people are aware that pH has something to do with acids or bases and that it is important for things like water quality, food-making, aquatic life and plants, but not many really understand what it means.
In simplest terms, a pH value indicates how acidic or basic water is. A pH of 7 indicates a neutral solution, a pH greater than 7 indicates a basic solution, and a pH of less than 7 indicates an acidic solution. The farther a pH value is from 7, the more strongly acidic or basic it is. For example, water with a pH of 6 is mildly acidic, and water with a pH of 3 is highly acidic.
If you are satisfied with this answer, you can stop reading here. But if you want to know what pH really means, read on…
So what does pH really mean?
The ‘H’ in pH is the elemental symbol for hydrogen. The ‘p’ can refer to different things in different languages, but the ‘pH’ is most commonly said to mean ‘power of hydrogen’. So what does this mean?
The hydrogen referred to in the pH symbol is actually the hydrogen ion, which is written as H+. A hydrogen atom has one proton (which has a positive charge) and one electron (which has a negative charge). The hydrogen ion is a hydrogen atom that has given up its electron. Because there is no longer a negative charge from the electron to balance the positive charge from the proton, H+has a net positive charge (hence the ‘+’ in the symbol).
Technical note: Sometimes the hydrogen ion is written as the hydronium ion (H3O+).
The presence of H+is what makes water acidic. An acid is a compound that contributes H+to water, either directly or indirectly. For example, hydrochloric acid (HCl) dissociates into the hydrogen (H+) and chloride (Cl-) ions in water, and it is the H+that contributes to the acidity of the water. The higher the concentration of H+, the more acidic the water is.
HCl ® H++ Cl-
The concentration of H+can be expressed in terms of moles per liter of water (one mole is 6.022×1023objects, also known as Avogadro’s number). The value of pH reflects the concentration of H+as follows:
(concentration of H+in mol/L) = 10-pH
For example, at pH 6, the concentration of H+is 10-6or 0.000001 mol/L. At pH 3, the concentration is 10-3or 0.001 mol/L.
Technical note: Professional chemists may use slightly different definitions of pH that account for the non-ideal behavior of the hydrogen ion or that use different concentration units (such as moles per kg of water), but the above relationship explains the meaning of pH well enough for many people.
What about neutral and basic solutions?
Water molecules break apart into two ions, the hydrogen ion (H+) and hydroxyl ion (OH-):
H2O = H++ OH-
This reaction is reversible, so that H+and OH-can recombine to form water molecules. At any given time, the concentrations of H+and OH-are very small compared to the amount of water molecules.
It happens that at pH 7, the concentrations of H+and OH-are equal. This is why pH 7 is considered the neutral pH. Below pH 7, the concentration of H+is greater than the concentration of OH-, making the water acidic. Above pH 7, the concentration of OH-is greater than the concentration of H+, making the water basic.
At high pH, the concentration of H+is very small. For example, at pH 11, the concentration of H+is 10-11or 0.00000000001 mol/L. The concentration of OH-, however, is correspondingly higher. The concentration of OH-can be expressed as follows:
(concentration of OH-in mol/L) = 10pH-14
At pH 11, the concentration of OH-is 1011-14(10-3) or 0.001 mol/L, which is the same as the concentration of H+at pH 3.
The pH scale: the power of hydrogen
The pH scale is a logarithmic scale where a difference of one pH unit represents a power of ten difference in the concentration of H+. For example, the concentration of H+at pH 5 is ten times higher than at pH 6, 100 times higher than at pH 7, and 1000 times higher than at pH 8. Because each pH unit represents a power of ten difference in the H+concentration, the pH value represents the “power of hydrogen”. In fact, pH is most often defined as the negative logarithm of the H+concentration:
pH = -log10[H+concentration]
The same general pattern applies to OH-. The concentration of OH-at pH 8 is ten times higher than at pH 7, 100 times higher than at pH 6, and so on.
A common misconception is that the pH scale begins at zero and ends at 14. Actually, negative pH values and pH values greater than 14 are possible. For example, the concentrated acids common in laboratories and industry have negative pH values. Even in the environment, pH values as low as -3.6 have been observed in extreme situations.
 Langmuir, D. (1997).Aqueous environmental geochemistry. Upper Saddle River, N.J: Prentice Hall.
 Nordstrom, D. K., Alpers, C. N., Ptacek, C. J., and Blowes, D. W. (2000) Negative pH and Extremely Acidic Mine Waters from Iron Mountain, California. USGS Staff -- Published Research. 479.