Why pH Varies
The ratio in uptake of anions (negatively charged nutrients) and
cations (positively charged nutrients) by plants may cause substantial
shifts in pH. In general, an excess of cation over anion leads to a
decrease in pH, whereas an excess of anion over cation uptake leads to an
increase in pH. As nitrogen (an element required in large quantities for
healthy plant growth) may be supplied either as a cation (ammonium - NH4+)
or an anion (nitrate - NO3-), the ratio of these two forms of nitrogen in
the nutrient solution can have large effects on both the rate and
direction of pH changes with time. This shift in pH can be surprisingly
fast.
Daylight photosynthesis produces hydrogen ions which can cause the
nutrient acidity to increase (lowering the pH). At dusk photosynthesis
stops and the plants increase their rate of respiration and this coupled
with the respiration of micro organisms and the decomposition of organic
matter uses up the hydrogen ions so the acidity of the solution tends to
decrease (pH rises).
Most varieties of vegetables grow at their best in a nutrient solution
having a pH between 6.0 and 7.5 and a nutrient temperature between 20 and
22OC.
In low light (overcast days or indoor growing environments) plants
take up more potassium and phosphorous from the nutrient solution so the
acidity increases (pH drops). In strong intense light (clear sunny days)
plants take up more nitrogen from the nutrient solution so the acidity
decreases (pH rises).
Extremes in pH can result in precipitation of
certain nutrients.
For plant roots to be able to absorb nutrients, the nutrients must be
dissolved in solution. The process of precipitation (the reverse of
dissolving) results in the formation of solids in the nutrient solution,
making nutrients unavailable to plants. Not all precipitation settles to
the bottom of the tanks, some precipitates occur as very fine suspension
invisible to the naked eye.
Plants can tell us their problems through leaf symptoms (e.g. Iron deficiency) when it's too late. Iron is one essential plant nutrient
whose solubility is affected by pH which is why it is added in a chelated
form (or daily). Iron deficiency symptoms occur readily. At pH values over
7, less than 50% of the Iron is available to plants. At pH 8.0, no Fe is
left in solution due to iron hydroxide precipitation (Fe(OH)3 - which
eventually converts to rust). As long as the pH is kept below 6.5, over
90% of the Iron is available to plants.
Varying pH of summer lettuce nutrient solutions also affects the
solubility of calcium (Ca) and phosphorus (P). Due to calcium phosphate
precipitation (Ca3(PO4)2) the availability of calcium and phosphorus decreases at pH
values above 6.0. All other nutrients stay in solution and do not
precipitate over a wide pH range. Poor water quality could exacerbate any
precipitation reactions that may occur.
Generally in the pH range 4.0 to 6.0, all nutrients are available to
plants. Precipitation reduces Iron, Calcium and Phosphorus availability at pH 6.0 and over
.
Adjusting pH
The addition of acids or alkalis to nutrient solutions is the most
common and practical means to adjust pH, and can be easily automated. See
our hydroponics controller
. There are
ways to minimize pH variations and they are worth some consideration.
Nitrogen is the essential inorganic nutrient required in the largest
quantity by plants. Most plants are able to absorb either nitrate (NO3-)
or ammonium (NH4+) or both. NH4+ as the sole source of nitrogen or in
excess is deleterious to the growth of many plant species. Some plants
yield better when supplied with a mixture of NH4+ (ammonium) and NO3-
(nitrate) compared to NO3- alone. A combination of NH4+ and NO3- can be
used to buffer against changes in pH.
Plants grown in nutrient solution containing only NO3- as the sole
nitrogen source tend to increase solution pH - hence
the need to add acid. But when approximately 10%-20% of the total nitrogen
is supplied as NH4+, the nutrient solution pH is stabilized at pH 5.5.
