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O, that egg again!

We’ve arrived at Part 6 if this extraordinary saga of how calcium arrives and behaves in the soil (if I’ve occasionally typed “soul” instead of “soil”,  is it really a typo?). Click to catch up on part 1, part 2, part 3 and part 4 and part 5.

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6. Soil base saturation and soil pH

The term “soil acidity” expresses the quantity (expressed in meq/100g) of the acidic cations (cf. part 3) that the soil can hold on to. The percent base saturation – another important term on your soil test results – is the percentage of the soil’s cation exchange capacity (CEC) occupied by the basic cations.

This is from our soil test:

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This means that calcium occupies 50.6% of the total exchange sites. In other words, in 100g of my soil, 15.6 meq can hold on to cations, both basic and acidic. Of that, 7.9 meq is occupied, or saturated, by calcium, 1.65 meq by magnesium, 0.64 meq by potassium. So, as far as I can learn from the test results (*), 10.19 meq/100g of soil, or 65.3% of the CEC, is saturated by bases. That leaves 35.3% of the CEC (*) for the acidic cations (hydrogen and aluminum).

(*) Sodium (also a base cation) is not listed on my test results, which means its levels are low, so I don’t have a sodic soil (cf. part 5).

Not surprisingly, the greater the percent base saturation, the higher the soil pH. Because calcium is normally the major cation, by virtue of its abundance taking up about half the CEC (as in our soil), we can say that there is less calcium in acid soils and more in alkaline soils.

But if the soil is very alkaline (pH > 7.0), the high levels of calcium may have negative effects. For one, more calcium taking up the CEC very simply means that there is less room on the colloid for everything else. Secondly, an excess of calcium can no longer be adsorbed onto the colloid. This “free” or unadsorbed calcium begins to accumulate in the soil water and goes on to react with what other nutrients are present.

For instance, the free calcium will readily attract soluble boron (B-), which is an an-ion (a negatively charged ion), and form a nearly insoluble compound with it, thus making the boron less available to plants.

Excess calcium will also tie up, or immobilize into insoluble compounds, cations like iron (Fe++), phosphorus (P+++) aluminum (Al+++), zinc (Zn2+), copper (Cu2+), cobalt (Co2+), and manganese (Mn2+), as well as magnesium (Mg ++) and potassium (K+).

Lastly, calcium also increases the pore space in the soil by flocculation, which, as we saw in part 5, is desirable. But when pore space exceeds 50% of the total soil volume, the soil can dry out much easier, like sand.

In short, too much calcium in your soil and many nutrients become insoluble and thus unavailable to plant roots, and the soil structure is damaged to boot.

But, on the other hand, if the soil is very acidic, and thus if there is not enough calcium, many of the other cations can become excessive and thus toxic. Then calcium applications with limestone are called for. The aim when attempting to adjust soil acidity is never so much to neutralize the pH as to replace lost cation nutrients, particularly calcium.

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Next time, in Part 7, I promise, we’ll finally meet the plants, and discover by what magical means they get the calcium out of the soul soil. 

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This is part 5 of a series on how nutrients, mainly calcium, get into our soil and vegetables (click for part 1, part 2, part 3 and part 4). It is the longest and most difficult part of my expose, and the least “popular” one, judging by the fact that the issues discussed will not show up on the average soil test. Still, I include it because it gives us something to think about when we irrigate our garden and – I admit it – because it introduces that most enchanting of words in soil science. Flocculation. Come on, say it, out loud, taste it! Now you have to find out what it means.

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5. Soil structure, flocculation, salinity and sodicity

As we saw, it’s good to have some amount of clay, as clay particles are negatively charged and thus able to attract and hold on to nutrients, which are positively charged. Now clay particles can either be unattached and dispersed, or clumped together, “flocculated” into aggregates (flocs = flakes).

Flocculation happens because opposites attract and like repels like. Thus one negatively charged clay particle will repel another negatively charged clay particle. But the positively charged cations create bonds between them, shaping them into clumps or flakes.

