Bleached leaves. Too much light or nutrition problem?

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Rick said:
But a plant in a pot is always going to loose the arms race to an active and growing bacteria population. As long as you keep shoveling in ammonia, and adding alkalinity the bacteria multiply at rates plants (or plant growers) can only dream about. Nitrifiers operate 24/7, and respond over temperature ranges better than the plants. Plant activity is focused on daylight activities
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As I understand it, the decomposing bacteria are always present when there is a carbon source in the mix. They are always there and nothing can be done to supress them. As they consume the carbon they use N both ammonium and nitrate. Therefore you must supply enough N for the both the bacteria and the plant.
If you are growing in say peat/perlite, the requirement of N for decomposers will be rather small so the N usage would be roughly equal for bacteria and plant. Now try to grow the same plant in fresh bark and you will need to supply 5 times the N just for the bacteria. The plant of course requires the same amount but if you give less than the 5 times N, the plant will starve as the microbes satisfy themselves first.
The only way of really knowing how much N a particular mix uses is to perform a Nitrogen draw-down test which involves testing ammonium and nitrate leachates after a series of incubation times. But you can pretty well bet that composted bark or chc mixes will use 2-3 times the N that an inert mix does and fresh bark would probably need 6 times the N for the same plant growth not taking CEC into account of course.
The nitrifiers are totally different species and don't really compete with the plant for overall N.
The option of running higher nitrate/ammonia ratios and operating at lower alkalinity starves the bacteria population
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INCREASES the bacteria poulation unless Ca becomes almost non existant in which case the plant will not grow well.
Note that the decomosers operate at pH values from 5 to 9 and nitrifiers above 6.
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Rick said:
http://www.google.com/patents/EP2229054A1?cl=en

Crazy

Check out this patent for use of amino acids as plant/mycorrhizae fertilizer!!


All these amino acids are in kelp, so don't know how you can patent this.
Click to expand...

I think this is specific patent for the use of exactly content of pure aminoa acids.
The content in kelp is extremly viable depending on which lep you use and where it has grwon year .....
It is like a good wine. Here they use mostly extracted amino acids, it is the same that soem companys use amino acid for chelating trace elements for better bio availability. You can make a patent on nealy everything if you can pay the cost, show that you are the first one who did it and register it, its not only the ingredient it is the idea or maybe the use you can make a patent on.
 
I know this is unrelated, but these last few posts about bacteria reminded me of something interesting:

Think oils spills and "oil eating super-bugs". The technology works, but the "bug" population fades away quickly, so they need to be replenished. Having worked in the petrochemical and refining industry for 30 years, I learned that in the Exxon Valdiz spill, Elf Aquitaine (my company) utilized an updated method that was far more effective: spray an emulsion of olive oil and high-nitrogen fertilizer on the spill. The N sparks an explosion in the population of indigenous bacteria, who start by consuming the light, and easy to devour olive oil. As the population expands and the olive oil supply dwindles, they go after the heavier oil! This will work on your driveway too, but don't slip on the olive oil.

So do we know of K-eating bacteria we can seed the potting media with to clean it up?
 
Rick said:
Only if your mix is sterile.

But the nitrifying bacteria are very good at colonizing mixes and they "eat ammonia" and drop pH as they do so.

Basic waste water, and aquarium technology.
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Guess its ammonia they eat ;) you have a point, provided your bacterial(nitrosomonas) population is high enough to oxidize enough of the ammonia. Guess the net reaction going from NH3 to NO3- is one H+ in excess? But if your pH is 5.5 or so then your nitrosomonas would not be very active or? :confused: And of course enough Nitrobacter has to be present as well since we are not interested in nitrite that is the product of the nitrosomonas. Nitrobacter enjoy pH levels of slightly above 7 so perhaps they do not like the acid bark after all?
Btw. I do water frequently, but drawing a comparison to aquariums is perhaps to exaggerate a bit :poke:
Joke aside;
Another thing is the fact that those of you that uses well water should pay a bit attention to the alkalinity, or carbonate hardness as its called over here. It is capable of altering the pH balance of your mix more that the possible urea will. At my place I have water from a well that gives an alkalinity of 250ppm. Fresh the pH is around 7 but that is due to dissolved CO2. after a while it attains values of 8-8.5. If I would use that for irrigation, it would totally dominate the pH given my fertiliser addition is at only some 200ppm. I try to control the pH of my irrigation water to be slightly below 6. That is easy by using rainwater, but with the well water I would probably run into trouble; as I did with my fresh-water aquarium.:eek: Its ok now:)
 
