Fertilizers

Achievement Standard here.
These notes wil be undergoing some further revision but should now be fairly concordant with the Achievement standard.
Pardon my American spelling of fertilizer. If someone really wants to go through and fix it all, just be careful not to break links etc.
Video - cows and nitrates (from Hot Science website)
Note: Wikispaces have done something to their interpreter which creates an extra line break after superscript or subscripts which has done funny things to this page. I will go through and fix it when I have time. If anyone knows what the problem is, I would appreciate their contacting me. At the moment I am just removing the line breaks, which undoes the subscript, so chemical formulae don't display properly unless I do them as graphics.
Query for teachers: I am assuming a discussion of pKa is not required. Any other opinions?

1. Plant nutrients


tree2.pngFertilizers are sometimes termed 'plant food', but plant food is not “food” in the sense of supplying energy to the plant. Plants make their own food from sunlight, carbon dioxide and water (i.e. the elements carbon, hydrogen and oxygen).
However, plants require additional elements to grow and function. They acquire these through their roots along with the water they uptake. The elements are divided into two groups:
*
  • Macronutrients: elements which make up the bulk of the plant’s mineral requirements and
  • Micronutrients, which are: required in smaller quantities These are also known as trace elements.

The three main macronutrients are nitrogen, phosphorus and potassium, sometimes known as “NPK”.. However, sulfur and magnesium are also sometimes considered macronutrients, and fall between macro- and micronutrients in terms of what plants require
Micronutrients are also known as trace elements. They are used for particular plant functions, such as in enzymes. Examples of micronutrients include iron, copper, selenium and cobalt. Usually the soil has enough of these, but in some places they are lacking e.g. the Volcanic Plateau is usually short of cobalt.
The prescription specifies N, P, K, S, Mg, Co, Mo and B as being the elements about which specific knowledge is required. A few other important elements will be mentioned here.

How are these elements present?

The macro and micronutrients are not present as elements. You don’t put calcium metal on your plants. They are present in compounds, combined with other elements. These can be ionic or covalent. This will be discussed in more detail later.

What do plants need nutrients for?

Most plant material is carbohydrates such as cellulose, lignin and starch. These are made from carbon, hydrogen and oxygen which plants get entirely from carbon dioxide and water. There are other materials in plants which contain elements from plant nutrients.These include protein (found in all cells), amino acids, enzymes, chlorophyll, DNA and ATP (an energy carrier). Plants use the nutrients in some of the folllowing ways:
Macronutrients:
Nitrogen is found in all proteins, amino acids, nucleic acids including DNA and enzymes.
Phosphorus is used in DNA and is also found in an important molecule called ATP (adenosine triphosphate), which all living things use for respiration..
Potassium: plants and animals rely on charged ions to move electrical charges across cell membranes and create electric 'action potential'. Animals use a combination of sodium and potassium ions, but plants use potassium and chloride ions. We say that potassium ions are an electrtolyte in plants.
Secondary nutrients
Sulfur: In plants the amino acids cysteine and methionine contain sulfur, as do all polypeptides, proteins, and enzymes which contain these amino acids.
Magnesium: Magnesium ions are essential to the basic nucleic acid chemistry of life, and thus are essential to all cells of all known living organisms. Plants have an additional use for magnesium in that chlorophylls are magnesium-centered porphyrins. Many enzymes require the presence of magnesium ions for their catalytic action, especially enzymes utilizing ATP, or those which use other nucleotides to synthesize DNA and RNA. Magnesium deficiency in plants causes late-season yellowing between leaf veins, especially in older leaves, and can be corrected by applying epsom salts (which is rapidly leached), or else crushed dolomitic limestone to the soil.
Trace elements/Micronutrients
Cobalt: in small amounts is essential to many living where it has a role in enzymes..
Molybdenum: has a role in nitrogen uptake, and is particularly important in legumes (such as clover) that use nitrogen fixing bacteria.
Boron:is primarily for maintaining the integrity of cell walls

