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Understanding Soils and Nutrients

Part 3 - The Chemistry of Plant Nutrients

Jim Barrow

What is all this about? Native plants don't need fertilisers - or do they?

Well they wouldn't if all native plants were the same; if they all grew naturally on very low fertility soil; if all the soils on which we try to grow them provided sufficient nutrients; and if all gardeners were satisfied with very slow growth and often a somewhat unimpressive appearance. Not all of these conditions are going to be filled, so it behoves us to learn just a little about fertilisers and what happens to the nutrients.

The last time I checked it out, there were 18 nutrient elements needed by plants - the list keeps growing as techniques become more refined, with nickel being only recently added. Some of the elements come from water and air (hydrogen, oxygen and carbon). Others are so plentiful in the environment that deficiency does not occur (chlorine and (seemingly) nickel). Three others we have already discussed in Parts 1 and 2 (nitrogen, iron and manganese). This leaves a mere 10 to worry about.

The problems of charge

But before we do this we have to worry about something else. Remember, way back in Part 1 I said that some of the soil particles carried an electric charge? Before we talk about how ions react with soil we need to learn something about that charge. And we have to start by thinking about crystals.

All crystals are made from neatly organised atoms. The atoms may be all of one kind as carbon atoms in diamond, or of several kinds as in the minerals found in soils which may contain silicon, oxygen, aluminium, magnesium and other atoms.

Now there are two tricky things about soil minerals. First, we need to know that the correct "recipe" for some of the minerals in soil doesn't always get followed. Some of the spots which "should" have a silicon atom have instead an aluminium one. Some which should have an aluminium atom have a magnesium. These "wrong" atoms just don't have enough charge so the mineral as a whole has a negative charge. These substitutions are deeply buried in the mineral so they are called "permanent" charge.

The other tricky thing concerns the atoms at the edge of the mineral crystal. The atoms in the middle of the crystal are all happy because they nicely linked in a neat pattern with (mostly) the correct partners.

But what about the ones right at the edge? They are OK on one side but rather lonely on the other. Especially if the minerals are oxides, they end to grab water molecules to compensate. These water molecules then tend to react with hydrogen ions in the soil water (remember the article on pH). So they become positively charged if there are a lot of hydrogen ions about (low pH) or negatively charged if there are not many (high pH).

The charge varies with pH - and, appropriately enough, is called "variable" charge. There is yet another source of charge in soils, which we will consider when we get to talk about phosphate - but for now, let's just accept that the soil particles do have a charge.

Potassium, Calcium and Magnesium

Now, back to the nutrients. For three of them charge is (almost) everything. They are rather independent critters and don't like holding hands - but they are really fond of the opposite charge. They are potassium, magnesium and calcium.

They are not clones of each other. There are indeed important differences between them in the way they react. But all three are held by the negative charge on the soil particles.

On "real" soils there is plenty of negative charge (enough cation exchange capacity in the jargon). These nutrients are held nicely in the soil and there is little chance of them being leached out. Alas, some areas (eg Perth) don't have real soils, just sands and there often isn't enough charge to hold things. Of the three, potassium is the one likely to be least tightly held.

"....potassium is a nutrient you may have to apply fairly often.....if you cut the plants and cart the material away, you lose a lot of potassium."

So in such areas potassium is a nutrient you may have to apply fairly often. And plants need a lot of it; if you cut the plants and cart the material away, you lose a lot of potassium. This is what happens when you throw away lawn clippings. Not only are they a good source of organic matter but they are high in potassium. However as soon as the leaf cells die, the potassium is freed and can be easily washed out. It will find its way rapidly back to the soil if the clippings are used as a mulch but may get lost from compost.

Lawn Mower

Potassium has another trick which we should mention ...... Some of the soil minerals are built rather like stacks of ham sandwiches and potassium can get in among the ham and hide. Like many things, this is both good and bad.

It means that there is a nice reserve of potassium which can be rationed to plants but it also means that fertilizer potassium may not be as rapidly available. It seems that you only get ham sandwiches in young soils. In Western Australia almost all of the soils are very old and the ham sandwiches have "gone off". They have turned into a different kind of mineral constructed more like a loaf of sliced bread than a stack of ham sandwiches and there isn't anywhere for potassium to hide. This means that our soils don't have big reserves of potassium and repeated removal of clippings can get the lot.


If potassium, calcium and magnesium are boys, then sulphur is a girl.

What I mean by that is that the first three are positively charged but sulphur is present as sulphate and is negatively charged. It is only mildly interested in holding hands with the soil particles but will do so if they are attractive enough. That means that the soil particles have to have some positive charge.

Usually this means that the soils have to have some iron oxides in them so they look reddish and they are most likely to have positive charge if they are acid. This is not a good description of the Perth sands and on these, sulphate is another one that gets lost in winter rain or if too much water is applied.

Sulphur has a few other tricks. It is an important component of soil organic matter and there are therefore some analogies with nitrogen. But such analogies are never straight forward. The trick with sulphur is that a fair bit is present as organic sulphates - that is, the S is linked to the C of organic matter through an O atom - unlike N which mostly has C-N bonds. The effect of this is that the sulphate can be broken off more easily, for example just by our summer heating. So as we go through our seasons we find highest sulphate levels in soil in autumn. They drop as winter rains wash through the soil and rise again through the next summer.


Boron is a bit of a surfy.

It comes to us naturally courtesy of sea spray. So you find boron deficiency a long way from the sea. (As is also the case for Iodine deficiency.)

What is more boron is a bit flighty - not very keen on strong attachments in the soil. If the rainfall is pretty high it gets washed out easily, especially on sandy soils. The southern highlands of NSW is one place that fits these criteria and boron deficiency is well known in pine plantations.

