What is ORP (Oxidatio-Reduction Potential)

Monday, March 12, 2018 8:55 PM

What is Oxidation-Reduction Potential (ORP)?

Oxidation-reduction potential (ORP) or redox is a measurement that indicates how oxidizing or reducing a liquid is. For example, water may be moderately oxidizing (such as aerated water), strongly oxidizing (such as chlorinated water or hydrogen peroxide solution), or reducing (such as an environment where anaerobic microbes are active). In short, ORP is a measure of the cleanliness of the water and its ability to break down contaminants. This measurement has a variety of applications, such as checking for safe sanitation of drinking water or monitoring fluid for the suitability for anaerobic microbial processes.

What are oxidation and reduction?

Oxidation and reduction are related chemical processes that refer to the exchange of electrons in a reaction. Oxidation refers to when a chemical loses electrons. Reduction refers to when a chemical gains electrons, so reduction is the opposite of oxidation. Both oxidation and reduction can happen in the same reaction, which is why reactions involving oxidation and reduction are often called redox reactions.

As an example, let’s look at the reaction of oxygen gas with hydrogen gas to form water:

O2 + 2H2  -- 2H2O

If we look closer at the water molecule,  writing it as (H+)2(O-2), it can be viewed as a combination of two ions, O-2 and H+, that have electrical charges because they gained or lost electrons:

2H+ + O-2  -- (H+)2(O-2)

Electrons have a negative charge, so the oxygen atom in the water molecule gained two electrons to end up with a −2 charge:

O + 2e- -- O-2

In the above reaction, the oxygen atom was reduced because it gained electrons.

Each of the two hydrogen atoms in the water molecule lost an electron to end up with a +1 charge:

H2 -- 2H+ + 2e-

In this reaction, the hydrogen atoms were oxidized because they each lost an electron.


Oxidation or Reduction?

O + 2e- -- O-2

The oxygen atom gains electrons.
The oxygen atom is reduced.

H -- H+ + e-

The hydrogen atom loses an electron.

The hydrogen atom is oxidized.

O2 + 2H2 --2H2O

The oxygen atoms are reduced.

The hydrogen atoms are oxidized.


In the reaction of oxygen and hydrogen gas to form water, the oxygen accepts electrons from the hydrogen, so we can say that the hydrogen is oxidized by the oxygen. Likewise, we canalso say that the oxygen is reduced by the hydrogen.

Some common oxidation processes include decomposition of organic matter and conversion of iron to rust (iron oxide).

Electrons and the ORP scale

From the above discussion, one might guess where the word “oxidize” comes from. Oxygen gas is very good at accepting electrons from other atoms, and this is indeed the most common type of oxidation process that occurs in the environment. From this, we might also suppose that an environment that contains oxygen gas is an oxidizing environment. In such an environment, iron will turn to rust, and aerobic respiration can occur.

One might also guess that a reducing environment is an environment without oxygen gas. Such an environment often includes dissolved gases that are products of anaerobic activity, such as methane, hydrogen sulfide, and hydrogen.

Chemicals (such as oxygen) that accept electrons from other compounds are called oxidizing agents, and substances (such as methane or hydrogen) that give up electrons are called reducing agents.

The degree to which a fluid is oxidizing or reducing (represented by ORP) depends on the presence and strength of various oxidizing and reducing agents. ORP can also be thought of as representing the availability of electrons. Because reducing agents give up electrons, a reducing environment is one where electrons are relatively available. In contrast, an oxidizing environment is one where electrons are relatively unavailable.

ORP is expressed as an electrical potential (a voltage). Generally speaking, a reducing environment is indicated by a negative reading, and an oxidizing environment is indicated by a positive reading. The most common unit for expressing ORP is the millivolt (mV), and most meters can read values ranging from -1000 mV to +1000 mV. The more extreme the negative or positive value, the more reducing or oxidizing the fluid is.


Different oxidation-reduction processes and conditions have different ORP values, with aerobic conditions having higher ORP values and anaerobic conditions having lower ORP values. \


Applications of ORP measurement

One of the biggest applications of ORP is in water disinfection. Municipal drinking water supplies, for example, use strong oxidizers such as chlorine to kill bacteria and other microbes and to prevent their growth in water supply lines. Higher ORP values are associated with higher concentrations of the disinfectant, so ORP is used to monitor and control disinfectant levels in water supplies. In swimming pools and spas, disinfectants are used to kill microbes that may transmit diseases. In outdoor swimming pools and cooling towers, disinfectants are also used to prevent the growth of algae.

ORP is also used for monitoring and control of many oxidation-reduction reactions in industrial processes. For example, in automated industrial systems, ORP is often used to maintain a slight excess of oxidizing chemicals such as chlorine, hydrogen peroxide and ozone, or reducing chemicals such as sulfur dioxide and sodium sulfite.