Bjorn said:
Another thing is the fact that those of you that uses well water should pay a bit attention to the alkalinity, or carbonate hardness as its called over here. It is capable of altering the pH balance of your mix more that the possible urea will. At my place I have water from a well that gives an alkalinity of 250ppm. Fresh the pH is around 7 but that is due to dissolved CO2. after a while it attains values of 8-8.5. If I would use that for irrigation, it would totally dominate the pH given my fertiliser addition is at only some 200ppm. I try to control the pH of my irrigation water to be slightly below 6. That is easy by using rainwater, but with the well water I would probably run into trouble; as I did with my fresh-water aquarium.:eek: Its ok now:)
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This is actually what I have been doing. I have hard water and I always try to bring down the ph to 6 using phosphoric acid. I took a sample of the drip water after three days and guess what, its 7.5 - 8.0. I think this is due to the calcium bicarbonates in the water. If this is the case then I take that the acidity is not a problem.
 
Ray said:
So do we know of K-eating bacteria we can seed the potting media with to clean it up?
Click to expand...

There are not. They do not eat calcium either as Stone eluded to in a previous post.

Since I work in biological waste treatment the world would go crazy if someone came up with a bug that "consumes" the 4 major cations (Na, Ca, K, and Mg).

The bacteria (like plants and humans) need a very small amount of these materials for biomass. But that amount is insignificant compared to the amount of carbon consumed or the ammonia that is converted to nitrate by nitrifiers.

Yes nitrifying bacteria are different from general heterotrophic species, but are significant in potting media. They probably convert most of the ammonia to nitrate in a fert mix. The actually do not "consume" ammonia, but use it as a source of electrons (along with oxygen). Protons are a waste product of the reaction, and requires alkalinity OH and HCO3 (not calcium) to neutralize it (to keep from poisoning themselves out).

Nitrifying bacteria are the crux of municipal waste water treatment. These guys handle litterally tons of ammonia per day (converting to nitrate). Your basic aquarium opperates on the exact same principal to keep the ammonia produced by fish excreetion to pollute your home aquarium.

They are everywhere in damp soil conditions as long as pH is > 5.0 and O2 is present. They work in fresh and salt water. Temps almost to freezing and up to 30+C.
 
Bjorn said:
Guess its ammonia they eat ;) you have a point, provided your bacterial(nitrosomonas) population is high enough to oxidize enough of the ammonia. Guess the net reaction going from NH3 to NO3- is one H+ in excess? But if your pH is 5.5 or so then your nitrosomonas would not be very active or? :confused: And of course enough Nitrobacter has to be present as well since we are not interested in nitrite that is the product of the nitrosomonas. Nitrobacter enjoy pH levels of slightly above 7 so perhaps they do not like the acid bark after all?
Btw. I do water frequently, but drawing a comparison to aquariums is perhaps to exaggerate a bit :poke:
Joke aside;
Another thing is the fact that those of you that uses well water should pay a bit attention to the alkalinity, or carbonate hardness as its called over here. It is capable of altering the pH balance of your mix more that the possible urea will. At my place I have water from a well that gives an alkalinity of 250ppm. Fresh the pH is around 7 but that is due to dissolved CO2. after a while it attains values of 8-8.5. If I would use that for irrigation, it would totally dominate the pH given my fertiliser addition is at only some 200ppm. I try to control the pH of my irrigation water to be slightly below 6. That is easy by using rainwater, but with the well water I would probably run into trouble; as I did with my fresh-water aquarium.:eek: Its ok now:)
Click to expand...