But it is not an element that worries us much in native gardens. In plants, boron gets moved around with the water. It readily goes to the leaves which of course loose a lot of water. Deficiency occurs in parts of the plant that don't lose much water - such as bulky storage organs like cauliflowers or apple fruits. These are the plants and plant parts that are most susceptible.


Phosphorus is the king of the nutrients. Nitrogen is perhaps the queen.

This royal couple are the ones most likely to limit growth in everybody's garden. The chemistry of soil phosphate has filled soil science text books for many decades and has been treated with almost religious reverence as a great mystery. But it really is pretty simple.

Phosphate loves iron and aluminium oxides. If there are oxides of these in the soil, phosphate just can't stop holding hands with them. This means that there is only a very small amount of phosphate left in the soil solution and this is the source of P for most plants. The soil therefore rations the supply to plants.

"The more oxides in the soil, the more slowly the P gets to the plants....."

The more oxides in the soil, the more slowly the P gets to the plants, the less the chance of leaching loss, and the more P you have to apply to get an adequate rate of supply to plants.

The oxides are present in soil as very fine particles. It is the surface area of the particle that matters not the total amounts. Don't get fooled by the colour of the soil. Some of our desert soils are brilliant red. But this colour is so fiery because they don't have much organic matter to mask the colour and they have an iron oxide called haematite in them. This oxide isn't quite such a good sink for P as another oxide called goethite. (Yes, it was named for the poet!). It is a scungy yellow/brown colour and boy does it grab P.

The oxides and the P do a complicated dance to the tune of pH changes. The dance involves changes in the charge on the surface of the oxide (which we have already talked about) and changes in the form of the P present in the soil solution. This is a subject that has worried soil scientists for decades. I think it is all pretty simple - but not everyone agrees!

The really tricky thing about the oxides is that P loves them so much that it just has to bury itself in them. It can do this because the crystal structure of the oxides formed in soil is far from perfect so there are plenty of faults for the P molecules to go into. This is why P fertilizers have to be re-applied - and again we have already talked about it.

There is, however, a complicated twist to this tale. We can express it quixotically as follows. "You cannot re-apply P to the same soil".

Once you have put some P on, it becomes a different soil! The burying of P into the soil particle changes their charge and they become less able to react with more P. So the next application of P fertilizer is a bit more effective. This process becomes really important when soil is used to accept high P inputs such as sewerage. Eventually the soil becomes "saturated" because the charge becomes so negative that it can't accept any more P.


Molybdenum is the scientist's nutrient.

I say that because the need for Mo was discovered in a purely academic exercise. It was so difficult to purify solutions sufficiently to show that Mo was needed that the scientists concerned were sure that their discovery was of no practical importance. They hadn't reckoned on Australian soils!

Within a year or two it was shown that some Australian soils were so deficient that spectacular responses occurred to minute applications. Rather than the tens of kg of P needed per ha, you need a couple of hundred grams of Mo. And this is such a heavy dressing that it is sufficient for many years. Indeed you have to be careful to not walk from the +Mo plots to the -Mo plots lest your footprints show up!

Such was the response that the energy released per g on Mo (in terms of better growth) was compared to that released by uranium (shows that this was some time ago when uranium was reckoned to be a good thing). By the way, Mo's place as a super trace element has been usurped by selenium, which is needed by animals, and for which a couple of g per ha is enough.

In soil, Mo reacts pretty much like P. It too loves oxides and it too can burrow into them becoming less effective with time. The big difference is that the ions present in solution change with pH in a very different way. The result is that Mo reacts weakly with alkaline soils - and so can be made more available by applying lime. This was surely the reason for some responses to lime. In contrast, P is much less strongly affected by pH (a very brief summary of a complex question!).

The "heavy" metals...Copper, Zinc and Cobalt

We hear so much about heavy metals and toxicity that we can forget that they are essential nutrients. All plants (and animals) need copper (Cu) and zinc (Zn). Animals need cobalt (Co) and those plants that fix nitrogen from the air need it too.

But don't forget they are toxic in excess. If you have a high concentration anywhere in the soil, roots just won't go there. This is good way to keep roots out of drains but if you are trying to supply them to plants they have to be well spread.

Much of the spectacularly-flowering WA wildflowers grow on land which is deficient in both Cu and Zn. Because of these deficiencies (and P and N) the vegetation is restricted to low scrub. It was only in the late 1920s and early 1930s that it was discovered that plants needed Cu and Zn. When this knowledge was applied in WA after the second World War it lead to the clearing of huge areas and their conversion to agriculture - a process wildflower enthusiasts may view with mixed emotions.

Perhaps this means that these plants need little of these nutrients, but, in a garden it is wiser to be on the safe side and small applications of a complete fertilizer are a good insurance.

In soil, these heavy metals react with oxides and with organic matter. When they react with oxides, the process is similar to that for P and Mo but the charge on the metals is positive rather than negative.

This means that increases in pH increase the reaction because they make the oxides more negative and more attractive to the metals. Oddly, raising the pH only decreases the uptake by plants a little. This is because, although the soil holds the metals more tightly, the plants are better able to get it as the pH is increased.


I hope I haven't made this story too complex! I have tried to keep it simple - but almost every sentence could be expanded to a dozen pages, yet still be incomplete. If this has whetted your appetite, I have achieved something.

Jim Barrow is a member of the Wildflower Society of Western Australia and is a "semi-retired" CSIRO soil scientist.

This series of articles was first published in the Newsletter of the Wildflower Society of Western Australia during 1998-1999. Parts 1 and 2 of the series appeared in the March 1999 and June 1999 issues, respectively, of "Australian Plants online".

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Australian Plants online - September 1999
The Society for Growing Australian Plants