Why Choose a pH Meter instead of a pH Test Strip

Wednesday, November 29, 2017 8:13 PM

Test strips or sticks used to measure pH are inexpensive and easy to use. So why would you want to use a pH meter, which is more expensive and takes more time to use?

1. Resolution

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.


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Why pH is important?

Wednesday, November 15, 2017 12:00 AM

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[1]. 

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.

Water treatment

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.

Food Industry

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.


[1] Brady, Nyle C., and Ray R. Weil. 2002. The nature and properties of soils. Upper Saddle River, N.J.: Prentice Hall.


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What is pH?

Friday, October 27, 2017 12:00 AM

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 6.022×1023 objects, 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 10-6 or 0.000001 mol/L. At pH 3, the concentration is 10-3 or 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)[1], 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 10-11 or 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.


H+ concentration, mol/L

OH- concentration, mol/L














































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[2].


[1] Langmuir, D. (1997). Aqueous environmental geochemistry. Upper Saddle River, N.J: Prentice Hall.

[2] 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.


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