Where these other elements come from: In a natural environment (no fertiliser) they are obtained from the weathering of minerals in the soil, from recycling of material by decay (e.g. dead plants and animals and wastes such as urine), from natural processes (lightning producing nitrogen oxides) and the action of microorganisms (nitrogen-fixing bacteria). These can be sufficient to sustain a natural ecosystem, but when crops are continuously removed in agriculture these natural processes have difficulty replacing the elements lost with the harvest.
Nitrogen cycle. Click for source.
Nitrogen cycle. Click for source.
Nutrient cycling is particularly important for nitrogen, as this is never produced by the breakdown of rock. Nitrogen is found n the atmosphere in the form of N
2
, a gas. Nitrogen gas is not particularly reactive, but small amounts are formed into nitrate ions by reaction with oxygen and water during lightning. More is converted by nitrogen fixing bacteria (left of bacteria). However, recycling by decomposers is really important. Some nitrogen is returned to the atmosphere by denitryfying bacteria. These bacteria are really useful for removing nitrogen from waste (e.g. sewage) to prevent overloading waterways such as lakes with nitrogen, which could lead to eutrophication (see later section).



2. Chemistry of solutes present in the soil

For plants to access elements in the soil they must be dissolved i.e. present as solutes. Generally they are present as ions.

Ionic and covalent bonds

Although we talk about a soil needing “potassium”, it is not potassium metal which is present. Potassium metal is highly reactive, and when it reacts with water:
Potassium + water --> potassium hydroxide + hydrogen gas
potassiumreaction.png

The potassium atoms have given their electrons to the hydrogen ions to make potassium ions and hydrogen atoms. Hydrogen atoms are only stable if two of them stick together to make a hydrogen molecule. When they do this, they share their electrons with what is called a covalent bond. The hydrogen molecule formed in the chemical reaction shown above is held together by a covalent bond. However, few inorganic fertilizers involve this type of bonding to any great degree and the AS concentrates on ionic bonding in questions
external image Sodium_chloride_crystal.png
Ionic solids: Ionic substances form solids which are held together by their electrostatic charge in a regular arrangement as shown in the diagram. For example, a crystal of potassium chloride) will consist of regularly arrahged potassium ions (blue in diagram) and chloride ions (green in diagram)

Dissolving ions: When ionic substances, such as potassium hydroxide or sodium chloride, dissolve in water, the ions come apart (or dissociate).They can do this because water has a special property of being able to cancel out the electric force fields that hold them together in solid form.


water_dipole.png
These dipoles arrange themselves around the ions as below. This cancels out the electric force field holding the ions together, and allows them to move apart, or dissociate. This causes them to dissolve.

water_around_ions.png



A solution of potassium chloride therefore contains potassium ions and chloride ions.
Not all ionic substances are soluble
;Some ionic substances are not soluble or only weakly soluble, so do not completely dissociate. Calcium phosphate (rock phosphate) is ionic, but insoluble. It must be chemically changed into calcium dihydrogen phosphate (which is soluble) before plants can use the phophorus.
How plants absorb fertilisers: Plants take up their nutrients in the form of these dissolved ions: The ions can pass into the roots by osmosis, along with the water. They then pass up into the plant through vessels called xylem. Because water is lost from the leaves by transpiration, this ensures continuous transport of water and dissolved minerals from the roots up to the stems and leaves and other parts of the plant where they are needed.
Covalent substances: Some covalent substances can dissolve. This is because of something called hydrogen bonding, which is caused by uneven electric charges in the covalent bond. These uneven charges are attracted to the uneven charges in water (water is termed a polar substance), allowing the substance to dissolve. An example of a substance that does this is is urea (link has more information).
When I get time I will create an illustration of this to put up.

Ions for plants
These are the ionic formula for the elements mentioned in the Achievement Standard.
ionic_formulae_fertilizers.png
There are other ions which can be important in fertilizers and for plants. Click here for a more complete list.
Ion charges: Ion charge indicates how many extra or fewer electrons the ion has. The nitrate ion has one extra electron (one more electron than the number of protons in the whole ion). Positively charged ions have lost electrons.
You are likely to be asked for the charges on some ions from this list, and what this means in terms of electrons, in the exam.

