Okay, so there's all this talk about pings liking lime in their medium, making it alkaline for them.... If this is the case, why can't I water them with regular tap water? If they're so indiferent to the medium and actually can tolerate alkaline conditions, is it possible to forgo using RO and just water them from the tap?
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Clue
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I was wondering the same thing about Drosophyllum and them liking alkaline conditions.
For those that don't know, our water in this part Bay Area is horribly hard and filled with calcium and other minerals.
Clue
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What does letting it sit do? Remove chlorine?
Jefforever
A yellow M&M
I believe that anyone interested in cultivating this fascinating species should read this report:
Below are two quotes taken from this report:
"Soil on the south coast of Spain is sandy or loamy, slightly acid to neutral, lime-less and poor in nutrients. The geologic underlay consists of sandstone. Drosophyllum also grows directly in sandstone crevices."
"During our observations in its native habitat, we have never seen Drosophyllum growing in alkaline soils, despite several intensive searches at suitable places."
It appears that the belief that Drosophyllum, in situ, grows in alkaline (high pH) soils, is mistaken.
I think you're confusing 2 different variables (pH and ppm.)
It's true that some CPs can thrive in slightly alkaline soils. The pH of a water source is a measure of hydrogen ions, and has little to do with other dissolved material in the water. If your tap water happens to be slightly alkaline and otherwise high-quality, knock yourself out. Most water sources are actually slightly alkaline, and run around a pH of 8.0 to prevent corrosion of old pipes. This will be canceled out if there is peat moss in your soil mix, as it is acidic.
However, the real killer in tap water in most cases is the total dissolved solids. This is what you read on a TDS meter (in ppm) and can be anything from salt, to calcium, to whatever else. You could have completely neutral pH water and have 35000 ppm salt (this is also called sea water) that would immediately destroy any CP. In my experience, pings are pretty sensitive to high ppm, so you should exercise caution before making a decision based only on pH.
How do you change the H+ or OH- concentration without adding something other than H+ or OH-?
Acids add H+ ions to the solution, such as CO2, acetic acid, citric acid, etc. Adding a base, such as calcium carbonate, potassium hydroxide, etc. will increase the availability of OH- ions. I don't know of an easy way to just add the H+ or OH- ions. I know that you can use electrochemically charged plastic resins (called: ion exchange resins) to tip the balance. Resins charged with H+ will adsorb anions and replace them with H+ ions, and likewise with OH- charged resins, they will release their OH- ions in exchange for other cations.
I wish I knew more chemistry, its been a while since I studied buffers.
it is possible to use top water , but attention to the chlore .
for the drosophyllum see here their condition 'in situ'
Pnguicula & Cie
for the drosophyllum see here their condition 'in situ'
Pnguicula & Cie
Without running thorough test for what is dissolved in your tap water, everytime you want to use it, the only way of telling if your plants are OK with it is to risk it and give it a try. If you can't take the risk then stick to RO water. pH and TDS vary from area to area and even season to season, so you can't really get a conclusive answer. One persons experience can be different from anothers due to locale. Personally, I think that there are a lot of myths with regards to how intolerant CPs are, many of them based on how they grow in their natural environment. For instance, it is often assumed that because they naturally grow in a typically nutrient poor environment then any soil nutrients will kill them but this is not necessarily true. However, enevitably there will be a limit to their tolerance, as with all plants.
Can you give an example? Let us say we start off with tap water made alkaline by dissolved calcium carbonate and we want to decrease the pH to 7 and get rid of the mineral ions.
I'll try to give a general (hopefully correct) explanation...
Basic buffers, like calcium carbonate work by keeping hydronium ions (H+ or H3O+) "occupied," which would normally contribute to lowering the pH of the solution.
For example, normally, if you added hydrochloric acid (HCl) to a glass of water, this reaction would occur:
HCl+H20--->Cl- (aqueous) + H30+
Notice that since the H30+ is "free" on the right side of the equation, it contributes to lowering the pH (a measure of the concentration of hydronium ions in solution).
First, adding calcium carbonate (CaCO3) to water, it will lower the pH slightly, but it won't contribute OH- ions, so its effects won't be as dramatic on raising the pH compared to if you poured the same mass of of sodium hydroxide into solution (see the paragraph below this for an example).
So if you add HCl to CaCO3, the main reaction would be the following (ignoring other side-reactions that would take place in solution, for simplicity):
2HCl+ CaCO3(s)---> CaCl2(aqueous) + H20 +C02
HCO3- and H2CO3 could also be formed in solution, but in all of these cases, the hydronium ions would be kept "occupied" more often and would not contribute to raising the pH as much.
Just to give another perspective- modifying an idea from wikipedia:
Water can dissolve up to 15.9 grams/Liter of CaCO3 and still maintain a pH of 7. Converting this to molarity, gets you a 0.16 mole/L solution of CaCO3 that will still have a pH of 7.
Compare that to sodium hydroxide, which with the same molarity of 0.16 mole/L (6.4 grams/L) would equate to a pH of 13.2. So calcium carbonate is able to counteract the addition of acids to solution without significantly raising the pH, themselves.
Please, NaN, or anyone- correct me if I'm wrong on anything!!!! I took chemistry 2 and a half years ago, so it's quite a bit fuzzier than it used to be.
And here's a helpful resource I used back in the day to try and understand buffers: http://www.chemguide.co.uk/physical/acidbaseeqia/buffers.html
Basic buffers, like calcium carbonate work by keeping hydronium ions (H+ or H3O+) "occupied," which would normally contribute to lowering the pH of the solution.
For example, normally, if you added hydrochloric acid (HCl) to a glass of water, this reaction would occur:
HCl+H20--->Cl- (aqueous) + H30+
Notice that since the H30+ is "free" on the right side of the equation, it contributes to lowering the pH (a measure of the concentration of hydronium ions in solution).
First, adding calcium carbonate (CaCO3) to water, it will lower the pH slightly, but it won't contribute OH- ions, so its effects won't be as dramatic on raising the pH compared to if you poured the same mass of of sodium hydroxide into solution (see the paragraph below this for an example).
So if you add HCl to CaCO3, the main reaction would be the following (ignoring other side-reactions that would take place in solution, for simplicity):
2HCl+ CaCO3(s)---> CaCl2(aqueous) + H20 +C02
HCO3- and H2CO3 could also be formed in solution, but in all of these cases, the hydronium ions would be kept "occupied" more often and would not contribute to raising the pH as much.
Just to give another perspective- modifying an idea from wikipedia:
Water can dissolve up to 15.9 grams/Liter of CaCO3 and still maintain a pH of 7. Converting this to molarity, gets you a 0.16 mole/L solution of CaCO3 that will still have a pH of 7.
Compare that to sodium hydroxide, which with the same molarity of 0.16 mole/L (6.4 grams/L) would equate to a pH of 13.2. So calcium carbonate is able to counteract the addition of acids to solution without significantly raising the pH, themselves.
Please, NaN, or anyone- correct me if I'm wrong on anything!!!! I took chemistry 2 and a half years ago, so it's quite a bit fuzzier than it used to be.
And here's a helpful resource I used back in the day to try and understand buffers: http://www.chemguide.co.uk/physical/acidbaseeqia/buffers.html
