Table 6.2
Names of the Monatomic Anions
| Anion | Name | Anion | Name | Anion | Name |
| N3- | nitride | O2- | oxide | H- | hydride |
| P3- | phosphide | S2- | sulfide | F- | fluoride |
| Se2- | selenide | Cl- | chloride | ||
| Br- | bromide | ||||
| I- | iodide |
The names of monatomic cations always start with the name of the metal, sometimes followed by a Roman numeral to indicate the charge of the ion. For example, Cu+ is copper(I), and Cu2+ is copper(II).
Table 6.3
Common Polyatomic Ions
| Ion | Name | Ion | Name | Ion | Name |
| NH4+ | ammonium | 3- PO4 | phosphate | 2- SO4 | sulfate |
| OH- | hydroxide | NO3– | nitrate | C2H3O2– | acetate |
| 2- CO3 | carbonate |
Some polyatomic anions are formed by the attachment of one or more hydrogen atoms. In fact, it is common for hydrogen atoms to be transferred from one ion or molecule to another ion or molecule. When this happens, the hydrogen atom is usually transferred without its electron, as H+. If an anion has a charge of -2 or -3, it can gain one or two H+ ions and still retain a negative charge. For example, carbonate, CO 2-, can gain an H+ ion to form HCO3-, which is found in baking soda. The sulfide ion, S2-, can gain one H+ ion to form HS-. Phosphate, PO 3-, can gain one H+ ion and form HPO 2-, or it can gain two H+ ions to form H PO -. Both HPO 2- and H PO – are found in flame retardants. These polyatomic ions are named with the word hydrogen in front of the name of the anion if there is one H+ ion attached and dihydrogen in front of the name of the anion if two H+ ions are attached.
Arrhenius recognized that when ionic compounds dissolve, they form ions in solution. (For example, when sodium chloride dissolves, it forms sodium ions and chloride ions. He postulated that acids dissolve in a similar way to form H+ ions and some kind
of anion. For example, he predicted that when HCl is added to water, H+ ions and Cl- ions form. We now know that H+ ions do not persist in water; they combine with water molecules to form hydronium ions, H3O+. Therefore, according to the modern
form of the Arrhenius theory, an acid is a substance that produces hydronium ions, H3O+, when it is added to water. On the basis of this definition, an acidic solution is a solution with a significant concentration of H3O+.
To get an understanding of how hydronium ions are formed when Arrhenius acids are added to water, let’s consider the dissolving of gaseous hydrogen chloride, HCl( g ), in water. The solution that forms is called hydrochloric acid. When HCl molecules dissolve in water, a chemical change takes place in which water molecules pull hydrogen atoms away from HCl molecules. In each case, the hydrogen atom is transferred without its electron, that is, as an H ion, and because most uncharged hydrogen atoms contain only one proton and one electron, most hydrogen atoms without their electrons are just protons. For this reason, the hydrogen ion, H, is often called a proton. We say that the HCl donates a proton, H, to water, forming hydronium ion, H3O, and chloride ion, Cl (Figure 6.1).
Because HCl produces hydronium ions when added to water, it is an acid according to the Arrhenius definition of acids. Once the chloride ion and the hydronium ion are formed, the negatively charged oxygen atoms of the water molecules surround the hydronium ion, and the positively charged hydrogen atoms of the water molecules surround the chloride ion. Figure 6.2 shows how you can picture this solution.
Types of Arrhenius Acids
In terms of chemical structure, Arrhenius acids can be divided into several different subcategories. We will look at three of them here: binary acids, oxyacids, and organic acids. The binary acids are HF(aq), HCl(aq), HBr(aq), and HI(aq); all have the general formula of HX(aq), where X is one of the first four halogens. The formulas for the binary acids will be followed by (aq) in this text to show that they are dissolved in water. The most common binary acid is hydrochloric acid, HCl(aq).
Oxyacids (often called oxoacids) are molecular substances that have the general formula HaXbOc. In other words, they contain hydrogen, oxygen, and one other element represented by X; the a, b, and c represent subscripts. The most common oxyacids in the chemical laboratory are nitric acid, HNO3, and sulfuric acid, H2SO4. Acetic acid, the acid responsible for the properties of vinegar, contains hydrogen, oxygen, and carbon and therefore fits the criteria for classification as an oxyacid, but it is more commonly described as an organic (or carbon-based) acid. It can also be called a carboxylic acid.