3. Fertiliser composition

One of the things you will need to be able to do is give the composition of various fertilizers. This can be weight percent or NPK value.

3.1 Weight Percent

This is the percentage by weight of the elements that make up the compound. To work it out, you neet do know the relative mass of the elements that make up the compound. You will always be given this.
Example: what weight percent of nitrogen is potassium nitrate KNO3 (apologies for strange notation; it is hard to use equals sign in wiki notes because it thinks it means what you are writing is a heading)
Information: Mr : K is 40; N is 14; O = 16
You will be given this information in the examination; Mr means relative mass
Steps to working out the answer:
Step 1. Calculate Mr of the whole molecule:
working out: add up the mass numbers 40 + 14 + (3 x 16) which is 102.
Step 2: Divide the amount of nitrogen into the total and convert to a %
14 / 102 x 100 = 14 %


3.2 NPK values

Many fertilisers have information on the packet such as “NPK 12-5-3”. These numbers are the percentages of nitrogen, phophorus and potassium respectively in the fertilizer. However, the phosphorus and postassium are reported as percenteges of their weight as an oxide i.e equivalent to P2O5
and K2O. This means that the above "12-5-3" fertiliser is 12% by weight nitrogen, 5% P2O5
and 3% K2O. For a fertiliser that contains only one chemical this can be easily calculated. Most fertilisers are mixtures.
You are only likely to be asked for the number ov grams of nitrogen on based on NPK values.
Example: A fertiliser bag give the information: NPK 24-10-6. How many grams of nitrogen will there be in 2 kg of this fertilizer?
Solution: N is 24%, so 24/100 x 2000g gives 480g.

If you are unfortuante enough to get a question about P or K, you may need to work out exactly what you are being asked and I am not sure what kind of question the NZQA is likely to throw at you. They have a habit of sticking in nasty questions when they want to get pass rates down, but the statistics on this standard suggest that is not too likely at the moment. The standard just says 'percentage composition' which is (probably deliberately) ambiguous.


4. Problems with using and applying fertilisers

4.1: Leaching:
This occurs when the soluble component of fertilizers is removed by water (e.g. after rain). They can wind up in waterways, lakes or the sea. It causes two problems
  • Less availability of nutrients to the plants under fertilization
  • Pollution of waterways. This can lead to eutrophication (next section) or the growth of undesirable plants or algae. For example, toxic algae growth in the Rotorua lakes has killed some fish and limited recreational use of the lakes. Fertiliser runoff is suspect in causing damage to the Great Barrier Reef off Queensland
The effects of leaching can be reduced by
  • Application of excess fertiliser, which reduces the first problem but greatly increases the second.
  • Applying fertilisers in times of lower rainfall – not always possible
  • Using slow release fertilisers. These can be organic, but a variety of synthetic slow release fertilisers are manufactured e.g. osmocote, or fertilisers with limited solubility and applied in pellet form

4.2 Nutrient overload in waterways
This arises the overloading of waterways with nutrients, particularly nitrogenous ones. It promotes the growth of water plants and algae. This can cause several problems
  1. Some algae are directly toxic (see above)
  2. When too much growth is too rapid, the decay of dead material removes all oxygen from the water leading to anaerobic conditions. This kills most animal life, and produces foul smells. The term for this is eutrophication.
Eutrophication does not only arise as a result of over fertilisation; animal and human waste contributes nitrogen to the water also. Solutions to the problem include
    • Modifying the land use in the catchment e.g a switch from dairy to forestry
    • Restrictions on fertiliser application and stock density
    • Use of slow release fertiliser
    • Denitrification of sewerage in a treatment plant and removal of low level treatment options such as septic tanks. Many environmental authorities require milkshed or similar waste to be sprayed on pasture and have strict limits about how much and when this can be applied.

