Now, on to the grumpy stuff:
I am very curious about soil science, and would particularly love to learn indicator species for identifying soil resources / deficiencies without resorting to lab tests.
However, I have done a bit of basic chemistry education, and I had some problems with some of the info presented.
The errors and typos that I can recognize based on my limited knowledge, make me nervous about taking anything else in this resource book as gospel.
I estimate that I'm finding one error about every two pages - or about 10% of the info I am qualified to critique. In the classroom, 90% is an A/B. But in actual biology, 98% (DNA similarity) is where you get a chimpanzee instead of a homo sapiens. I don't feel that 90% is enough accuracy to rely on for a subject this complex.
If every other person reading these threads will submit corrections / caveats when their own experience or training raises questions about the manual, it will give me more confidence about which info is useful.
(I suppose it would be even more confidence-boosting to give credit to the accurate bits, so feel free to be a non-grumpy critic if you feel like it!)
p. 182 - section 8.1 paragraph 6 and 7 of the introduction
Mollison introduces this chapter with some caveats about scientists not being involved in field work.
Many scientists _are_ involved in field work. Many citizens are also involved in wide-scale scientific research - great examples of 'citizen science' include birdwatching clubs and the Audubon Society providing detailed info about migration routes, and returning banded birds to research programs; amateur astronomers discovering a vast proportion of the known comets and near-earth asteroids; and amateur weather-watching data collected for example the 4" Rain Gauge collection project
in the USA (exact society and website forgotten, but look it up if you want to participate - there are far more farmers with rain gauges than there are official weather stations, so more fine-grained data can be collected by citizens).
enthusiasts are equally dismissive of scientific and field work ("farmers are destroying our soils"), yet believe in wildly unproven experiments.
I keep hearing from permies who started with herb spirals and raised mulch gardens, killed a lot of plants both pre-existing and cultivated, then made a second transition
to more locally-appropriate practices (like integrated range management for drylands).
From Mollison's writing, I can't tell which of the proposals he advocates have actually proven effective; let alone in what climate and biome.
Is every method being advocated because it has produced good results?
Or are half of these ideas unproven "good ideas" that will be tried, without feedback, by hundreds of permies long after some of them have proven to be very bad ideas for specific situations?
I'm looking forward to specific chapters on specific conditions.
I wish more of the chapters were arranged with conditions first, treatments and techniques divided by appropriate conditions.
My chemistry caveats for this chapter:
p. 190, Nitrogen discussion: Confusing labeling of terms.
paragraph 3 - "More than 80 ppl in waters are lethal to young animals and children" - What's a PPL? Typo for PPM (parts per million), or for g/l (grams per liter)?
Paragraph 4: EU standard allows 50 mg/l as acceptable in drinking water (50-80 ppm).
Why the broad range? it does suggest that the 80 figure above may be either 80 ppm or 80 mg/l.
Then he offers this chart:
Some nitrite (or nitrate) levels (ppm)
0-50 Safe for human consumption
500 Spinach grown from compost
2000-3000 Chemical fertiliser-grown spinach
Now, that is just plain misleading.
- There is a difference between nitrite and nitrate; I'd love to know more about it. If the tests can't tell the difference, reports about human safety will be contradictory and misleading.
- Show me one child, anywhere, ever,
who routinely consumes the same weight or volume of spinach compared with their consumption of water (and water derivatives like reconstituted juices and Popsicles). It would be seriously unnatural. Children in general are well-protected from the dangers of overeating spinach, as any parent who has tried to feed
it to infants can attest. They are usually good for about 6 bites, tops, and that's only if it's well-paired with a juicy sandwich or salad-dressing.
We use water in our bodies to dilute and distribute other nutrients; we eat foods for specific nutrients, flavor, etc.
ANY food should
contain much higher levels of specific nutrients than the local
water does, in order for both food and water to serve their proper functions.
- What are the nitrite (and nitrate) levels for other common foods, like preserved meats? What is a reference level for spinach grown in sand instead of compost
This chart is incomplete and misleading, and seems intended to frighten people
into thinking that spinach is lethal.
That implication can be disproven with a moment's thought, by reflecting on the number of spinach salads (both compost-grown and fertilised) which you have already eaten without lethal effects.
