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Water hardness
IntroductionWater hardness provides a very simple indication of whether our water is ‘hard’ – richly mineralized, or ‘soft’ – sparsely mineralized.
Hard water usually forms when rain percolates through base rich rocks and a small proportion of that rock is dissolved and goes into solution as ions. If rain falls onto a catchment that is entirely without base rich rocks the water remains “soft” and sparsely mineralized.
Water hardness does not measure other dissolved solids or the bicarbonate content.
This is a water hardness map for UK tap water.
Permanent Hardness
“Permanent” or “General” water hardness refers to the multivalent ion content of the water. In practice these ions are nearly always calcium (Ca) and magnesium (Mg). This is because calcium and magnesium are usually the most abundant cations dissolved in water.
Nothing else contributes to permanent hardness, and when just “water hardness” is mentioned, we are referring to the “permanent hardness”.
Ca and Mg are just two of the six alkaline-earth metals found in group 2 of the periodic table. The other four, beryllium (Be), strontium (Sr), barium (Ba), and radium (Ra), do not make a significant contribution. The same applies for all other metals with multivalent ions (such as iron (Fe), manganese (Mn), aluminum (Al) etc).
There are exceptions, though. The most significant among them is perhaps the water we use for keeping Rift Lake cichlids. This has little calcium, slightly more magnesium and is rich in sodium (Na), which is alkaline but does not contribute to the permanent hardness (it is a monovalent ion (Na+)).
Units
Traditionally we measure water hardness in German degrees (°dGH).
Water processing plants often express hardness in ppm (mg/L) or mmol/L (milimol per litre). Using molar or mg/L values is to be preferred over using derived units like °dGH, but we can convert from one unit to another.
1 mmol/L = 5.61 °dGH
1 °dGH = 0.18 mmol/L
Ca: 1 mg/L = 0.14 °dGH = 0.03 mmol/L = 24.95 µmol/L (µM)
Mg: 1 mg/L = 0.23 °dGH = 0.04 mmol/L = 41.14 µmol/L (µM)
(Identical amounts by weight of Mg and Ca do not contribute equally to water hardness because their relative atomic masses (RAM) differ: RAM Ca = 40.1, Mg = 24.3.)
Alkalinity
Alkalinity is often – inaccurately – called carbonate (or temporary) hardness. In fact, alkalinity or acid neutralizing capacity (ANC) refers to the amount of strong acid needed to change pH from current value to a different (lower) value.ANC4.5 is used regularly in water processing plants’ reports, and it denotes ANC to reach pH value of 4.5 – which is a point where no bicarbonate ions (HCO3-) remain in water. ANC is not often used in the UK, but is a very useful concept.
The “drop tests”, commonly used by hobbyists, measures ANC4.5, and the result is expressed in German degrees (°dKH).
This may be a bit confusing in two ways:
- The units are similar, as well as the obsolete term ‘carbonate hardness’, but alkalinity is not necessarily related to water hardness. A sample may contain lots of dissolved calcium and magnesium and still feature (near-)zero alkalinity. And vice versa, it is easily possible to make a solution with high alkalinity without any calcium or magnesium. - It is only because in practice, higher content of Mg and Ca is often accompanied by higher content of bicarbonates, and these numbers are often provided side by side.
- Alkalinity (ANC4.5) does not consist solely of bicarbonates, but also phosphates, silicates, and various organic compounds.
Like water hardness, alkalinity too is best expressed in mmol/L or mg/L.
1 mmol/L = 2.80387 °dKH
1 °dKH = 0.356650 mmol/L
HCO3- : 1 mg/L = 0.016389 mmol/L = 0.0460 °dKH
Carbonate system
Carbonates are of paramount importance in natural waters as well as in our tanks. When carbon dioxide (CO2) dissolves in water it also reacts with water, and following compounds appear: loosely hydrated CO2·H2O (commonly abbreviated as CO2), H+ + HCO3- (bicarbonate), and 2H+ + CO32- (carbonate). The distribution of each compound depends on pH and, to a lesser extent, on temperature.
We can see that in neutral-to-alkaline water, plants are forced to use bicarbonates instead of much-preferred carbon dioxide for photosynthesis. Decades ago, scientists believed that among submerged higher plants, some species are able to use bicarbonates while others are not. Recently the opinion prevails that various species differ in degree of ability to use bicarbonates, and there is no strict boundary either/or. Without exemptions, all higher plants prefer carbon dioxide if available.
If we know pH and alkalinity, we can calculate the CO2 content using following equation:
where A stands for alkalinity in German degrees (°dKH) and 𝛠(CO2) stands for CO2 content in mg/L.
Alternatively, you can use the attached chart Tillmans.xlsx.
In planted tanks, we mostly prefer slightly acidic water. There are several reasons for that. First, as we can see, in acidic water larger share of inorganic carbon takes up the form of dissolved CO2. Second, most micronutrients (B, Fe, Mn, Zn, Cu, Ni) and phosphorus are better accessible to plants’ roots in moderately acidic environment.
There is one more reason of which hobbyists are often unaware – bicarbonates can hinder assimilation (in contrast to uptake) of iron (and perhaps other transition metals) inside the plant. So, pH and bicarbonate content (alkalinity) are tightly tied together but have differing effects on plants. Various species differ significantly in their tolerance to both.
This may be of significance when injecting CO2. Addition of CO2 to water lowers its pH but does not change its alkalinity, since the reaction produces the same number of positively contributing species (H+) as negative contributing species (HCO3− and/or CO32−).
Let’s have a look at the following pic (note the alkalinity axis is logarithmic):
The blue field encompasses combinations of pH and alkalinity feasible in natural habitats (CO2 content ranging from 0.5 to 4 mg/L). If we increase CO2 content to 10 to 30 mg/L, we enter the red field. We reach pH=6.5, for example, but the alkalinity remains high (from 0.9 to 2.7 °dKH). Such a coincidence is inevitably outside the blue field – the conditions our plants experience in nature.
Is it of practical significance? That remains open to discussion, observation, and research.
Electric conductivity
Electric conductivity is indicative of the quantity of dissolved compounds. As a rule, the more dissolved compounds the higher conductivity. The SI unit of electrical conductivity is S/m (siemens per metre).1 S/m = 1000 mS/m (mili Siemens) = 10000 µS/cm (micro Siemens).
1 µS/cm = 0.1 mS/m
1 mS/m = 10 µS/cm
Because our water is a “dilute solution” conductivity values are usually shown in micro Siemens (µS) /cm.
Electric conductivity is neither identical, nor linearly related to, total dissolved solids (ppm TDS). Electric conductivity measures the concentration of electrically charged ions (cations and anions) in water. Non-charged dissolved compounds – typically many organic compounds – do not change conductivity.
Total dissolved solids meters (TDS meters) measure electric conductivity and convert it by a constant k: TDS [ppm] = k · EC [µS/cm], where k is either set by the manufacturer or can be adjusted by the user. The conversion factor k usually varies between 0.5 and 0.8. There is no generally accepted value for conversion coefficient. Thus, the deviations among TDS meters are quite significant, and aquarium hobbyists should be discouraged from measuring total dissolved solids.
Measuring TDS is problematic, even in laboratory conditions. When boiling the sample (105 °C), some compounds are lost by evaporation (NH3 etc.), some change their nature (bicarbonates turn to carbonates and even to hydroxides), and some compounds keep water of crystallization (MgSO4·7H2O, etc.). Even more complications occur in the case of organic compounds.
<Many thanks to @dw1305 , @Guest , @Zeus., @Geoffrey Rea, @Hanuman , @LondonDragon , @X3NiTH for their valuable input and criticism.>
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