Being in a company that specializes in biological waste water treatment, we've learned the limits of what nitrobactor and nitrosomonas are capable of. Actually they live in joint colonies proximal to each other. It is possible to kill one but leave the other alive, but since they noramally are found together the switch from NH3 to NO2 to NO3 is almost instantaneous, and for engineering purposes consider a single step. As noted earlier, every mg of NH3 converted to nitrate requres over 7mg of alkalinity (over 8 mg of bicarbonate) to neutralize the proton production. That's a fair amount of acidity produced. Also the pH at absolute 0 alkalinity in a system is a little over 4.0 su. So yes the efficiency of nitrification is reduced as pH goes down, but it is proportionate to the alkalinity availability (rather strictly pH). So you can still get nitrification at pH's below 6.0

I'm not sure if you have heard of wet/dry or trickle filters. Instead of having your bacteria colonizing submerged media in the aquarium, you grow your bacteria on a high surface media (often Leca!!) that you spray water over so it is wet but not submerged. The nitrifiers are much more efficient in this system. (and very analogous to a semi hydro orchid growing system).
 
Bjorn said:
Guess its ammonia they eat ;) you have a point, provided your bacterial(nitrosomonas) population is high enough to oxidize enough of the ammonia. Guess the net reaction going from NH3 to NO3- is one H+ in excess? But if your pH is 5.5 or so then your nitrosomonas would not be very active or? :confused: And of course enough Nitrobacter has to be present as well since we are not interested in nitrite that is the product of the nitrosomonas. Nitrobacter enjoy pH levels of slightly above 7 so perhaps they do not like the acid bark after all?
Btw. I do water frequently, but drawing a comparison to aquariums is perhaps to exaggerate a bit :poke:
Joke aside;
Another thing is the fact that those of you that uses well water should pay a bit attention to the alkalinity, or carbonate hardness as its called over here. It is capable of altering the pH balance of your mix more that the possible urea will. At my place I have water from a well that gives an alkalinity of 250ppm. Fresh the pH is around 7 but that is due to dissolved CO2. after a while it attains values of 8-8.5. If I would use that for irrigation, it would totally dominate the pH given my fertiliser addition is at only some 200ppm. I try to control the pH of my irrigation water to be slightly below 6. That is easy by using rainwater, but with the well water I would probably run into trouble; as I did with my fresh-water aquarium.:eek: Its ok now:)
Click to expand...

http://aem.asm.org/content/70/11/6481.short

Bjorn This is a new one on me. Efficient nitrification at pH 4.3!!! They are using a fixed film reactor (similar to WET/Dry, trickle system I mentioned earlier).
 
Rick said:
Being in a company that specializes in biological waste water treatment, we've learned the limits of what nitrobactor and nitrosomonas are capable of. Actually they live in joint colonies proximal to each other. It is possible to kill one but leave the other alive, but since they noramally are found together the switch from NH3 to NO2 to NO3 is almost instantaneous, and for engineering purposes consider a single step. As noted earlier, every mg of NH3 converted to nitrate requres over 7mg of alkalinity (over 8 mg of bicarbonate) to neutralize the proton production. That's a fair amount of acidity produced. Also the pH at absolute 0 alkalinity in a system is a little over 4.0 su. So yes the efficiency of nitrification is reduced as pH goes down, but it is proportionate to the alkalinity availability (rather strictly pH). So you can still get nitrification at pH's below 6.0
Click to expand...