Mineral content in water is dangerous because it becomes part of EVERYTHING you eat and drink over time, and it is impossible to go without water for more than a few days. The minerals concentrate in soups, teas, irrigation waters, as coatings on rinsed dishes, and so on. When the water you drink contains excess concentrations of the same minerals your body is trying to excrete, like salty seawater, it becomes impossible for your body to excrete it faster than it comes in. Shipwrecked sailors who drink seawater go delerious and then die of kidney failure.
If you are over-exposed to potassium from food sources, you can go for years without eating a banana. If it's in your water, it's hard to avoid.
p. 194, Iron: "Meats such as veal, liver ... can provide dietary iron."
These examples are from opposite ends of the red-meat spectrum.
Veal is calf's meat from calves that have only nursed and never eaten grass - or in the US, from older beef cattle
that have been continuously and deliberately deprived of iron in order to produce a pale, mild-flavored meat for delicate dishes. (It's the subject of ethical protests, because the commercial method involves caging them away from sunlight, with seriously limited exercise, to prevent muscles developing to mature toughness. The cage must be free of iron bolts or latches, because the deprived cattle will lick at any iron source as normal ruminants lick at salt or other minerals.)
Liver, on the other hand, tends to be a concentrated source of digestible iron whether from beef, lamb, pork, or other sources. Good, healthy liver meat is so nutrient-concentrated that most people only need a few bites to feel satisfied, but when my husband gets anemic he tends to cook and eat about 2 lbs.
Listing veal instead of beef or lamb, red meat, or another organ such as heart, is a weird choice.
Antimony: As for arsenic.
(Really? Horses, chickens
, all that? Does the similarity extend to matters not discussed, such as the use of arsenic for pressure-treated lumber or pest poisons, or the concentration of this mineral in apple
8.8 pH and Soils:
The discussion gives % silica (SiO2
) for acid, intermediate, basic, and ultrabasic soils.
What is the other proportions? I would assume calcium-rich materials like limestone and chalk, since they are basic and can be plentiful in soils.
If a soil was, for example, 50% silica and 50% humus, without any alkaline minerals, wouldn't it still be really acidic?
p. 196: Acid and Alkaline Waters and Soils
"soft" waters include not only acidic waters, but also neutral waters without dissolved minerals. "Hard" waters are generally alkaline, as described. I believe they're considered "hard" not only because they don't readily dissolve soaps, but because they are more likely to leave hard deposits on pots and plumbing. Acidic waters are more likely to dissolve metal plumbing than to leave deposits. But neutral, soft water neither leaves deposits, nor dissolves plumbing as quickly as acids do.
I'm going to plug in some basic chemistry education here, because I think this section was unnecessarily hard to understand. For this chapter's discussions of ions and elements, it helps to know what ions are, and what forms most salts and minerals will commonly take. If you are already comfortable with this, skip
to the next elipses (....) mark.
Ions are charged particles. They are classed as positive (having more protons than electrons, and a positive electrical charge) or negative (having more electrons, and a negative electrical charge). Some are made up of single atoms of a given element (like Na+ sodium, or H+ hydrogen, or Cl- chlorine). Others are made up of little bundles of atoms that cling together very strongly (SO4-- sulphate, NH3- ammonium ion). In solid form, they tend to stack up into crystals with each + ion surrounded by - ions. We call these crystals by the smallest neutral unit they are made of (NaCl, CaSO4), but in reality it would be difficult or impossible to assign particular ions in particular pairs. They aren't molecules like we think of a molecule of benzene, oxygen, or carbon
dioxide, which float around in air and bump into other molecules. They are just neighbors with an extremely strong grip. (This also makes it possible to have crystals with random inclusions of "other" molecules - like different-colored gems with the same basic chemical structure.)
Water is an ionic fluid, or solvent. Charged particles dissolve easily in water. Once dissolved, they swim around separately from each other, and don't re-combine into crystals or neutral molecules until the party's over.
Acid and alkaline ONLY apply to solutions in water.
Water's chemical formula is H2O; if you look at the molecule it's kinda like a Mickey-mouse head, with two little hydrogen "ears" on a bigger oxygen "head."
