[az2008]

I dissolve eggshells in white (distilled) vinegar for 3 weeks (it won't dissolve any more). I believe this liquid is called calcium acetate.

I strain out the undissolved material and let the liquid dehydrate, resulting is a crusty cake which I break up into a powder. I use it as a soil/fertilizer amendment which I believe supplies calcium faster than typical calcium carbonate would (oyster shells, dolomite lime). For example: it dissolves almost immediately in water. So, I assume it's a purer, more readily available form of calcium because it's already broken down(?)

My questions:

1. What is the dehydrated form called. Is it near-pure[1] calcium? Or, is it still calcium acetate?

2. If I dissolve 1mg in water, what percentage of the resulting PPM increase is calcium?

[1] I know eggshells contain trace minerals like phosphorous. I'm not trying to be that exact when I refer to purity or PPMs.

[az2008]

I did a little googling and found medical-grade tablets. Their list of active ingredients imply that calcium acetate contains 1/4 elemental calcium. (E.g., "1000mg calcium acetate, equivalent to 250mg elemental calcium").

Based upon that, it sounds like the dehydrated material I make is still calcium acetate. And, I can estimate that 25% of the PPMs dissolved in water are calcium(?).

I don't know anything about chemistry. What I'm trying to get at is

1. My Ca:Mg ratio should be between 2:1 and 4:1.
2. For Mg I use epsom salt which is MgSO4*7H2O.
3. I have read epsom salt contains about 10% Mg by weight.

So, it sounds like if I use 1 part of my dehydrated calcium acetate and 1 part epsom salt, that would be roughly 2:1 ratio of elemental calcium to magnesium?

What confuses me is how weight differs from ppms. I'm thinking more in terms of ppms, not weight.

[electronpusher96]

1)PPM is weird, it sounds like it means "number of solute molecules per 1,000,000 solvent molecules", but I don't believe that it does.  Whenever I've had to use PPM, I was told it was based on mass, so: PPM = mg solute/kg solvent = mg solute/1,000g solvent = mg solute/1,000,000mg solvent. 

2) Chemists use mols for this type of thing.  A mol of any given molecule is defined as 6.022 X 10^23 molecules of said substance(or if they're individual atoms: 6.022 x 10^23 atoms of the substance).  So, say I have a reaction: A + B  C, if I reacted 1 mol of A with 1 mol of B, I'd have 1 mol of C, however the mass of C would be equal to the mass of A + mass of B.

Once this is realized, determining elemental composition of molecules is very easy.  Look at the periodic table, the molar mass is listed on each element.

So: Let's determine how much calcium is in Calcium Hydroxide, Ca(OH)₂:

Molar Mass Ca = 40g/mol
Molar Mass O =16g/mol
Molar Mass H = 1g/mol

So each molecule contains 1 Calcium, 2 Oxygen and 2 Hydrogen =>
40 + 2(16 + 1) = 74g/mol

Using this, let's determine the percent of Calcium =>

(40g/mol) / (74g/mol) = .5405 => .5405 x 100 = 54.05% = % Calcium in Ca(OH)₂


3) Calcium is a metal, pure elemental calcium looks like any other silver colored metal, but is reactive with the air and must be stored in oil or something similar.  So, no, it's not pure calcium.  Calcium acetate is calcium acetate, if it's got 1 water it is "Calcium Acetate Monohydrate", if 2 waters it's "Calcium Acetate Dihydrate" and so on...


4) Solubility does not equal purity, your calcium acetate is dissolving faster simply because it's much, much more soluble than calcium carbonate(which is nearly insoluble.).  I don't think it's necessarily any better for plants, any maybe even could be problematic because the rain could wash it away more easily.  Furthermore, bioavalibility is not the same as purity or solubility; organisms absorb things at different rates for a myriad of reasons and I wouldn't make any assumptions.


