CATIONIC SPECIATION IN NUTRIENT SOLUTIONS AS A FUNCTION OF pH
G. De Rijck and E. Schrevens
Faculty of Agricultural and Applied Biological Sciences, Department of Applied Plant Sciences K.U.Leuven, Willem de Croylaan 42, B-3001 Heverlee (Belgium)
ABSTRACT
The pH of a nutrient solution is a property that is inherent to its composition. Changing the pH of a nutrient solution affects its composition, elemental speciation and bioavailability. The term “speciation” indicates the distribution of elements among their various chemical and physical forms like: free ions, soluble complexes, chelates, ion pairs, solid and gaseous phases and different oxidation states.
For a standard nutrient solution elemental speciation is calculated for a pH range from 2 to 9. For each cation the formation of precipitates, ion pairs, complexes and chelates is illustrated as a function of pH. The calculation of elemental speciation in nutrient solutions is an indispensable tool in the design, the analysis and the interpretation of experiments with the mineral composition of nutrient solutions in plant nutritional research.
INTRODUCTION
According to Steiner (1961), the chemical composition of a nutrient solution is determined by the relative cation proportions, the relative anion proportions, the total
ionic concentration and the pH. The pH of an aqueous solution is determined by the initial concentration of acids and bases. So once the relative anion proportions, the relative cation proportions and the total ionic concentration are determined the pH is also determined. This means that the pH is an inherent property of the mineral composition of the nutrient solution that can not be changed independently (De Rijck and Schrevens, 1997a).
The pH of a nutrient solution can be changed by adding a strong acid (HNO3) or a strong base (KOH). In commercial practice this way of pH adjustment is used, resulting in a changed mineral composition. Another way to increase the pH of a nutrient solution is adding a mixture of KOH, Ca(OH)2 and Mg(OH)2. In this mixture the relative proportions of K+, Ca2+ and Mg2+ are the same as their proportions present in the nutrient solution.
To reduce the pH a mixture of HNO3, H3PO4 and H2SO4 can be used, in the same relative proportions as the proportions of NO3-, H2PO4- and SO42- present in the nutrient solution.
In both cases the pH is changed without changing the relative cation or the relative anion proportions of the initial solution, respectively. However, the ratio of each cation to each anion and the total ionic concentration changes.
In the speciation calculations the pH is changed in this way for the standard nutrient solution of the Research Centre for Soilless Cultures at a total ionic concentration of 32.34 mmol/L (50 meq/L) (Schrevens et al., 1988a) (Table 1).
MATERIAL AND METHODS
This research only carries out speciation calculations, using the computer program Geochem PC version 2.0 (Parker et al., 1995). The effects of pH changes on elemental speciation and on the total ionic concentration are elaborated for a pH range from 2 to 9. The effect of temperature, atmospheric pressure and composition is not taken into account. All the calculations are made
for closed systems (no CO2 exchange with the atmosphere) at a temperature of 25 °C and an atmospheric pressure of 0.1 Mpa.
Both real complexes (Fe-HEDTA), ion pairs (Mg-SO4) and acidification (NH3 -H) are considered as soluble complexes.
The standard nutrient solution of the Research Centre for Soilless is successfully used for the cropping of tomatoes, lettuce, cucumbers, peppers, celery, leek and strawberries in hydroponics (Schrevens et al. 1988a; 1988b; De Rijck et al., 1993a and 1993b).
Na and Cl present in the solution are remnants of the preparation of the iron chelate. The iron chelate Fe(HEDTA) is prepared by dissolving equimolecular amounts of Fe(Cl)3 and Na3(HEDTA).
The pH of the considered standard solution is 4.2. In the speciation calculations the pH is changed by adding a mixture of KOH, Ca(OH)2 and Mg(OH)2 or a mixture of HNO3, H3PO4 and H2SO4 in the same proportions as their proportions in the initial nutrient solution. So increasing the pH above 4.2 also increases the K+, Ca2+ and Mg2+ concentration, while reducing the pH below 4.2 increases the NO3-, H2PO4- and SO42- concentration. Changing the pH from 4.2 to 9 or from 4.2 to 2 increases the total macro-cation concentration from 25 to 29.9 meq/L and the macro-anion concentration from 25 to 44 meq/L respectively.
For each cation the effect of pH on elemental speciation is represented by its change in free concentration, due to the formation of precipitates, soluble complexes or ion pairs.
RESULTS AND DISCUSSION
The speciation behaviour of the standard nutrient solution is represented in Figures 1 to 8
.
The left vertical axis represents the total and the free concentration of the cation considered. The right vertical axis indicates the concentration of the different complexes or precipitates that are formed.Potassium
Increasing the pH from 4.2 to 9, increases the total K+ concentration from 11 to 13.1 mmol/L (Figure 1). In the pH range 2 to 9, potassium is almost completely present as a free ion. According to the pH, between 0.09 and 0.12 mmol/L of K+ forms a soluble complex with SO42- (Kf = 100.62). An even smaller amount of K+ is bound to Cl-.