When you dissolve other things in water, sometimes it produces an excess of ears (H+) or heads (OH-). (Oxygen is too greedy to let go of both hydrogens at once without a seriously better offer.)
Acidic (excess H+) and alkaline (excess OH-) solutions turn out to have really different properties, so they're very important concepts in chemistry.
Metals dissolve easier in acidic solutions. Acids also have big effects on proteins.
Alkaline solutions don't tend to dissolve metals, so slightly alkaline tap water helps prevent lead and copper from the pipes. But alkaline solutions do dissolve (saponify) fats, and thus can be very disruptive to most forms of life (most cell membranes are oil-like).
We humans make a very strong acid in our stomachs (HCl), and can digest acidic fruits down to pH 2 to 4 (lemon and tomato
). But most of our food is mildly acidic (milk
) to neutral (water, bread).
We can also digest some slightly alkaline materials (baking soda is pH
. Strong alkali are very caustic, causing severe skin burns and poisoning reactions.
Acid and alkali neutralize each other, but that doesn't change the total amount of minerals present.
If perfectly equal quantities of acidic (excess H+) and alkaline (excess OH-) solutions, they combine to make water (H2O), and a salt.
For example, table salt could be made by combining NaOH (lye, an alkali) and HCl (hydrochloric or stomach acid) to make NaCl and H2O).
Some salts dissolve easily in water (like NaCl), others don't (like BaSO4).
Plants uptake most minerals more easily if they are dissolved, as opposed to sitting around as little pellets of rock or metal. But easily-dissolved minerals also wash away easily in rains, unless they are held by something (like the gels mentioned).
Where rainfall is more than evaporation, salts wash downstream and down into the deep soils. Where evaporation exceeds rainfall, water carries salts upward in the soils, and can leave a crust of soluble salts on the top of the soils.
Excess salts can suck water out of plant roots, just like salt helps us dry and preserve foods like bacon
or salt cod. (This is true of ANY dissolved mineral, including fertilizer; this is part of how excess nitrogen "burns" plant roots, by totally drying them out.) So excess salts can be a bigger problem than just scarcity of water.
It's also possible for a particular mineral to change how easily-dissolved other minerals are - like acid making iron more available.
This is a subject I have had to explain in public; many, many times; to children of all ages, often in the presence of their doting PhD parents or grandparents.
When I screwed up or mis-spoke, the PhD's would quickly correct me, and if unchecked could make the subject very, very complicated for said children. (I enjoyed these discussions when children were not present, but also strove to keep them from getting bored.)
So I learned to be very accurate, while giving the minimum information necessary to explain a particular fun experiment (like launching a rocket using H2 and O2).
There are many parts of basic science where I readily admit ignorance, but in simple explanations of inorganic chemistry (ions, metals, salts) I'm pretty confident.
So, back to the chapter:
"In solution, metals release positive ions (H+). Nonmetals release negative ions (OH)."
This at first appeared nonsensical.
If I dissolve a little lump of metallic sodium in water, it floats around spitting and releases a lot of both H and OH. The H is released not as ions, but as hydrogen gas (flammable bubbles). If the lump is about half the size of a pea
, it emits enough heat to light the hydrogen on fire. We didn't tend to use bigger lumps, on account of they can release so much hydrogen and heat that the flask explodes. but back to the ions....
The OH is released as ions (OH-), and makes the water very alkaline. The sodium loses an electron, and goes from neutral (shiny metal) to being a positive ion (a charged particle, in this case invisible as it is dissolved in water).
Did the sodium "release" positive ions of H+? No, it released negative ions of OH-, and positive ions of its own self (Na+).
So, I would say it like this:
In solution, metals become
positive ions (Fe++, Na+, Ca++, K+, Cu++). Many of these are associated with alkaline conditions (excess OH-).
negative ions (O--, Cl-). Many of these will form neutral or acidic conditions (excess H+).
When elements combine into more complex ions (SO4--, CO3--), their properties change. Ammonium ions (NH4, NH3-) create alkaline solutions.
Carbonate ions can act as buffers by changing from (CO2 to CO3--), keeping pH more stable in specific ranges. Carbonate ions help buffer our blood, keeping it from changing pH when we eat a meal or breath more heavily.