I'd recommend reading up on the basics from a general chemistry book, and or khan academy(CrashCourse is also good for the basics.).  This stuff is all covered in the first semester of gen chem.

https://www.khanacademy.org/science/chemistry/atomic-structure-and-properties/introduction-to-the-atom/v/atomic-mass

https://www.khanacademy.org/science/chemistry/atomic-structure-and-properties/introduction-to-compounds/v/molecular-mass-and-molecular-weight

https://www.khanacademy.org/science/chemistry/chemical-reactions-stoichiome/balancing-chemical-equations/v/balancing-chemical-equations-introduction

[electronpusher96]

Also:  I would determine the quantity of Ca based on the mass of eggshells(assume they're 100% Calcium Carbonate for simplicity) and then forget about the hydrates.  Hydrates happen, often with varying numbers of water molecules.  So, if you have X mols of Ca and you solvate it, then evaporate the solvent, you'll still have X mols of Ca(Again, I'm assuming perfect conditions for simplicity.).  If the mass of the final product is different from the calculated value, that doesn't mean much because you could have a few different types of calcium acetate in there(I.e. anhydrous and multiple different hydrates.).

[az2008]

1)PPM is weird, it sounds like it means "number of solute molecules per 1,000,000 solvent molecules", but I don't believe that it does.  Whenever I've had to use PPM, I was told it was based on mass, so: PPM = mg solute/kg solvent = mg solute/1,000g solvent = mg solute/1,000,000mg solvent.

Thank you for replying, and especially the links. I will spend some time there trying to understand the topic better.

Your reply came at a great time because I was going to post another topic asking about PPMs.

I know I can mix 1g of ordinary table salt in 1L of pure (distilled) water, and produce about 1000ppm (accounting for how table salt may not be as pure as what would really be used to make calibration fluids. I calibrate my TDS meter this way. It's close enough for what I do.).

But, something I don't understand is when something like epsom salt (MgSO₄·7H₂O) is dissolved in water.

Can the expected PPMs from 1g be calculated? Does the O₄ and H₂O become "parts?" (Like the Mg and S?) Or, does it absorb into the water as dissolved oxygen which is then released into the atmosphere?

Does the Mg and S "parts" produce the same PPMs as NaCl?

For example, I know Epsom salt contains 10% and 13% Mg and S (by weight). So, should I expect 1g in 1L to produce 230 PPM (compared to 1g of NaCl producing 1000 PPM)?

I know TDS/PPM is a kludgey way of expressing electrical conductivity (EC) in gardening. But, the general rule is to multiple EC by 1000 and divide by 2. (1 EC = 500 PPM). I can live with that kind of vaguery.

But, what I have trouble with is how to predict (calculate) what a particular element will produce. I.e., the exact steps. It seems like it involves knowing the percent of mass, how many atoms are in a gram(?). But, what happens to things like H and O? Do they not count?

Sorry if I don't make sense. I'm trying to understand how, if I add 1g of K₂O (a common form of potassium in gardening), what "strength" will it be. (So, a ratio of K to N, which comes from calcium nitrate, Ca(NO3)₂, could be predicted.).

Thanks in advance!

[Borek]

But, something I don't understand is when something like epsom salt (MgSO₄·7H₂O) is dissolved in water.

MgSO₄ alone is what becomes the dissolved solid. MgSO₄ makes around 49% of the MgSO₄·7H₂O (you can calculate it using molar masses of the elements), the rest - 51% - is water. MgSO₄ dissociates into Mg²⁺ and SO₄^2- - the latter is so called polyatomic anion, it is quite stable and doesn't get further split in the solution. Water molecules on the other hand become just part of the solution and they can't be distinguished from other water molecules used to dissolve the compound. So if you add 1 g of the Epsom salt to 1 L of water expected TDS is around 0.49 g/(1000 g + 0.51 g) or 490 ppm (amount of water I added in the denominator is so low it can be ignored here, but that's not always the case).

[az2008]

MgSO₄ alone is what becomes the dissolved solid. MgSO₄ makes around 49% of the MgSO₄·7H₂O (you can calculate it using molar masses of the elements), the rest - 51% - is water.

