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Interpretive Significance Of Ph


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#1 synack

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Posted 13 February 2010 - 11:31 AM

From elsewhere - credit to the original author (unknown).

Quite comprehensive info on PH and testing of, maybe useful for some. Reformatted for clarity.



Interpretive Significance of pH and Solution Factors Controlling pH

pH values of natural waters are worthless!!

Yes! As a guide for determining how much acid or alkali needs to be added to change the pH by a required amount or as a measure of buffering capacity or corrosivity, experienced water chemists know that for reasons outlined below, that statement is true - especially when applied to natural uncontaminated waters such as scheme / tap waters and especially bore waters, with pH values within the range 4.5 - 8.2. Further, pH values within that range can be very unstable - i.e. variable over a short time period.

WHAT CHEMICAL CONSTITUENTS CONTROL pH?

a) In Scheme Waters and Bore Waters

Because of its solubility in water, the presence of carbon dioxide (CO2) in the atmosphere has a major influence on the chemistry and pH of water.
The pH of natural, i.e. uncontaminated, waters with pH values between 4.5 - 8.2 is controlled by the concentrations of bicarbonate anion, HCO3- (sometimes referred to as combined CO2) and free carbon dioxide - where the presence of free carbon dioxide in water lowers pH and bicarbonate elevates pH. It is the amount of bicarbonate (i.e. its alkalinity) in a natural water that determines its buffering capacity. Buffering capacity is the amount of resistance the pH of a water shows to additions of acid or alkali.
The presence of free, i.e. uncombined, carbon dioxide tends to lower the pH because it reacts with water to form carbonic acid thus:
CO2 + H2O = H2CO3
Contrarily, the presence of bicarbonate anion elevates pH because it mops up hydrogen ion thus:
H+ + HCO3 - = H2CO3
The overall reaction is represented by:
CO2 + H2O = H2CO3 = H+ + HCO3-
Thus at high CO2 concentrations the reaction is pushed to the right with the production of more H+ (i.e. pH is lowered). High bicarbonate levels (compared to CO2) mop up H+ with the result the reaction shifts to the left and a higher pH value is produced.

However, a complicating factor is that free carbon dioxide concentrations above about 0.5 mg/L in water are unstable when such waters are exposed to the atmosphere at sea level pressures. Under that condition carbon dioxide in excess of 0.5 mg/L will slowly escape from the water into the atmosphere. This is particularly the case with groundwater's which typically have carbon dioxide contents around 50 - 200 mg/L - as a result of biological activity within the aquifer. When these waters are pumped to the surface, the observed pH rises because the excess (acidic) carbon dioxide escapes. The pH will then rise to a stable value solely dependent on the water's bicarbonate content. For example, a bore water with 100 mg/L bicarbonate and 100 mg/L of free carbon dioxide will have an initial pH of 6.3 gradually rising to 8.2 after it has been exposed to the atmosphere and after which the carbon dioxide content has dropped to around 0.5 mg/L.
The same phenomenon although to a much lesser extent (because of their much lower CO2 contents), occurs with scheme (tap) water. Thus the conclusion - because the pH of natural waters are only stable after aeration, it is only the "after aeration" pH value which is stable and has any interpretative significance. To determine that value, aerate the water by tumbling a sample of it from one container to another, 30-40 times prior to measuring its pH.
In conclusion: interpret pH values with caution because a natural water with a lower pH than another may produce the higher pH after both are aerated!!


b) In Hydroponic Nutrient Solutions

Most commercial hydroponic nutrient concentrates contain no artificial pH buffers, free acids and negligible alkaline impurities such as bicarbonates and carbonates. Therefore hydroponic system pH values are essentially determined by the bicarbonate content of the water supply and the Phosphate content of the nutrient mixture only. Thus, with twin pack liquid nutrient mixtures, the pH of the working nutrient solution is usually determined by the pack containing the Phosphate i.e., the final pH is essentially not influenced by the presence or absence of the other pack in the diluted nutrient mixture.

Optimum pH for Hydroponics


The solubility of nutrients and their availability for uptake by plants in a hydroponic system is greatly influenced by the pH of the nutrient solution. Contrarily, although soil grown plants can grow successfully at relatively high soil pH values this is not the case in hydroponics where for nutrients to remain dissolved and suspended in the solution and therefore mobile, it is important to maintain the pH between 5.0 and 6.0 with an absolute maximum of 6.5.