Because of their varying charge, certain cations are good flocculants, like calcium (Ca++) and magnesium (Mg++), whereas others are poor flocculants, like sodium (Na+) and potassium (K+). Add water in the mix, and the cations’ flocculating power diminishes, because the cations will also spend some of their positive charge on attracting the hydrogen ions (H-).

Flocculation is a good thing. Unattached, dispersed single particles sit together in a dense cement that allows no air pockets, called pores. Clumpy aggregates, on the other hand, will not fit together so perfectly and create pores. It is in these pockets that the rapid exchange of air, water and colloidal cations with plant roots can take place. It is also in and through these spaces that roots grow.

But in such a lively realm as soil, flocculation is a transitory thing. It is best if the aggregates are stable, which stability depends on (1) the amount of soluble salts in the soil, and (2) the balance between calcium and magnesium (the more powerful flocculators) and sodium (the weak one).

As for (1), had I known about it, I would have shelled out the extra $5 for a soluble salt test to be done on my sample. Soluble salts are any dissolved ions, be it calcium, sodium or potassium. Ions in solution conduct electricity. The extra test would have given me the electrical conductivity (EC) of my soil, which would have given me another indicator of its nutrient richness.

As for (2), that extra test would have enlightened me about the balance between calcium and magnesium on the one hand, and sodium on the other, as it would have given me the Sodium Adsorption Radius (SAR):

[Na+]
————–
[Ca++] + [Mg++]

How do EC and SAR matter?

Well, flocculation or aggregate stability occurs (1) if the amount of soluble salts (calcium, magnesium as well as sodium) in the soil is increased: more positive ions means more electrical conductivity (EC), which means more binding of clay particles into clumps. Conversely, soil particle dispersion occurs when the amount of soluble soils and thus the EC is decreased.

Soil particles also flocculate (2) when concentrations of Ca and Mg are increased relative to the concentration of Na ions (that is, when the SAR is decreased), because Ca and Mg are much stronger flocculants. Conversely, soil particles will disperse when the SAR is increased. (I recommend this powerpoint presentation for a more visual explanation of these interactions.)

As we saw, hydrogen anions (H-) diminish the soil’s cations’ flocculating power, so irrigating with “pure” water – water that has low amounts of soluble salts and is thus a very poor conductor of electric current (EC) – can destabilize soil aggregates.

If you irrigate with so-called saline water – water with a high EC, or high amount of soluble salts – then that soil will have a good structure.  However, as can be expected, if there is an excess of salts in the root zone, it will hinder plant roots from withdrawing water from the soil (this will be further explained in part 7).

Another word of caution: if you have sodic irrigation water, that is, if it contains a high amount of sodium (Na), it could damage your soil structure, making life difficult for plant roots and causing problems with irrigation.

That is because Na ions are larger than Ca and Mg ions. When too many large sodium ions (with their low flocculating values) come in between the clay particles, they act like wedges, separating the particles, breaking up their aggregation. This soil dispersion causes the clay particles to plug the soil pores and create cement.

If you soil cracks when it is dried up, you have a sodic soil. One of the solutions is to decrease the SAR by introducing calcium (mostly in the form of gypsum), which will compete with the same spaces on the colloids as the sodium, and flush them out.

Something to think about, when we water our garden!

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I really did enjoy that – no kidding. I used to study metaphysics in grad school and this reminds me of it, a bit. Let me know what it did for you!

 

Undaunted, let’s move on to Part 6.

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We’ve reached part 4 of this riveting story of how calcium and other nutrients make it into into the soil and thence into our vegetables and thence into our own bodies (and into chicken eggs). We’ve had some cliffhangers already, so be sure to check out parts one, two and three.

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4. Solubility, carbonation and chemical weathering

It is the reaction of calcium and calcium compounds with water (see last part) that makes them soluble. Solubility or dissolution is the process by which a “solute” forms a homogeneous mixture with a “solvent” (here water).

This happens as the solute molecule breaks down and its ions dissociate. The positive ions attract the partially-negative oxygen in H2O and the negative ions attract the partially-positive hydrogen in H2O. The ions thus get spread out and become surrounded by the water molecules. The dissolution is complete, or in equilibrium, once it’s all been spread around. (source)

One calcium compound is more soluble than another. Calcium carbonate (our eggshell) has a very poor solubility (47 mg/L at normal atmospheric CO2 partial pressure and 25 degrees C). As we shall see, this is important for gardeners who plan to enrich their soil with eggshell calcium, but I will come to that later.