Calcium
Increasing pH from 4.2 to 9, increases the total calcium concentration from 5.5 to 6.6 mmol/L (Figure 2). As the pH increases, also the fraction of phosphate present as PO43- and HPO42- increases (De Rijck and Schrevens, 1997b). Both effects increase the amount of solid beta-calciumphosphate (Ca3(PO4)2, Ksp = 10 -28.9) and reduce the free calcium concentration from 5 to 3 mmol/L. From pH 5 on Ca3(PO4)2 precipitates. At a pH above 7, 1.1 mmol/L of Ca3(PO4)2 is precipitated. The precipitated calcium is not available for uptake by plants. As a function of the pH, between 0.35 and 0.5 mmol/L of calcium forms a soluble complex with SO42- (Kf = 101.73). 0 to 0.1 mmol/L of Ca2+ is bound to HEDTA and H2PO4-. A negligible part of calcium is bound to chloride and nitrate.
Magnesium
Increasing pH from 4.2 to 9 increases the total magnesium concentration from 1.5 to 1.8 mmol/L (Figure 3). Magnesium is almost completely present as a free ion. As a function of the pH only 0.12 to 0.16 mmol/L of magnesium is bound to SO42- (Kf = 101.63). Magnesium forms soluble complexes with HPO4-, HEDTA, OH- and Cl-.
Sodium
The 0.4 mmol/L of sodium is in the nutrient solution almost completely present as a free ion (Figure 4). Only about 2 µmol/L of sodium is bound to SO42-. An even smaller concentration of sodium is bound to Cl-.
Iron
Due to the stability of the iron chelate and the precipitation of iron almost no iron is present as a free ion (Figure 5). For the pH range 2 to 7 most of the iron is bound to the chelate, indicating the suitability of HEDTA to chelate iron in this pH range. At pH 2 and 3, most of the iron is bound to HEDTA. Increasing pH above 2, gradually decreases the amount of iron bound to HEDTA in favour of the amount of iron bound to (HEDTA)OH. Changing pH from 5 to 8 decreases the amount of iron bound to (HEDTA)OH from 0.09 to 0 mmol/L.
At pH 2 and 3 small amounts of iron form a soluble complex with H2PO4-, HPO42-, NO3- and B(OH)4-.
From pH 2 to 6, the amount of iron precipitated as FePO4.2H2O (strengite) increases from 0.01 to 0.05 mmol/L. Increasing pH above 6, the amount of iron precipitated as Fe(OH)3 gradually increases. As the pH exceeds 8, iron is completely precipitated as Fe(OH)3.
Manganese
For a pH lower than 4 only 0.005 mmol/L of manganese forms a soluble complex with SO42- or Cl-, while the largest part is present as a free ion (Figure 6). Increasing the pH above 4, manganese present as a free ion gradually binds HEDTA. At pH 8 to 9 manganese is almost completely bound to HEDTA.
Copper
For the complete pH range copper is almost completely bound to the chelate, resulting in a very low free Cu2+ concentration (Figure 7). At pH 2 most of the copper is bound to H(HEDTA). Increasing pH reduces this amount if favour of HEDTA. Small concentrations of copper form soluble complexes with H2PO4-, PO43- and NO3-.
Zinc
Increasing pH from 2 to 4, reduces the 1.5 µmol/L zinc present as a free ion (Figure 8), while the amount of zinc bound to HEDTA increases to 2 µmol/L. In the pH range of nutrient solutions zinc is almost completely bound to HEDTA. At a low pH, small concentrations of zinc are bound to SO42-, H2PO4- and NO3-.
CONCLUSION
The cations K+, Ca2+, Mg2+ and Na+ are mainly present as a free ion in the nutrient solution. At a higher pH range some calcium precipitates as Ca3(PO4)2. The other cations (Fe3+, Mn2+, Cu2+ and Zn2+) are mainly bound to the chelate in the pH range of nutrient solutions.
The Fe3+ that dissociates from the chelate, precipitates as FePO4.2H2O (strengite) or as Fe(OH)3 at a pH above 8.
The pH of the nutrient solution is a determining factor in the amount of iron accumulated by plants (Jeffreys et al., 1961). The iron status of a plant influences the form in which iron is absorbed. Iron deficient plants absorb iron more readily than chelating agent (Hill-Cottingham and Lloyd-Jones, 1965). This results in more free chelate, competing with plants for the uptake of metals other than iron from the nutrient solution (Halvorson, 1971).
As illustrated by the results elemental bioavailability in nutrient solutions is seriously affected by speciation reactions. The equilibrium of dissociation,
complexation and precipitation reactions in nutrient solutions results for each ion in a complex balance between all its forms present.
Considering also the following phenomenons results in an even more complex equilibrium. Plants are selective in their absorption of ions (Epstein, 1972; Morard et al., 1990; Morard and Benavides, 1990 and Morard et al., 1993).