Acids can have more than one acidic hydrogen. If each molecule of an acid can donate one hydrogen ion, the acid is called a monoprotic acid. If each molecule can donate two or more hydrogen ions, the acid is a polyprotic acid. A diprotic acid, such as sulfuric acid, H2SO4, has two acidic hydrogen atoms. Some acids, such as phosphoric acid, H3PO4, are triprotic acids. Most of the phosphoric acid produced by the chemical industry is used to make fertilizers and detergents, but it is also used to make pharmaceuticals, to refine sugar, and in water treatment. The tartness of some foods and beverages comes from acidifying them by adding phosphoric acid. The space- filling model in Figure 6.4 shows the three acidic hydrogen atoms of phosphoric acid.
Acid Review
according to the modern form of the Arrhenius theory, an acid is a substance that produces hydronium ions, H3O+, when it is added to water, and an acidic solution is a solution with a significant concentration of H3O+. Acids can be binary acids—such as HF(aq), HCl(aq), HBr(aq), and HI(aq)—oxyacids, which have the general formula HaXbOc, and organic acids, such as acetic acid, HC2H3O2.
An acid, such as hydrofluoric acid, HF(aq), whose molecules can each donate one proton, H+, to a water molecule is called a monoprotic acid. The acids, such as sulfuric acid, H2SO4, that can donate two protons are called diprotic, and some acids, such as phosphoric acid, H3PO4, are triprotic acids.
A strong acid, such as hydrochloric acid, HCl(aq), is a substance that undergoes a completion reaction with water such that each acid particle reacts to form a hydronium ion, H3O+. Thus strong acids form nearly one H3O+ ion in solution for each acid molecule dissolved in water.
Sulfuric acid, H2SO4, is a strong diprotic acid. When added to water, each H2SO4 molecule loses its first hydrogen ion completely.
H2SO4(aq) + H2O(l ) ® H3O+(aq) + HSO4-(aq)
The hydrogen sulfate ion, HSO4– that forms is a weak acid. It reacts with water in a reversible reaction to form a hydronium ion and a sulfate ion.
HSO4-(aq) + H2O(l ) H3O+(aq) + SO 2-(aq)
A weak acid is a substance that is incompletely ionized in water because of the reversibility of its reaction with water that forms hydronium ion, H3O+. Weak acids yield significantly less than one H3O+ ion in solution for each acid molecule dissolved in water.
When ammonia, NH3, dissolves in water, some hydrogen ions, H+, are transferred from water molecules to ammonia molecules, NH3, producing ammonium ions, NH4+, and hydroxide ions, OH-. The reaction is reversible, so when an ammonium ion and a hydroxide ion meet in solution, the H+ ion can be passed back to the OH- to reform an NH3 molecule and a water molecule (Figure 8.1).
Ammonia is an Arrhenius base because it produces OH- ions when added to water. Because the reaction is reversible, however, only some ammonia molecules have acquired protons (creating OH-) at any given time, so an ammonia solution contains fewer hydroxide ions than would be found in a solution made using an equivalent amount of a strong base. Therefore, we classify ammonia as a weak base, which is a base that produces fewer hydroxide ions in water solution than there are particles of base dissolved.
Combination
The reaction between the strong acid nitric acid and the strong base sodium hydroxide is our first example. Figure 8.5 shows the behavior of nitric acid in solution. As a strong acid, virtually every HNO3 molecule donates an H+ ion to water to form a hydronium ion, H3O+, and a nitrate ion, NO3-. Because the reaction goes essentially to completion, you can picture the solution as containing H2O, NO3-, and H3O+, with no HNO3 remaining. The negatively charged oxygen ends of the water molecules surround the positive hydronium ions, and the positively charged hydrogen ends of water molecules surround the nitrate ions.
Like a water solution of any ionic compound, a solution of sodium hydroxide (NaOH) consists of ions separated and surrounded by water molecules. At the instant that the solution of sodium hydroxide is added to the aqueous nitric acid, there are four different ions in solution surrounded by water molecules: H3O+, NO3-, Na+, and OH-
The ions in solution move in a random way, like any particles in a liquid, so they will constantly collide with other ions. When two cations or two anions collide, they repel each other and move apart. When a hydronium ion and a nitrate ion collide, it
is possible that the H3O+ ion will return an H+ ion to the NO3– ion, but nitrate ions
are stable enough in water to make this unlikely. When a sodium ion collides with a hydroxide ion, they may stay together for a short time, but their attraction is too
weak and water molecules collide with them and push them apart. When hydronium ions and hydroxide ions collide, however, they react to form water (Figure 8.7), so more water molecules are shown in Figure 8.8 than in Figure 8.6.