However, if carbon dioxide is also present, that carbon dioxide will react with the water to form carbonic acid (H2CO3), which is a weak acid – it’s the bubbly in our soft drinks. This carbonic acid will in turn react with the calcium carbonate to form calcium bicarbonate (or calcium hydrogen carbonate). So

CaCO3 + CO2 + H2O → Ca(HCO3)2

Calcium bicarbonate is five times more soluble in water than calcium carbonate—in fact, it exists only in solution.

This is the main process by which carbonate rocks of the Earth’s crust are weathered. As we saw, if water is saturated with carbon dioxide, it produces a mild carbonic acid. This is what happens with (unpolluted) rainwater (water plus atmospheric CO2), which has a pH of around 5.6 (polluted, “acid” rain has a pH of as low as 3.0), and with water in aquifers underground, where it can be exposed to CO2 levels much higher than the ones in the atmosphere.

In a process called carbonation, this water’s carbonic acid reacts with the solid calcium carbonate in rocks like limestone or chalk, forming calcium bicarbonate and dissolving it. This solution of water and mineralized calcium is then borne off into the soil, where it is deposited on the colloid, and where it waits to be again dissolved in water and made available to plant roots.

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Well. That just brings us back to the beginning!

On to Part 5

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This is the third article in a series on how calcium and other nutrients end up inside our vegetables, and on how to interpret certain soil test results. It is preceded by part 1 and part 2.

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3. Water and pH

Let’s investigate the water in the soil. For one, water brings the minerals to  the colloid, and it can take them away again (but so do the soil critters). Also, for reasons that will become clear later, calcium is available to plants only in dissolved form, that is, as part of a solution in water. Thirdly, this watery context heavily impacts the lives of soil critters. The most important factor in all these matters is the water’s pH or acidity or alkalinity.

As we saw, a molecule of water is composed of one oxygen atom and two hydrogen atoms: H2O. In a vat of pure water, most water molecules remain intact, but a very small amount of them react with each other in the following manner:

H2O + H2O ===> H3O+ + OH–

Water + Water ===> hydronium ion+ (an acidic cation) + hydroxyl ion– (a base)

The hydronium ion ( H3O+) is the chemical unit that accounts for the acidic properties of a solution, and the hydroxyl ion (OH–) is the chemical that accounts for the basic or alkaline properties of a solution. How?

Well, in pure water, the amounts of H3O+ and of OH– are equal, so the acid and the base cancel each other out, so pure water is said to be neutral, with a pH close to 7.0. Also, in pure water the concentration of H3O+ and OH– are in balance, so that an increase in the concentration of H3O+ causes a proportional decrease in the concentration of OH–.

This means that, if you add an acid like hydrochloric acid (HCl) to water, it reacts with some of the water molecules like this:

HCl + H2O ====> H3O+ + Cl–

And this increases the H3O+ or the acid concentration, throwing off the balance and lowering the solution’s pH to below 7, making it acidic. But if you add a strong base, such as calcium, to the water, it ionizes as follows:

Ca + H2O ====> Ca(OH-)2 + H2

Thus, the addition of calcium to water increases the OH- or alkali concentration of the resulting solutions, making the solution alkaline.
The cations (positively charged ions) we’re interested are either bases or acids:

  • Basic cations: calcium (Ca++), potassium (K+), magnesium (Mg ++) sodium (Na +)
    Acidic cations: aluminum (Al+++) and hydrogen (H+)

The pH of the water that saturates the soil (see 4) regulates the solubility of minerals in that soil (see 4), thus their availability to plant roots (see 5), as well as the activity of soil bacteria (see 6).

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Mm, on to Part 4: that pesky problem of solubility, which took me a while to understand. For now let me add that I forgot all the chemistry I learned in secondary school (way back), and that this excursion has proved to be a fantastic rediscovery of all that magic.