Figure 8.5 - Mineral availability:
"The curves represent pH values" - is this two diagrams overlaid on top of each other, with the white and shaded curves representing the logarithmic scale of pH, and the black bars representing mineral availability?
In real solutions, H+ and OH- ions are both available at all ends of the pH spectrum; they are just more prevalent at their respective ends. The curve representation is neither clear nor accurate. It'll have to get by on its good looks.
Table 8.5 COMPOSTING:
What exactly does he mean by 'noxious?' I've always understood it to mean dangerous to health - usually with an implication of contagious; a noxious weed is one that spreads unchecked, and is usually considered dangerous to livestock or native species. Noxious wastes might include the categories commonly called biohazards (blood, bodily fluids, contagious matter), and perhaps also any material that is toxic, stinky, or difficult to dispose of (animal offal).
The chart seems at first glance to have one, or possibly two, reversed labels. Composters, please confirm:
- The second level of the flow chart, which asks "noxious," leads to a central column with minimal processing (pulverise) and two side columns with elaborate processing (feeding to poultry, hot composting, fermenting).
At the box marked "noxious," answering "N" leads to the complex paths, while a "Y" answer leads to the simple (pulverise and apply) path.
Shouldn't the elaborate processing be applied to "NOXIOUS" materials, and the simple methods for direct application to non-noxious materials?
Or is pulverizing meant to involve extremely fine grinding and spray-drying, to achieve something like blood-meal that can be applied before it rots?
Examples as currently labeled: The only "noxious" materials described are at (e).
(a) nut husks and shells, bark, woodchips, canvas, are examples of dry, non-noxious materials that would be mulched.
with seed heads, weeds in flower, bulb, or root
of weeds - not noxious, but would be fed to poultry before mulching.
(c) sewage and sullage, liquid manure and urine
, meat and animal paunches... would be considered not noxious (??), sprayed or hot composted.
(d) sludge from digesters and weed-free manures fall under the tree at "Y" noxious, no weed seed
(e) chicken manures, blood, bone, etc. are considered noxious, would be pulverized and used as manures on gardens.
(f) dissolved minerals, compost teas are considered suitable for spray
I think that (c) and possibly (d) would also be considered noxious - if blood and chicken
manures are noxious, surely animal paunches and meat are as well?
(d) digester sludge could be considered non-noxious if the digester is verifiably working propertly to dispose of noxious materials and convert pathogens to safe, soil-type bacteria by the end of the digestion process.
Should the upper-right label be reversed?
I agree that the info about organic materials making gels is pretty cool.
I suspect that mycorrhizal webs are an even more important, gel-like functional storage and transport for soluble nutrients.
I disagree that synthetic water-sol gels are an appropriate replacement - I don't know anything horrible against them, it just seems unnecessary.
I've also heard people talk about 'biochar
' as a nutrient sink, holding nutrients and water somewhat like gels, wood fiber, or perlite.
I think living humus is probably better than biochar; biochar makes a difference where soils are very specifically hostile to humus formation or retention (humid tropics). I've also seen somewhat reliable reports of biochar helping to improve heavy clay soils in temperate rainforest conditions.
If you don't live in a rainforest, please do small-scale testing before embracing biochar as a soil amendment, and if results seem promising extend the testing incrementally (not all at once).
Biochar production at home tends to consume time and energy, produce a lot of smoke, and waste a lot of the nutrient value of the substance being burned (fragile proteins, acids, fats, etc). So I'd want very good evidence that charcoal is somehow superior to ordinary mulching methods. One of the few appropriate ways to make biochar, in my opinion, is to sterilize noxious wastes that are in serious excess, like rhizominous weeds (blackberries, thistles), or chicken manure in the case of very large poultry farms producing excess nitrogen at near-coastal (urban belt) locations. If you burn a dirty stove and have charcoal left in the ashes, by all means use it. But burning dirty fires has never been my favorite way to save the world.
Grumpy chemistry updates concluded!
It's entirely possible that things just aren't making sense to me because soil science is such a big, complex field, and I don't have much experience testing for specifics or observing interactions.
I've gardened in two very different climates, and not had any major problems with nutrients. In any household that imports food, the compost can easily replace more than the harvest consumes. I haven't had to face nutrient replacement for a market garden, or large-scale gardening
to the point of erosion problems.