Thank you for that explanation! I've been playing at convertunits[dot]com/molarmass and I think I see a pattern I can follow.

For example, applying it to calcium acetate Ca(CH₃COO)₂*H₂O (which is 176.18 g/mol):

• The dissolve'able part (to the left of "*") is 56.26% (158.17 g/mol) of its dry weight.
• 1g dissolved in 1L should produce 563ppm
• Of those PPM, 40.43% are Ca (228ppm)

Do I have it right? I hope I understand this now. It's been a question in the back of my mind for over a year. I feel like I can explain it to other people now.

More questions:

1. What about something like potassium oxide (K₂O) which doesn't appear to have the part to the right (*Xx_nXx_n)? If I understand this correctly, it is like NaCl (table salt) and 1g should produce 1000ppm? And, 83.02% of those PPMs will be K?

2. Why do some things like epsom salt, gypsum and calcium acetate have that discardable part to the right (*Xx_nXx_n), but other things like K₂O or calcium nitrate (Ca(NO₃)₂) don't?

Thanks!

[Borek]

For example, applying it to calcium acetate Ca(CH₃COO)₂*H₂O (which is 176.18 g/mol):

176.18 g/mol it is.

The dissolve'able part (to the left of "*") is 56.26% (158.17 g/mol) of its dry weight.

158.17 g/mol is a correct number, but it is not 56.26% of the 176.18, check your math. It should be

[tex]\frac {158.17}{176.18} \times 100%%[/tex]

Of those PPM, 40.43% are Ca (228ppm)

Again, this percentage is not correct.

1. What about something like potassium oxide (K₂O) which doesn't appear to have the part to the right (*Xx_nXx_n)? If I understand this correctly, it is like NaCl (table salt) and 1g should produce 1000ppm? And, 83.02% of those PPMs will be K?

This is tricky. 1 g of K₂O would not produce 1000 ppm. To be honest, I am not sure what is the correct number. K₂O reacts with water producing KOH, and I would calculate mass of KOH produced to use it to calculate TDS. We never use K₂O directly, but for every potassium containing compound we can easily calculate equivalent amount of K₂O, then we can use it to compare how much potassium is present in different substances.

Why do some things like epsom salt, gypsum and calcium acetate have that discardable part to the right (*Xx_nXx_n), but other things like K₂O or calcium nitrate (Ca(NO₃)₂) don't?

It reflects their composition. Some salts are hydrated - that is, they contain water molecules in their crystal structure (and the water is present there is a more or less constant amount). Some compounds can be easily written as a mixture of oxides, for example sodium metasilicate Na₂SiO₃ can be though of as combination of two oxides - Na₂O·SiO₂. In general these formulas are a mix of systematic approach and non-systematic historical conventions, they can be confusing.

[az2008]

1. What about something like potassium oxide (K₂O) which doesn't appear to have the part to the right (*Xx_nXx_n)? If I understand this correctly, it is like NaCl (table salt) and 1g should produce 1000ppm? And, 83.02% of those PPMs will be K?

This is tricky. 1 g of K₂O would not produce 1000 ppm. To be honest, I am not sure what is the correct number. K₂O reacts with water producing KOH, and I would calculate mass of KOH produced to use it to calculate TDS. We never use K₂O directly, but for every potassium containing compound we can easily calculate equivalent amount of K₂O, then we can use it to compare how much potassium is present in different substances.

Thanks! I understand this topic much better. One thing I realize I was stuck on is where the "O" and "H" goes. I thought the "parts" measured were the K, and the OH's were dissolved in water.

Now I believe they remain bound to the K and contribute as ions(?) to the electrical conductivity (EC), which is roughly interpretable as "parts" (per million).

I wasn't distinguishing between the reduced K available to the plant versus the PPMs which might be in solution. I thought everything in the water was elemental. That gets me as close as I needed to be, anyway.