The sodium and nitrate ions are unchanged in the reaction. They were separate and surrounded by water molecules at the beginning of the reaction, and they are still separate and surrounded by water molecules after the reaction. They were important in delivering the hydroxide and hydronium ions to solution, but they did not actively participate in the reaction. In other words, they are spectator ions, so they are left
out of the net ionic chemical equation. The net ionic equation for the reaction is therefore
H3O+(aq) + OH-(aq) ® 2H2O(l )
Most chemists are in the habit of describing reactions such as this one in terms of H+ rather than H3O+, even though hydrogen ions do not exist in a water solution in the same sense that sodium ions do. When an acid loses a hydrogen atom as H+, the proton immediately forms a covalent bond to some other atom. In water, it forms a covalent bond to a water molecule to produce the hydronium ion. Although H3O+ is a better description of what is found in acid solutions, it is still convenient and conventional to write H+ in equations instead. You can think of H+ as a shorthand notation for H3O+. Therefore, the following net ionic equation is a common way to describe the net ionic equation above.
H+(aq) + OH-(aq) ® H2O(l )
pH
The scientific term pH has crept into our everyday language. Advertisements encourage us to choose products that are “pH balanced,” while environmentalists point to the lower pH of rain in certain parts of the country as a cause of ecological damage (Figure 8.3). The term was originated by chemists to describe the acidic and basic strengths of solutions.
We know that an Arrhenius acid donates H+ ions to water to create H3O+ ions. The resulting solution is called an acidic solution. We also know that when you add a certain amount of a strong acid to one sample of water—say the water’s volume is a liter—and add the same amount of a weak acid to another sample of water whose volume is also a liter, the strong acid generates more H3O+ ions in solution. Because the concentration of H3O+ ions in the strong acid solution is higher (there are more H3O+ ions per liter of solution), we say it is more acidic than the weak acid solution. A solution can also be made more acidic by the addition of more acid (while the amount of water remains the same). The pH scale can be used to describe the relative acidity of solutions.
If you take other chemistry courses, you will probably learn how pH is defined and how the pH values of solutions are determined. For now, all you need to remember is that acidic solutions have pH values less than 7, and that the more acidic a solution is, the lower its pH. A change of one pH unit reflects a ten‑fold change in H3O+ ion concentration. For example, a solution with a pH of 5 has ten times the concentration of H3O+ ions as a solution with a pH of 6. The pH of some common solutions are listed in Figure 8.4. Note that gastric juice in our stomach has a pH of about 1.4, and orange juice has a pH of about 2.8. Thus gastric juice is more than ten times more concentrated in H3O+ ions than orange juice.
The pH scale is also used to describe basic solutions, which are formed when an Arrhenius base is added to water, generating OH- ions. When you add a certain amount of a strong base to one sample of water—again, let’s say a liter—and add the same amount of a weak base to another sample of water whose volume is the same, the strong base generates more OH- ions in solution. Because the concentration of OH- ions in the strong base solution is higher (there are more OH- ions per liter of solution), we say it is more basic than the weak base solution. A solution can also be made more basic by the addition of more base while the amount of water is held constant.
Basic solutions have pH values greater than 7, and the more basic the solution is, the higher its pH. A change of one pH unit reflects a ten‑fold change in OH- ion concentration. For example, a solution with a pH of 12 has ten times the concentration of OH- ions as does a solution with a pH of 11. The pH difference of about 4 between household ammonia solutions (pH about 11.9) and seawater (pH about 7.9) shows that household ammonia has about ten thousand (104) times the hydroxide ion concentration of seawater.
In nature, water contains dissolved substances that make it slightly acidic, but pure water is neutral and has a pH of 7 (Figure 8.4).
Active and exchangeable acidity
The pH of a soil is a measure of the hydrogen ion concentration in the soil solution. pH is a the negative logarithm of H+ concentration in moles / liter:
pH = – log [H+]
and is therefore a solution measurement which only reflects the presence of acid cations adsorbed on soil colloids. A pH scale is shown below along with some reference points.
Hydrogen ion concentration is acidic soils is largely determined by the number of hydrogen ions that disassociate from the cation exchange complex. Dissociation of hydrogen is directly related to the fraction of the exchange complex that is occupied by hydrogen and aluminum ions. pH decreases (or acidity increases) as percentage saturation of H+ and Al3+ increases. Hydrogen ion in soil solution is termed active acidity and is the acidity measured by common pH tests. Hydrogen and aluminum ions adsorbed on soil colloids are termed exchangeable (or sometimes reserve) acidity. Exchangeable acidity is much larger than active acidity.