In wastewater treatment, ORP is used to determine the types of microbial processes that are occurring and to help operators manage the treatment system by promoting or preventing certain reactions. For example, ORP may be controlled in various parts of a system to digest organic matter, remove nitrate or phosphorus, and control odors.

Because low values of ORP indicate anaerobic conditions, ORP can be used to detect anaerobic microbial activity in the environment, such as in the water column or in sediment. ORP can also be used to indicate soil saturation, which makes it useful for mapping wetlands[1].

In other environmental applications, ORP measurements can be viewed as an extension of the dissolved oxygen (DO) scale[1]. DO meters can cover the range of aerobic conditions, but they cannot indicate how reducing an anaerobic environment is. The ORP scale, on the other hand, covers a wide range of reducing conditions. Because of this, ORP can provide insight into the chemistry of anaerobic environments, such as the types of microbial processes in sediments or reactions involving pollutants in contaminated aquifers.

ORP can also be used in conjunction with membrane DO sensors to identify conditions where the DO measurements may be faulty[1]. Under anaerobic conditions, membrane-type DO sensors may give false readings if sulfides are present. If the ORP measurement indicates anaerobic conditions, then positive DO measurements taken from these types of sensors should be considered suspect.


ORP is a fast and inexpensive measurement of the oxidizing and reducing conditions in an environment or system. This makes ORP measurement suitable for a wide range of industrial and environmental applications where oxidizing and reducing conditions vary. ORP is especially useful for routine or continuous monitoring situations where slower and more expensive chemical tests would not be as practical.


 [1] U.S. Environmental Protection Agency (2017) Field measurement of oxidation reduction potential (ORP). SESDPROC-113-R2.


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How to Test Dissolved Oxygen (DO)?

Monday, February 26, 2018 8:29 PM


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What is dissolved oxygen (DO)?

Monday, February 12, 2018 8:20 PM

Dissolved oxygen (DO) is oxygen gas (O2) that is dissolved in water. Gases in the atmosphere, such as oxygen, nitrogen and carbon dioxide, naturally dissolve in water to some degree. Like salt or sugar, these gases are invisible in water once they become dissolved.

The element oxygen exists in many forms in nature. Although most people know that oxygen is part of the water molecule, most would be surprised to hear that oxygen is also the most abundant element in rocks. In these forms, oxygen is bound to other elements such as hydrogen, silicon or carbon. Molecular oxygen (O2), which is in air, is different than other forms because it is not bound to other elements. In nature, the O2 that we breathe is chemically much more reactive than the more abundant forms of oxygen that we come in contact with. This is what allows plants, animals and other organisms to use O2 to metabolize their food through the process of respiration.

Concentration and solubility

The amount (concentration) of O2 dissolved in water is most often expressed in terms of milligrams per liter of water (mg/L). This concentration is referred to as the dissolved oxygen (DO) content of the water. There is a natural tendency for water in contact with air to dissolve O2 until the saturation concentration is reached. For example, the DO in fresh water at 25°C in contact with air is 8.3 mg/L, assuming that equilibrium between water and air is reached and that nothing is removing the O2 from the water. 

DO concentrations are sometimes expressed as % of saturation. If the DO of the water is at the saturation concentration, then it is said to be 100% saturated. If the DO is 5.0 mg/L in fresh water that is at 25°C, for example, then it is 60% saturated (5.0 divided by the saturation level of 8.3 mg/L, multiplied by 100%).

This saturation concentration is known as the solubility of O2, which is the amount of O2 that water can hold. The solubility of O2 changes with temperature, salinity and pressure. The solubility of O2 in water increases as the temperature decreases, meaning that cold water can hold more O2. For example, cold water at 5°C (12.8 mg/L) holds about 55% more dissolved oxygen than warm water at 25°C (8.3 mg/L)[1]. 

Because the temperature of water varies with the seasons, DO levels tend to be higher in the cooler months because the solubility of O2 is higher in cold water. In the summer, water levels tend to be lower and the air is warmer, which leads to warmer water and lower DO levels.

The salinity of water also affects the solubility of O2, such that seawater can hold about 20% less O2 than fresh water[1].

Source: [1]

Dissolved oxygen solubility changes with temperature and salinity.

Pressure also affects the solubility of O2. The water pressure at a certain depth depends on the height of the water column above it, so pressure increases with depth. Water at greater pressure can hold more O2, meaning that the solubility of O2 increases at greater depths. For example, water at 4 m (13.1 ft) depth can hold about 40% more O2 than water at the surface[2]. 

It is possible for water to have a DO level that is higher than the solubility of O2 (more than 100% saturation). This condition is called supersaturation, which can happen under special circumstances (see below).