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This is the second article in a series on how calcium and other nutrients end up inside our vegetables, and on how to interpret certain soil test results. You can read the first part here.

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2. The soil’s ability to hold on to this calcium: colloids and cation exchange capacity

The soil needs to store the weathered bedrock calcium, keep it from leaching away to where the plant roots can’t reach it. Certain clays (very fine inorganic particles) and organic or humus particles form colloids: thin, flat plates with a large surface area that has a negative electrical charge.

As such they attract or adsorb large quantities of positively charged ions, or cations (pronounced cat-eye-ons), which comes to adhere to it with a weak electrochemical bond. As the cations are taken from this storehouse by plant roots (see 6), other cations in the soil water replace them on the colloid. This is called cation exchange.

For example, hydrogen (H+) is such a cation, and it’s part of H2O, so water molecules are attracted by the clay’s negative electrical charge.

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That is why clay is good to have, since it holds on to water so well. But if your soil is mainly clay, it’s not so good, because the clay will lock the water molecules into bond that is so strong that the plants’ roots will not be able to break it. So the soil may hold a large amount of water, but much of it is actually not available to plants. On the other hand, a soil that is too sandy) will not be able to hold on to anything, water or nutrients, because sand particles have no electrical charge.

Cations will compete with one another for a place on the colloid. Some are charged more than others, for instance, calcium (Ca++) is charged twice as much as hydrogen (H+), and aluminum (Al+++) is even stronger. An aluminum cation borne along by the soil water can easily knock a hydrogen cation off the colloid and take its place. If the soil water bears a high concentration of a particular cation, those cations will replace other cations on the colloid, for instance, a swarm of calcium ions against one sodium (Na++) ion.

This image (source) shows the flat, plate-like structure of the colloid and the negative charges along its edges:

cec-diagram

http://www.dpi.nsw.gov.au/agriculture/resources/soils/structure/cec

The stronger the colloid’s negative charge, the greater its capacity to hold the positively charged cations. The kind of clay (for more details, see here), its mixture with sand and organic matter and its crumb structure, all determine the soil’s capacity to hold cations against leaching. This capacity is termed the soil’s cation exchange capacity (CEC) – an important term on your soil test results. This CEC is expressed in milliequivalents per 100 grams (meq/100g) of soil or in centimoles of positive charge per kilogram of soil (cmol(+)/kg).

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On our soil test results the CEC is given as 15.6 meq/100g. So, in every 100 grams of our soil, 15.6 meq of soil can hold onto the goodies, calcium (Ca++), potassium (K+) and magnesium (Mg++), that come along in the soil water, as well as hydrogen (H+), and sodium (Na +) and aluminum (Al+++), which are not plant nutrients.

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That explains at least one unknown on that soil test! Next up, in Part 3, the role of water in all of this: it is, after all, the soil water that brings and that can take away these cations.


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It was the combination of finding an eggshell in the compost and staring at our soil test results that did it. I started researching and one thing led to another. But I figured it out, the basics of it, anyway. The result is a long text, so I’m serializing it over the next couple of days. I hope you find this sort of thing as fascinating as I do, and that it will help you with your own soil test results.

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Digging and moving my compost heaps over a month ago, it was interesting to find eggshells and bones, all over a year old and barely decomposed. I keep sifting them out and putting them back into the heap. I figured it is the high amount of calcium in them that makes them so hardy.

This made me wonder how that calcium will ever make it into my garden vegetables. In what shape or form can the calcium in the eggshell, and the calcium originating from the bedrock, be taken up by plants?

It took me a long time to find out, but I realized early on that we first need to get clear on what this “calcium” is that we’re talking about.

Calcium in its elemental state is a metal – a soft gray alkaline earth metal (Wikipedia). But it is so reactive that it is never found in its elemental state, that is, all by itself. It readily combines with whatever it comes in contact with and becomes part of a compound.

Thus in the eggshell, calcium occurs as calcium carbonate (CACO3): an eggshell consists of 94 to 97% calcium carbonate. Calcium carbonate is also the active ingredient in agricultural lime, which is mined from limestone. Calcium can also occur in the soil as part of the compound calcium bicarbonate (Ca(HCO3)2), or as part of the compound calcium nitrate (Ca(NO3)2). And so on.