5. Types of fertiliser
Fertilizers can be organic or inorganic.
Compost, an organic fertilizer
Compost, an organic fertilizer
Organic:
this means fertiliser made from plant or animal remains. It includes manures, compost and so on. Bird manure is more nitrogen rich than mammal manure because birds do not urinate – their nitrogen waste is passed out in their droppings along with the products of digestion. In mammals, droppings only contain undigested food.
Compost is broken down plant material. Sometimes it needs to be made less acid e.g. by adding lime (ground limestone, calcium carbonate)
It is now common for dairy farmers to spray the waste from the milking sheds (dung and urine) onto the pasture. This not only avoids nutrient overload in waterways, but also reduces the amount of nitrogenous fertiliser that needs to be added to pasture. However, the amount sprayed needs to be balanced to avoid runoff into streams and this can be a problem (e.g. during wet weather). Some farmers now use specially designed sheds to reduce this problem by using the animal waste to create a high-nitrogen fertiliser that can be stored and applied later.
Advantages of organic fertilisers:
  • Slow release of nutrients – fertilises for a long time
  • Improves soil structure (even more when used with techniques like 'no-till')
  • Less likely to cause pollution by leaching, because they tend to be 'slow release' (usually, not always)
  • can involve recycling and uses renewable resources - see comment above about spraying cowshed waste

Disadvantages of organic fertilisers:
  • Cannot easily control exactly the balance of nutrients
  • Not suited to promote ‘quick growth’ as they have to break down first
  • Not (usually) suited to hydroponics
  • Sometimes it is difficult to supply sufficient quantities of all nutrients needed to maintain productivity
  • Tends to be short on potassium
  • cowshed waste amounts can be inconvenient amounts at the wrong time - for example, the exceptionally wet winter/spring of 2010 meant it couldn't be applied to already saturated pasture; few dairy farms have sufficient effluent holding facilities to stockpile it until drier months (though this would be ideal). Excessive application in very dry conditions can 'burn' pasture if inadequately diluted (e.g. when irrigation water is short)

Note that new technologies are being developed which slow the release of nitrogen from effluent or stock waste (there is a lawsuit going on about this at the moment).

Ammonium nitrate, an inorganic fertilizer
Ammonium nitrate, an inorganic fertilizer
Inorganic fertilizers

Include minerals, and manufactured chemicals.
Advantages of inorganic fertilisers
  • Instant release of nutrients – quick growth
  • Can tailor nutrient balance to needs
  • Often less bulky, so suited to application by topdressing
  • can be made 'slow release' e.g. 'Osmocote', but this adds to expense (link is to a PDF from the company that makes this brand of 'controlled release' fertiliser).

Disadvantages of inorganic fertilisers
  • Inappropriate use leads to nutrient runoff and pollution
  • Manufacture uses non-renewable resources e.g. oil, gas, rock phosphate
  • can be expensive
  • sometimes they contain impurities which can build up in the soil with undesirable consequences

The key with many of the disadvantages is 'inappropriate use', which frequently means overuse. For example, so much fertliser is used of the putting greens of many golf courses that the soil itself could be used elsewhere as fertliser.
There is a HortResearch paper on fertilisers for horticultural use here, which includes compositions of various inorganic and organic fertilisers


6. Soil
soilhorizon.png
Diagram of soil horizons

6.1 What is soil?

Soil is a mixture of organic material, called humus, and broken down (weathered) rock.
A soil horizon goes from the soil, down to partly broken down rock called regolith, to unweathered bedrock.
During the process of chemical weathering minerals are changed. The minerals produced are mostly oxides and clays. Resistant minerals like quartz remain in the weathered material. The breakdown also releases some nutrients, notably potassium, calcium and magnesium, and phosphorus. Nitrogen must come from the breakdown of organic material, or from nitrogen-fixing bacteria, or from natural sources such as nitrogen oxides in air dissolved in rain.
Nitrogen fixing bacteria are found in the roots of legumes such as beans, peas, clover, lupin, gorse and broom.
Phosphorus comes from both weathered down rock and from recycling. All nutrients are recycled to some extent (of interest: the relative abundance of sodium ions in seawater arises partly from the fact that plants don't take up much of this; both potassium and sodium are released on weathering but plants quickly remove potassium, leaving sodium to wash into the sea)
Clay is one of the most important inorganic components of soil. Clay is a term for a group of minerals with a sheet structure. The sheets can slide over each other (which is why clays are slippery) and can hold water and ions between the sheets. This ability has important implications for the supply of nutrients to plants. Sodium rich soils are 'sticky', leading to poor drainage. Applying a source of calcium ions e.g. in superphosphate, can help make soils more friable (crumbly) by displacing sodium held between clay sheets and replacing them with calcium ions. The calcium versions of clays are less gluey in their structure.