NH4+ concentration needs to be monitored as it has been shown recently
that micro-organisms growing on plant root surfaces can convert the NH4+
to NO3-. Since hand-held ion-selective electrodes for measuring both NH4+
and NO3- are now available, it should be possible to accurately monitor
and maintain a predetermined NO3-/NH4+ ratio throughout the life of the
crop.
Phosphorus is required in large amounts by plants. Interestingly, there
are two forms of fertilizers containing both K and P - KH2PO4
mono-potassium phosphate (MKP) and K2HPO4 di-potassium phosphate. Equal
quantities of both can be used to maintain the pH at 7.0. Using a higher
proportion of K2HPO4 increases pH. MKP can be used to lower the solution
pH.
Buffers are solutions which resist pH change and are used to calibrate
pH electrodes. Buffers can be added to nutrient solutions in an attempt to
maintain pH stability. One such buffer is called 2-(N-morpholino)
ethanesulfonic acid - abbreviated to MES. Many of the companies who claim
better pH control with their 'specially' formulated nutrient solutions add
MES to their mixes. It is important to remember when using MES, that after
MES addition the pH is low and needs to be adjusted to your required level
with an alkali such as potassium hydroxide (KOH).
Another method of pH stabilization is to use ion- exchange and
chelating resins. Generally, these resins are small beads which have
nutrients absorbed or chelated onto them - the nutrient solution
circulates through the beads or the beads can be suspended in the nutrient
tank. As plants absorb nutrients, more nutrients are released by the
resins. The aim is to achieve controlled release of nutrients into the
solution in an attempt to mimic the way the soil releases nutrients.
Ideally, such release can adequately supply the growing plants'
nutritional requirements and maintain pH stability.
Is pH Adjustment Critical?
pH is not as critical as most hydroponicists believe. The main point is
to avoid extremes in pH. Plants grow on soils with a wide range of pH. For
most plant species there is an optimum pH in the region of pH 5 to pH 6
References - Nothing beats a good book
Asher, C.J. (1978). Natural and synthetic culture media for
spermatophytes. CRC Handbook Series in Nutrition and Food. Ed. Miloslav
Rechigl, Jr. CRC Press Inc., Cleveland, Ohio. Vol 3, 575- 609.
Bates, R.G. (1981). The modern meaning of pH. CRC Critical Reviews in
Analytical Chemistry 11; 247-278.
Errebhi, M. and Wilcox, C.A. (1990). Plant species response to ammonium
- nitrate concentration ratios. Journal of Plant Nutrition. 13 (8):
1017-1029.
Ficks, J. and Mitchell, C.A. (1993). Stabilisation of pH in
solid-matrix hydroponic systems. HortScience, 28 (10): 981-984.
Haynes, R.J. and Goh, K.M. (1978). Ammonium and nitrate nutrition of
plants. Biology Reviews, 53: 465-510.
Huett, D.O. (1994). Growth, nutrient uptake and tipburn severity of
hydroponic lettuce in response to electrical conductivity and K:Ca ratio
in solution. Australian Journal of Agricultural Research 45: 251-267.
Padgett, P.E. and Leonard, R.T. (1993). Contamination of ammonium-based
nutrient solutions by nitrifying organisms and the conversion of ammonium
to nitrate. Plant Physiology 101: 141-146.
Parker, D.R., Norvell, W.A., and Chaney, R.L. (1993). GEOCHEM-PC: a
chemical speciation program for IBM and compatible personal computers. In
R.H. Leoppert, ed, Chemical Equilibrium and Reaction Models, Soil Science
Society of America special publication. Madison, WI (in press).
Sposito, G. and Mattigod, S.V. (1980). GEOCHEM: A computer program for
the calculation of chemical equilibria in soil solutions and other natural
water systems. The Kearney Foundation of Soil Science, University of
California, Riverside, California.
Welch, R.M. (1990). Modern technologies for studying the requirements
and functions of plant mineral nutrients. Measurement Techniques in Plant
Science, Edited by Y. Hashimoto et al. Academic Press, San Diego. pp
319-342. #
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