Sources and sinks of O2 in water

The main source of O2 in water is the atmosphere. Oxygen molecules slowly enter water at the water surface. This process is aided naturally by turbulent flowing water, wind, and waves. Because of this, still water tends to have lower DO values than rapidly moving water. Aeration of water naturally by rapids or waterfalls, or artificially by bubbling air through water, turning waterwheels, or spilling through dams, greatly accelerates the transfer of O2 from air to water. O2 also enters water bodies from tributary streams and groundwater discharge.

O2 in water is also produced through photosynthesis, in which plants and algae convert dissolved carbon dioxide (CO2) into organic matter, releasing O2 into the water. Photosynthesis only takes place at times of day where light is present. The depth at which photosynthesis takes place depends on the clarity of the water. In murky water, light may not reach the bottom of a deep lake.

Aquatic plants, animals and microbes consume O2 by respiration, where organic material used as fuel is converted back into CO2; this is the opposite of photosynthesis. Many people are surprised to learn that plants consume O2 as well as produce it. Plants will actually consume O2 by respiration at night and release O2 through photosynthesis during the day. Because of this, DO in some aquatic environments will tend to decrease at night and increase in the daytime.

Microbes and fungi also consume O2 through the decomposition of dead organic matter. Often, this happens in deeper layers of the water column as dead material sinks toward the bottom. Because of this, deeper layers of water often have lower levels of DO than shallow layers.

DO and aquatic life

Different species of aquatic animals have different DO requirements. Animals that feed on the bottom of a water body, where DO levels tend to be lower, can typically tolerate lower DO levels that animals that dwell near the surface. Most fish are able to survive and grow at DO concentrations of 5 mg/L or higher, although spawning and optimal growth may require higher concentrations[3].

When DO levels are too low for a certain species, the animal can become lethargic or die. Hypoxia is a condition where DO is low enough to threaten aquatic animal species. Hypoxia can cause dead zones in water bodies, where fish and other aquatic life are absent. A DO level of less than 1-2 mg/L is generally considered hypoxic, and a level less than 3 mg/L is a cause for concern. These values are below the requirements for spawning and growth of most fish.

At the opposite extreme, supersaturation of water with O2 can lead to health problems in fish. Supersaturation arises when the solubility of O2 in water rapidly decreases or when O2 is rapidly produced by photosynthesis. The solubility of O2 can decrease when water temperature rises, for example, so a rapid rise in water temperature can lead to supersaturation. Supersaturation with O2 can cause a health condition in fish called gas bubble disease.

Environmental impacts on DO

Because dissolved O2 is needed by most aquatic organisms, the DO of a water body is often used to assess its health. DO levels in water bodies can be impacted by a number of different environmental problems. For example, runoff associated with clearcutting or agricultural wastes can carry excessive organic material into water bodies, which can result in the depletion of O2 as the material is decomposed. 

Another problem is excessive nutrients, which can enter water bodies through runoff associated with fertilizer application on agricultural or recreation land (such as golf courses) or from wastewater treatment plants. Excessive nutrients can result in algal blooms, a process known as eutrophication. Algal blooms can block light from reaching aquatic plants, and dead algae provide a source of organic matter that can deplete DO levels when it decomposes. Because the dead algae sink, this problem especially impacts deeper layers of water and animals that dwell on the floor or bed of the water body.

Riparian vegetation (plants that live along the banks of a stream or river) protects the DO of streams by providing shade that helps keep the water cool. When this vegetation is removed, however, the temperature of the water can increase, causing a corresponding drop in DO levels.

The temperature of water can also be affected by other human activities. When water is withdrawn or stored for drinking water, irrigation, or industrial use, especially during dry months, the water level in streams can decrease, making them especially susceptible to temperature fluctuations and warming. The resulting decrease in DO can harm aquatic life in these water bodies. When water is used for industrial cooling processes and then discharged back into a stream, its temperature is often higher than the water in the stream, resulting in warming of the stream and a decrease in its DO.


Dissolved oxygen is affected by many different factors and processes found in water bodies, and it can fluctuate over short time scales. Fortunately, most aquatic life can tolerate short periods where DO is low. However, persistent problems with low DO levels can lead to poor health of an aquatic environment. This is why routine monitoring of DO is important when there is concern about the health of aquatic life.


[1] American Public Health Association (APHA) (2005) Standard methods for examination of water and wastewater, 21st edn. APHA, AWWA, WPCF, Washington.

[2] FAO. (2014). Site selection for aquaculture: Chemical features of water. Washington, DC: Fisheries and Aquaculture Department, www.fao.org.

[3] U.S. Environmental Protection Agency (1986) Ambient water quality criteria for dissolved oxygen. EPA 440/5-86-003.