6.2 Soil pH

Soil pH is a measure of how acid or basic a soil is. This can affect the nutrient retention and release. It also affects plant growth; with different plants having different preferences regarding soil pH. For example, citrus trees prefer more acidic soils and olive trees prefer a higher pH.
Soils with a high organic (humus) content tend to be more acid. Soils weathered from limestone or calcareous soils can have a higher pH. The pH can be increased by adding lime (ground calcium carbonate or limestone), dolomite, or burnt lime or seashells. Soil pH can be lowered (made more acid) by adding organic content or acidic fertilisers such as superphosphate or ammonium sulfate.
pH is a measure of the concentration of hydrogen ions in the soil.

About pH

I know this is complex, but it was asked about in the 2008 exam.
Although you may have been taught that acids are 'things with hydrogen in them', reality is a bit more complex.
Hydrogen ions can't exist by themselves in solution. Instead, they react with water to produce 'hydronium ions':
hydroniumformation.png
However, water itself normally forms some hydronium ions as well as hydroxide ions. In neutral water, they have the same concentration.
hydroniumequation.png
If this becomes unbalanced - the hyronium ion concentration is greater than the hydroxide ion concentration, the pH is <7 and it is acid. If it is the other way around, the hydroxide ion concentration is higher and it is basic.
Acidic fertilisers have the ability to donate hydrogen ions (protons) to water to create the hydronium ion. For example, the ammonium ions in ammonium sulfate can donate a proton to water, creating ammonia and hydronium ions:
ammoniumacid.png
The hydronium ions thus created make it acid.
If you are asked about this, you can say that hydrogen ions are created for any acid fertiliser. This should be enough to get you achieved. To see why I am going into this detail, look at this marking schedule. However, even this explanation is rather simplified and a more detailed understanding of acid-base theory requires something called Brønsted-Lowry theory. Students doing chemistry should be familiar with it.

Acidic fertilisers - lowering pH
Some fertilisers are acidic because they have been formed by the reaction of strong acids with weak bases. Examples are:
  • Epsom salts (magnesium sulfate) – formed by reaction of magnesium oxide and sulfuric acid. Used as a source of magnesium.
  • Alum (aluminium sulfate) – does not supply any nutrients, but is used as a soil additive to acidify soils
  • Ammonium sulfate: formed by reaction of ammonium hydroxide with sulfuric acid. Used as an acidic nitrogenous fertiliser.
  • Ammonium nitrate: formed by reaction of ammonia with nitric acid. Also acidic, but not sold to public because it is explosive.
  • Superphosphate (see later note)
  • Gypsum (calcium sulfate) – also used to break up clay. It replaces the sodium in clays with calcium ions, which makes them less sticky
All acidic fertilisers increase the concentration of hydrogen ions in the soil.

Neutralizing acid soil - increasing pH
Acid soils can be made more neutral (pH increased) by -
  • calcium carbonate (lime)
  • calcium hydroxide (slaked like, usually only used very sparingly but sometimes added to compost - e.g. in making mushroom compost)
  • calcium magnesium carbonate (dolomite). This has the advantage of supplying magnesium
  • tripotassium phosphate supplies potassium and phosphorus, and has a pH above 7 because it is formed by reaction of potassium hydroxide and phosphoric acid (a weak acid). However, it is common to use potassium dihydrogen phosphate which is more acidic because not all the hydrogen ions are replaced with potassium ions and that is less basic.
  • ammonia, formed by breakdown of urea, is basic. Acid soils 'hold on' to ammonia, so effectively it partly neutralizes them. High pH environments result in ammonia gas being given off.
Of these methods, lime (crushed limestone) is by far the most common way of increasing soil pH and the one you are most likely to be asked about.