Plant roots release H+ if relatively more cations are taken up than anions and HCO3- and OH- if relatively more anions are taken up than cations (Willumsen, 1984), resulting in pH changes, reflected in elemental speciation.
Across root membranes and transversely along the root, the appearance of proton driven electrical currents changes the pH at the root surface (Marschner et al., 1987), resulting in a change in elemental speciation.
Besides the inorganic ions present in the nutrient solutions plants can secrete up to 12 to 18 % of the total carbon fixed from the atmosphere (Barber et al., 1976). In solution culture however, roots produce smaller quantities of exudates than in soil (Bowen and Rovira, 1976). These exudates will be dispersed through the bulk of a well stirred nutrient solution. All these organic substances can associate with the inorganic ions present, forming complexes, ion pairs or precipitates.
So it is clear that in function of the pH of nutrient solutions speciation reactions seriously affect elemental bioavailability. This should be taken into account in doing experiments with the mineral composition of nutrient solutions, especially in doing so called “univariate” pH experiments. Speciation calculations supply lots of useful information to interpret or even explain plant response to a certain treatment and can be considered as an indispensable tool in the design, the analysis and the interpretation of experiments with nutrient solutions in plant nutritional research
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TABLE 1. Standard nutrient solution for the Research Centre for Soilless Cultures
Nutrient mmol/L Salt mg/L
K+ 11 KH 2PO4 1298.825 Ca2+ 5.5 KNO 3 306.203 Mg2+ 1.5 K2SO4 631.938 Na+ 0.38151 Ca(NO3)2.4H2O 217.838 NH4+ 0.00054 MgSO4.7H2O 369.720 Fe3+ 0.12717 (NH4)6Mo7O24 0.110 Mn2+ 0.04375 ZnSO4.7H2O 0.539 Cu2+ 0.00062 CuSO4.5H2O 0.156 Zn2+ 0.00188 MnSO4.H2O 7.394 NO3- 17.25 H 3BO3 1.932 H2PO4- 2.25 Fe(HEDTA)4.5% 0.120 ml/L SO42- 2.75 Cl- 0.38 B(OH)4- 0.03126 MoO42- 0.0006 HEDTA 0.12717
K+ (mmol/l) Complex (mmol/l) pH Complex K+ Free Total KCl- KSO4 -0 2 4 6 8 10 12 14 2 3 4 5 6 7 8 9 0.00 0.03 0.06 0.09 0.12 0.15
Ca2+ (mmol/l) Complex (mmol/l)
pH Complex
Ca2+ Free Total
CaH2PO4+ Ca(HEDTA)- Ca3(PO4)2s
CaSO4 0 1 2 3 4 5 6 7 2 3 4 5 6 7 8 9 0.0 0.2 0.4 0.6 0.8 1.0 1.2
Mg2+ (mmol/l) Complex (mmol/l)
pH Complex
Mg2+ Free Total
Mg(HEDTA)- MgHPO4 MgOH+
MgSO4 0.0 0.5 1.0 1.5 2.0 2 3 4 5 6 7 8 9 0.00 0.05 0.10 0.15 0.20
Na+ (mmol/l) Complex (mmol/l) pH Complex Na+ Free Total NaCl NaSO4 -0.0 0.1 0.2 0.3 0.4 2 3 4 5 6 7 8 9 0.0000 0.0005 0.0010 0.0015 0.0020 0.0025
FeOH(HEDTA)
Fe3+ (mmol/l) Complex (mmol/l)
pH Complex Fe3+ Free Total 0.00 0.02 0.04 0.06 0.08 0.10 0.12 0.14 2 3 4 5 6 7 8 9 0.00 0.02 0.04 0.06 0.08 0.10 0.12 0.14
FeH2PO42+ FeHPO4+ FePO4s
Fe(HEDTA) FeOHs
Mn2+ (mmol/l) Complex (mmol/l) pH Complex Mn2+ Free Total MnCl+ Mn(HEDTA)- MnSO4 2-0.00 0.01 0.02 0.03 0.04 0.05 2 3 4 5 6 7 8 9 0.00 0.01 0.02 0.03 0.04 0.05
CuH(HEDTA)
Cu2+ (mmol/l) Complex (mmol/l)
pH Complex
Cu2+ Free Total
CuH2PO4+ CuPO4- Cu(HEDTA)
-0.0000 0.0001 0.0002 0.0003 0.0004 0.0005 0.0006 0.0007 2 3 4 5 6 7 8 9 0.0000 0.0001 0.0002 0.0003 0.0004 0.0005 0.0006 0.0007
Zn2+ (mmol/l) Complex (mmol/l) pH Complex Zn2+ Free Total ZnH2PO4+ Zn(HEDTA)- ZnNO3+ ZnSO4 0.0000 0.0005 0.0010 0.0015 0.0020 2 3 4 5 6 7 8 9 0.0000 0.0005 0.0010 0.0015 0.0020