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Testing Conductivity

Monday, January 29, 2018 8:19 PM

With an appropriate instrument, electrical conductivity (EC) measurements are relatively fast and simple. EC measures the ability of water to conduct an electric current, which in turn depends on the concentrations of ions in the solution. Because of this, EC provides useful information about the solution and can be used to estimate its total dissolved solids (TDS). 

The conductivity of water is measured using a probe that is inserted into the water. Using the electrodes in the probe and the electronics in the instrument, the instrument is able to measure the conductivity and report a temperature-compensated conductivity value (units of µS/cm are most typical). To ensure an accurate result, the instrument is usually calibrated with one of more standards prior to the measurement.

Calibration of the instrument

To make accurate measurements, a conductivity instrument is usually calibrated using potassium chloride (KCl) solutions of known concentration. Typically, a standard composed of 0.01 M KCl is used, which has a conductivity of 1412 µS/cm at 25°C [1], but a standard that has a conductivity similar to the solutions being analyzed is ideal. For greater accuracy over a wide range of conductivity values, up to 3-5 standards of different KCl concentrations can be used to calibrate the instrument.

Factors affecting conductivity

There are three main factors that affect the conductivity of a solution: the concentrations of ions, the type of ions, and the temperature of the solution.

1) The concentration of dissolved ions. An electrolyte consists of dissolved ions (such as Na+ and Cl-) that carry electrical charges and can move through water. As each ion is able to carry an electrical charge, water with more ions present is able to conduct a greater amount of current. This is the most important of the three main factors.

2) The types of ions in solution. Different ions have different abilities to transmit charge. Inorganic ions such as Na+, K+, Mg+2, Ca+2, HCO 3-, Cl- and SO4-2, tend to conduct electricity well, although each ion has a different ability to conduct electricity. This depends on factors such as the charge of the ion, its size, and its tendency to interact with water molecules. Heavier ions tend to move slower, but small ions can often attract water molecules more strongly, resulting in a slow-moving hydrated ion. For example, the lightweight ion Li + actually moves electricity only about half as well as the heavier K+ ion[2] because of its stronger interaction with water molecules.

Organic substances tend to make poorer electrolytes than inorganic substances largely because they have a relatively weak tendency to dissociate into ions. For example, acetic acid is a weak acid with a tendency to stay in its uncharged CH 3COOH0 form rather than separate into the hydrogen (H +) and acetate (CH3COO-) ions. Because many organic substances are weak acids, the conductivities of solutions containing them will tend to rise as pH increases. This is because organic acids tend to become converted to their ionic forms as the solution becomes more basic.

3) Temperature. This is a relatively small, but significant, effect. Because ions can move faster in warmer water, the conductivity of water increases with rising temperature. Conductivity will increase by approximately 1.9% for each 1°C increase in temperature [1] (or a little more than 1% for each 1°F difference), which makes it necessary to compensate for temperature so that different conductivity measurements can be compared.

Temperature compensation

To make it easier to compare results for samples tested at different temperatures, conductivity measurements are usually reported as temperature-compensated values. This means that the value reported is what the conductivity would be if the temperature was 25°C. For example, the actual conductivity of a solution tested at 20°C will be lower than the reported temperature-compensated value. Temperature compensation is usually done automatically with a built-in thermistor in the conductivity probe. If the conductivity readings are not temperature compensated, especially when the temperature is far away from 25˚C, the results would not be dependable.

Can conductivity be determined without using a conductivity instrument?

As described above, the conductivity of water depends on the type and amounts of charged ions in solution. If the chemical composition of a solution is known, and if the ions present are limited to well-characterized inorganic ions such as Na +, K+, Mg+2, Ca+2, HCO3-, Cl - and SO4-2 or some organic ions, the conductivity of the solution can be calculated based on the conductance properties of each ion. This is most easily accomplished using specialized chemical software such as PHREEQC [3]. However, it is usually simpler and more direct to measure the conductivity with an instrument.


[1] American Public Health Association (APHA) (2005) Standard methods for examination of water and wastewater, 21st edn. APHA, AWWA, WPCF, Washington.

[2] Haynes, W. M. (2009). CRC handbook of chemistry and physics: a ready-reference book of chemical and physical data. Boca Raton: CRC Press.

[3] Parkhurst, D.L., and Appelo, C.A.J. (2013), Description of input and examples for PHREEQC version 3--A computer program for speciation, batch- reaction, one-dimensional transport, and inverse geochemical calculations: U.S. Geological Survey Techniques and Methods, book 6, chap. A43, 497 p.

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Applications of Conductivity Measurement

Monday, January 15, 2018 8:16 PM

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[1].

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.


[1] Abrol, I. P., Yadav, J. S. P., & Massoud, F. I. (1988). Salt-affected soils and their management (No. 39). Food & Agriculture Org.

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