Thank you for explaining these things to me; but there arel details in your calculatuons that puzzles me: Please correct me if I am wrong. I am sincerely interested in understanding these issues a bit more in-depth.:)
Although Nitrification involves exchange of quite a few electrons and protons the net reaction going all the way from NH3 to NO3- produces only one proton. That should be capable of eliminating one HCO3-. or; nitrification of one gram NH3 should be able to consume 3.6gram hydrogencarbonate. Or if you prefer: 4.8 g Calciumhydrogencarbonate. But, the latter is hardly available as such in the water I guess:poke:
This use of alkalinity is another thing I do not really get the grip on: When you refer to alkalinity, is it as hydrogencarbonate or as carbonate? If it is as calciumcarbonate, the oxidation of one gram NH3 would neutralise 5.9 gram CaCO3 and not the 4.8 gram as predicted for calciumhydrogencarbonate.
Ok I have checked up a bit and it seems as if alkalinity is normally expressed as CaCO3. So If I used my well water of 250mg/l CaCO3 this is equivalent to approximately 2.5 mmol/l CaCO3. To neutralise that it would be necessary to have 2.5mmol/l ammonia from Urea that was converted to nitrate. Since urea decomposes to 2 NH3 this should be the same as 1.25mmol/l urea in the irrigation water which is equivalent to approximately 0.075 g/l. This equals to 75 mg/l. In this case I should not exceed 75mg/l urea in order not to acidify the mix. Or put differently: If I was using my well water and a fertiliser with 20%N from urea (=43%urea in the fertiliser) then any excess of 175ppm fertiliser would be capable of acidifying my mix. If everything was converted. Using a back mix of 1:10 with RO water should imply that some of that urea fertiliser(>17.5ppm urea fertiliser) would be beneficial. Of course if the urea is not 100% converted during its passage through the compost, more is needed.:p
The pH being 4 at no alkalinity is quite theoretical, it is totally dependent on the amount of dissolved solids and gases and as such highly variable. But due to CO2, it is normally acid and extremely hard to measure accurately.;)
PS: I found this one; at least for embryos of Cattleya, ammonium seems vital.
http://harvardforest.fas.harvard.ed...publications/pdfs/Raghavan_AmJBotany_1964.pdf
And this one explains it all:
http://forums2.gardenweb.com/forums/load/orchids/msg032222191547.html
 
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Bjorn said:
The pH being 4 at no alkalinity is quite theoretical, it is totally dependent on the amount of dissolved solids and gases and as such highly variable. But due to CO2, it is normally acid and extremely hard to measure accurately.;)
Click to expand...

However, the ASTM method for measuring alkalinity is a titration against a standardized acid solution, with the endpoint of pH 4.0~~ The titration and math are very easy to do.

Have you gone through this paper?

http://www.staugorchidsociety.org/PDF/IPAFertilizers.pdf
 
Bjorn said:
This use of alkalinity is another thing I do not really get the grip on: When you refer to alkalinity, is it as hydrogencarbonate or as carbonate? If it is as calciumcarbonate, the oxidation of one gram NH3 would neutralise 5.9 gram CaCO3 and not the 4.8 gram as predicted for calciumhydrogencarbonate.
Ok I have checked up a bit and it seems as if alkalinity is normally expressed as CaCO3.
Click to expand...

Bjorn Although expressed as CaCO3 equivalents, at the pH ranges and gas saturation levels we are working with, the predominat ion is bicarbonate (which I guess is what you are calling hydrogencarbonate, HCO3).

Calcium carbonate is esentially insoluble (until pH <<4.0) I can't get it to melt until pH 2.0 in a reasonable length of time. But if you want to generate measurable alkalinity fast, add some baking soda (NaHCO3) to some RO water, and the math is easy.

The conversion from alkalinity as CaCO3 to bicarbonate ion is just a factor of 1.22. So "alkalinity as CaCO3)"X1.22 will give you bicarbonate ion concentration.
 
Rick said:
However, the ASTM method for measuring alkalinity is a titration against a standardized acid solution, with the endpoint of pH 4.0~~ The titration and math are very easy to do.

Have you gone through this paper?

http://www.staugorchidsociety.org/PDF/IPAFertilizers.pdf
Click to expand...

Yes it puzzles me too. Today it seems widely accepted that nitrogen is used by the plant for production of aminoacids, and those are synthesized from ammonium, not nitrate. Quite a few investigations show that ammonium is assimilated quicker than nitrate. This is probably because it immediately goes into aminoacid synthesis. Nitrate can be stored for later use so a mix gives best results. In several tests, ammoniumnitrate is the preferred over nitrate or ammonium. In soil, Bill Argos presentation is probably correct, if all ammonium gets oxidised, but I am uncertain whether it is applicable for orchids.
Take Fig. 1 for instance. It may be read as if urea splits into ammonium and two protons, while reality is that urea yields two ammonium and two protons. This is possibly the root of quite a few misunderstandings regarding this topic.

The end point of the titration is just a matter of convention. It was chosen so to represent a value to keep to doing quick tests on that matter. Making real titration curves gets very difficult as the point of equivalence gets smeared out due to all the other salts except calcium carbonate that sums up the alkalinity. As a curiosity can be mentioned monosilicic acid that is the silicon species that is absorbed by plants.;)
 

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