Nutrients being constantly drawn from the nutrient reservoir and root exudates entering the nutrient solution can change the pH of the nutrient. Consequently the nutrient pH must be checked and adjusted on a regular basis. pH fluctuations are less at larger nutrient tank volumes.
Keep the pH Between 5.0 and 6.0
It is over this compromise pH range that all growth factors are catered for to produce optimal growth. If the pH is allowed to rise much above 6.0, some nutrients, including calcium, phosphorus, sulphate (and the trace elements copper, iron, manganese and zinc) can precipitate thus becoming immobile and unavailable for transport by the water flow to the roots.
The precise pH at which precipitation starts is determined by the combined concentrations of calcium, phosphorus and sulphate. Except for fertiliser water mixture combinations with low concentrations of these nutrients this problem commonly occurs at pH values of around 6.5.
In spite of this precipitation problem, some references advocate pH values well above 6.5 for some plant varieties i.e., conditions which risk depleted concentrations of the above mentioned elements.
Before following such advice you may wish to test for yourself whether or not this problem will occur with your water and nutrient mixture by performing the following simple test:

Adjust the pH of your diluted nutrient solution to your target pH of 6.5, 7.0 etc., and place about 200 ml in a clean, clear glass container. Stir the contents continuously for approximately 1 hour. Then, immediately after briefly stirring, place the glass in front of a bright light and closely examine the contents. The presence of fine white particles or flocculent/gelatinous particles verifies that precipitation has occurred and that that pH value is too high for optimal results. If uncertain of the results, cover the glass with a piece of paper and allow the mixture to stand for 24 hours. Any precipitate will then be evident by the presence of a white deposit either floating on the surface or on the bottom of the container.
Notes:

(a) The stirring over the 1 hour period simulates the water movement in an NFT system and accelerates the rate of precipitation.

(b) The stirring prior to the visual examination is to ensure that all particles lift up from the base of the glass and into the viewing zone for easier detection. Also, it is much easier to detect the presence of small particles when they are moving.
Comment on Common Recommendation of pH 6.2

Although this is a commonly recommended pH value, it has no scientific basis. It appears to have gained a sort of mythology status from the early days when the only cheap means for hobbyists to measure pH was by using the common bromothymol blue pH indicator sold by pet shops for maintenance of the pH of fish tank water. Because the lowest pH value able to be determined by that indicator is about 6.2, those values have unfortunately, become an entrenched recommendation by the hydroponic retail industry.

High pH values (i.e. above 6.0) are to be avoided more than low values - i.e. say, between 4.5 - 5.0.

Methods for Measuring pH


a) Colour Indicators

Although not precise, i.e. they will not distinguish between a pH of, say, 5.2 and 5.3, wide range colour pH indicators having good colour resolution (i.e. colour change red to orange to yellow to green to blue over the pH range 4 to 8) can be fast, extremely accurate, simple, reliable, user friendly and economical.
The method is based on the principle that the colour produced by the particular dye used in the indicator formulation is dependant on the pH of the solution.
A test is conducted by removing a small sample of nutrient from the nutrient tank, adding a drop of the indicator, mixing, then comparing the final solution colour with those on a coloured reference chart. Because of their fundamental accuracy and reliability, wide range pH indicators are the preferred method for measurement of pH in hydroponics.
Note that pool and aquarium pH indicators are usually not suitable for hydroponics because unlike Flairform's indicator, they produce the same colour at pH 4, 5, and 6 and therefore cannot warn of the need to add pH UP when the true pH is, say, 4.

b) Electronic pH Meters

pH meters employing a glass electrode are useful for precise pH measurement but require frequent calibration, proper storage and handling to ensure accuracy and reliability:
The principle on which such meters operate is based on the fact that when glass separates 2 aqueous solutions having different hydrogen ion concentrations, a voltage is developed between the two faces of the glass. The electronic meter itself is simply a very sensitive voltmeter to measure that voltage and is calibrated in terms of pH units instead of volts.
CALIBRATION
To ensure their accuracy, pH meters must be regularly 'calibrated' (i.e. tested and adjusted for accuracy) using, so called, standard pH buffer solutions which are made to an internationally agreed recipe. These are stable solutions that possess a specific pH value.
Proper calibration must always be conducted using at least two pH Buffers that differ in pH by at least three pH units over the pH range being used. This is to ensure that the correct, so called calibration "slope" is obtained.
Further, although it is a common practice amongst growers to only use one calibration buffer (usually buffer pH 7.0) it is a hazardous practice. This is because a normal/common result after calibrating using only pH buffer 7.0, is that if the electrode is placed in a pH buffer 4.0 solution the meter will read, say 4.4, instead of 4.0 - i.e. the software does not have the correct "slope".
This result simply empHasises that two pH buffers, differing by around 3 pH units, must always be used to guarantee proper calibration.
Note that although some manufacturers claim their pH meters require only a single point calibration that claim is simply not justifiable if high accuracy is desired.