The “calcium” we will be talking about here is the calcium ion in each of these compounds, because it is only this ion that can be taken up by plant roots and thus form a nutrient.

If our interest is in calcium as a nutrient for plants, we need to consider:

  1. its presence in the soil
  2. its continued present in the soil
  3. its interaction with water, and pH
  4. its solubility
  5. its relation to soil pH
  6. its uptake by the plant
  7. its availability when tied up in organic materials

1. Original presence of calcium in the soil: parent materials

First, of course, calcium needs to be present in the soil. So where would it come from in the first place? Calcium is abundant: the third most abundant metal in the earth’s crust, accounting for 3.64% of it. It is also, by the way, the fifth most abundant element by mass in the human body: what a beautiful correspondence!

Soil originates primarily from parent materials: rock. When rock is weathered or eroded, it is broken down. This can be a physical or mechanical process (abrasion by ice, for instance, or by biological agents such as lichens and mosses), or a chemical process (see 4). The left-over materials, when combined with organic materials, make soil, which is thus (almost) nothing but rock. I recommend the excellent soil page on Wikipedia.

Obviously, the more calcium (as well as N, P, S, K, Mg and all the micronutrients) is contained in the bedrock in your area (and the younger the soil), the more calcium will be present in the topsoil. So it pays to investigate what kind of rock your garden sits on: spend a day or two (or three) navigating the USGS map library, especially the surficial earth materials and bedrock lithological maps of your region.

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(http://www.mass.gov/mgis/bedlith.gif)

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Intrigued? This is only the first part of a series on calcium and other nutrients in soil and plants. Stay tuned!

Part 2 can now be read here.


While the dust mites in the bedding were freezing (to death, hopefully)…

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We warmed up this cold but sunny morning by splitting and sorting firewood, playing with a neighbor’s dog, and scouting out some animal tracks. Here are some tiny bird tracks next to my fingerprint:

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And these are some huge bird prints next to my footprints:

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And these claw marks are interesting:

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Then back inside for a steaming cup of tea and some reading of (library) books on bees, and chickens.

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It’s coming down hard: thick globs of melting snow. The wood stove is giving off enough heat to dispel any gloom: it’s merely cozy, as long as I don’t need to go out there.

Which I did have to, earlier on. One of the rain barrels was overflowing, and not through the overflow tube. In this weather I would have left it but the excess water was undermining the cinder blocks the heavy barrel is sitting on, slowly eroding away the soft soil. I didn’t relish the thought of it coming down right by the side of the house and the bed with the chard.

So out I went, and shook the overflow pipe, but nothing came out but a dreadful stink. O-ow, dead animal alert! I opened the barrel’s lid and saw the hind part of a chipmunk sticking out of the overflow pipe. It must have crawled up the pipe in drier weather, landed in the water, then made it back to the pipe only to get stuck.

It had that ghostly look of a thing dead in water. That half looked well preserved in the cold water, and I only considered for a second what the other half looked like. When I tried to dislodge it with a stick its skin just came off. I un-threaded the pipe and as the excess water suddenly rushed out all over me I shook the poor dead beast out in the bushes.

I usually take a picture of any dead animal I see (here and here and here) but this one, well, it was just too gruesome.

We’re spending the rest of the day inside, drawing animal tracks in snow. Squirrels, deer, chipmunks…

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Canned Bell Peppers (from Farmers Market)

Amie and I just ate our first three peas from the garden: delicious! She also ate seven green beans (from the garden) and four carrots (disks, that is, not the whole thing) (from the Farmers Market). She loves the beans especially. She also promised that when the time comes to break it out of the freezer, she’ll eat the pesto for which she harvested and trimmed the basil.

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A grasshopper was hiding out in the basil and luckily jumped out before I put the batch in the food processor. We caught it and let it outside, where it sat on Amie’s finger for a long time, then gave her a fright when it jumped.

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Later on a Ladybug came to visit and Amie fed it a blueberry.

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I do love those still-pudgy hands…