Clay and ion exchange
Most soils contain some clay. Clay is a layered silicate which contains water and ions between the layers. They can have the ability to
'soak up' and later release ions from fertilisers, improving the release time of fertilisers. This is called ion exchange.
Clay minerals can have several different types of ions between the layers. Many NZ soils contain clays rich in sodium ions, which tend to be sticky and easily become waterlogged. Adding an acidic calcium source, such as calcium sulfate, to the soil can help break up the sticky clays and make the soil more crumbly (or 'friable') by replacing the sodium ions with calcium ions making the clay less sticky. This can improve aeration and drainage of the soil. Care must be taken not to make the soil too acidic.
.
7. Manufacture of some inorganic fertilisers

7.1. Superphosphate

Nauro after rock phophate removal
Nauro after rock phophate removal
This is a commonly used fertiliser in NZ. It is made locally from sulfur and rock phosphate.
Rock phosphate: this mineral is formed by the reaction of phosphate in seabird droppings with limestone in raised coral atolls. It is mined from places like Christmas Island and Nauru. Rock phosphate by itself is unsuitable as a fertilizer because of its low solubility. It is turned into supephosphate by industrial chemical reactions. Superphosphate is suitable for adding phophorus for soil and for application by aerial topdressing, which is ideal for NZ's rugged landscape.
Superphosphate makes the phosphorus content soluble; it also is suitable as an 'acid' fertilizer for lowering soil pH which may release other nutrients.
Superphosphate manufacture:
1. the rock phosphate is reacted with sulfuric acid. The sulfuric acid is also made locally, from sulfur (mostly obtained during oil refining).
To make sulfuric acid: sulfur is burned in air to form sulfur dioxide:
sulfurdioxide.png

2. The sulfur dioxide is then reacted with oxygen in the presence of a catalyst (vanadium pentoxide) to form sulfur trioxide:sulfurtrioxide.png
3. The sulfur trioxide is dissolved in water to produce sulfuric acid:

sulfuricacid.png

4. Rock phosphate is calcium phosphate. This is reacted with sulfuric acid to form a mixture of calcium hydrogen phosphate and calcium sulfate:
Ca3(PO4)2 + H2SO4 ---> Ca(H2PO4)2 + CaSO4
Calcium hydrogen phosphate is soluble in water, making the phosphate available to plants. Rock phosphate is insoluble.

7.2 Urea and ammonium sulfate
These fertlisers require the production of ammonia for their manufacture.

Manufacture of ammonia: ammonia is made from natural gas and nitrogen from the air. This is called the Haber Process
Methane from natural gas is reacted with steam to make hydrogen:using a catalyst of nickel oxide. This is called steam reforming:
steamreformingequation.png

The hydrogen is then separated out and reacted under pressure with nitrogen, using an iron oxide catalyst:
ammoniareaction.png

(click here for more detail)
Urea: urea is made by reacting ammonia and carbon dioxide (the carbon dioxide can be produced from the carbon monoxide in the hydrogen production by reacting it with oxygen)

ammonia + carbon dioxide → urea (click here for more detail)

Ammonium sulfate:
Ammonia is reacted with sulfuric acid; usually, both are dissolved in water. Ammonia dissolved in water is called ammonium hydroxide

ammonium hydroxide + sulfuric acid → ammonium sulfate + water

This is a neutralization reaction. Ammonium hydroxide has a pH of about 13 and sulfuric acid a pH of 1. When they combine, the produce a fairly neutral salt, which has a pH of about 5 (it isn't 7 because sulfuric acid is a stronger acid than ammonia is a base).

Ammonium nitrate
This is similar to ammonium sulfate, except that nitric acid rather than sulfuric is used. Ammonium nitrate can be explosive, so although it is a better source of nitrogen, extreme caution must be used for storage and transport of this fertiliser.