pH Meter Problems

a) Accuracy Problems with pH Meters

Note that despite being properly calibrated it has been my frequent experience that different pH meters will produce significantly different pH readings on the same solutions - by as much as 0.5 pH units. This is especially the case with low conductivity waters and drinking waters with low total alkalinity.
This result simply shows that the specificity of the glass electrodes for hydrogen ion is not perfect.

b) Solving pH Electrode Problems

STORAGE

For a pH meter with a single combination electrode to work properly there must be an electrical connection between the electrode filling solution (usually potassium chloride - which is electrically conductive) and the sample solution. This is achieved via the use of a porous frit or wick in the glass wall separating the inner filling solution chamber from the sample. Both these devices allow the free (very slow) flow of the filling solution into the sample solution.
The most common cause of pH meter failure is due to failure of the electrode via blockage of this porous frit or wick. This is usually caused by dry storage of the electrode or leaving the electrode immersed for too long in supersaturated solutions. Dry storage causes dehydration of both the glass electrode and precipitation of salts within the interstices of the frit itself.
The results is reading drift, slow response times and decreased sensitivity.
To avoid these problems the electrode tip must be permanently stored in a specially formulated storage solution such as Flairform's pH Electrode Storage Solution. Further always perform all readings as quickly as practical. Also, do not immerse the electrode in the sample solution to a depth such that the level of the sample is above that of the filling solution on the inside of the electrode. Such a result will permit the sample solution to weep back into the electrode itself.
NOTE: Water or pH buffers are not suitable for long term storage of electrodes.

CLEANING

The glass tip of pH meter electrodes invariably become coated with impurities causing symptoms such as reading drift, slow response times and decreased sensitivity. Contamination may be so severe that re-calibration is not possible without prior cleaning. Minimise this problem by filtering potentially problem (i.e. greasy or highly turbid) water samples prior to performing a pH measurement. Inorganic based film deposits can be removed by using Flairform's pH Electrode Cleaner. Alternatively, grease etc. can be removed with acetone, methylated spirit or toluol. If toluol is necessary, rinse subsequently with acetone or methylated spirits and finally water.
Other more aggressive and successful methods of cleaning have included boiling (!!) nitric acid and caustic soda.
pH Adjustment

Before measuring the pH ensure that the nutrient is well stirred, especially after pH UP and DOWN are used.
If the nutrient pH is too high then add pH DOWN. Phosphoric acid is optimum for this purpose because it increases the nutrient's pH buffering capacity (helps minimise pH fluctuations) and is relatively safe to handle. Alternatively nitric acid has zero buffering capacity and is hazardous to use. Nevertheless, its use is preferable where high alkalinity waters are being used.
If the nutrient pH is too low then add pH UP. It is essential to pre dilute the pH UP dose into about 1 litre of raw water then stir the nutrient as you add this mixture. Failure to do this may cause permanent precipitation of essential nutrients. Also, when using "pH UP", precipitation problems can be minimised by ensuring the nutrient mixture is agitated rapidly near the point where the reagent strikes the surface of the nutrient. Also, if accidental overdosing to above 6.5 occurs, to prevent permanent precipitation and loss of essential elements, reduce the pH back, as quickly as possible, to below pH 6.0.
pH Up: pH Up contains 40% w/v potassium hydroxide. Contains no nuisance chemicals.
pH Down: pH Down contains 80% w/v phosphoric acid. This product is far safer than nitric acid and adds pH buffering capacity to the nutrient to help minimise pH fluctuations.

Edited by synack, 13 February 2010 - 11:33 AM.

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#2 highgrade

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Posted 14 February 2010 - 03:03 PM

interesting I get a fair bit of precipitation when using canna nutes next time I use them ill run them lower 5.8 or so as usually run 6.2


:yep: :P
the same old routine but it never gets old

#3 smokewhisperer

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Posted 04 June 2010 - 10:01 AM

A very informative article, exactly what I was looking for. :spliff:

Thanks synack. lol

#4 Takamine

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Posted 17 June 2010 - 09:43 PM

thanks Synack for that.

It is over this compromise pH range that all growth factors are catered for to produce optimal growth. If the pH is allowed to rise much above 6.0, some nutrients, including calcium, phosphorus, sulphate (and the trace elements copper, iron, manganese and zinc) can precipitate thus becoming immobile and unavailable for transport by the water flow to the roots.


if your res tank ph range does rise to much and precipitation happens does lowering the ph reverse this making nutrients mobile again?
is precipitation visible to the eye?

Edited by Takamine, 17 June 2010 - 09:44 PM.



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