Soil pH and Salt Concentration
Acidic cations on soil colloids will exchange with cations in the soil solution. The amount of exchange is proportional to the concentration of all cations in solution, since equilibrium conditions exist. Consequently, pH of a soil solution decreases as the concentration of neutral salts (eg. NaCl, CaSO4, etc.) increases.
This phenomenon has considerable influence on measurements of pH. Measurements of pH in a soil that has been dried will be lower than those measured in the same soil when wet. Measurements of soil pH in water will be higher in situ. Further fertilizer salts will lower pH measurements.
Several methodologies have been proposed for measuring pH. Measurement in distilled water is common, but its limitations in replicating field conditions must be recognized. Measurement in 0.01 M CaCl2 has advantages in that it replicates “typical” soil solution concentrations at “average moisture contents”.
———————
Single nutrient (“straight”) fertilizers
The main nitrogen-based straight fertilizer is ammonia or its solutions. Ammonium nitrate (NH4NO3) is also widely used. Urea is another popular source of nitrogen, having the advantage that it is solid and non-explosive, unlike ammonia and ammonium nitrate, respectively. A few percent of the nitrogen fertilizer market (4% in 2007)[23] has been met by calcium ammonium nitrate (Ca(NO3)2 • NH4 • 10H2O).
The main straight phosphate fertilizers are the superphosphates. “Single superphosphate” (SSP) consists of 14–18% P2O5, again in the form of Ca(H2PO4)2, but also phosphogypsum (CaSO4 • 2H2O). Triple superphosphate (TSP) typically consists of 44–48% of P2O5 and no gypsum. A mixture of single superphosphate and triple superphosphate is called double superphosphate. More than 90% of a typical superphosphate fertilizer is water-soluble.
The main potassium-based straight fertilizer is muriate of potash (MOP). Muriate of potash consists of 95–99% KCl, and is typically available as 0-0-60 or 0-0-62 fertilizer.
Multinutrient fertilizers
These fertilizers are common. They consist of two or more nutrient components.
Binary (NP, NK, PK) fertilizers
Major two-component fertilizers provide both nitrogen and phosphorus to the plants. These are called NP fertilizers. The main NP fertilizers are monoammonium phosphate (MAP) and diammonium phosphate (DAP). The active ingredient in MAP is NH4H2PO4. The active ingredient in DAP is (NH4)2HPO4. About 85% of MAP and DAP fertilizers are soluble in water.
NPK fertilizers are three-component fertilizers providing nitrogen, phosphorus, and potassium. There exist two types of NPK fertilizers: compound and blends. Compound NPK fertilizers contain chemically bound ingredients, while blended NPK fertilizers are physical mixtures of single nutrient components.
———————
Soil pH
| Denomination | pH range |
| Strongly acidic | 5.1–5.5 |
| Moderately acidic | 5.6–6.0 |
| Slightly acidic | 6.1–6.5 |
| Neutral | 6.6–7.3 |
| Slightly alkaline | 7.4–7.8 |
| Moderately alkaline | 7.9–8.4 |
| Strongly alkaline | 8.5–9.0 |
Sources of Soil acidity
- Root respiration and decomposition of organic matter by microorganisms releases CO2 which increases the carbonic acid (H2CO 3) concentration and subsequent leaching.
- Plant growth: Plants take up nutrients in the form of ions (e.g. NO− 3, NH+ 4, Ca2+ , H 2PO− 4), and they often take up more cations than anions. However plants must maintain a neutral charge in their roots. In order to compensate for the extra positive charge, they will release H+
ions from the root. Some plants also exude organic acids into the soil to acidify the zone around their roots to help solubilize metal nutrients that are insoluble at neutral pH, such as iron (Fe). - Fertilizer use: Ammonium (NH+4) fertilizers react in the soil by the process of nitrification to form nitrate (NO− 3), and in the process release H+ ions. Acid rain: The burning of fossil fuels releases oxides of sulfur and nitrogen into the atmosphere. These react with water in the atmosphere to form sulfuric and nitric acid in rain.
- Oxidative weathering: Oxidation of some primary minerals, especially sulfides and those containing Fe2+ , generate acidity.
Sources of Soil alkalinity
Total soil alkalinity increases with:[13][14]
- Weathering of silicate, aluminosilicate and carbonate minerals containing Na+ , Ca2+ , Mg2+ and K+ ;
- Addition of silicate, aluminosilicate and carbonate minerals to soils; this may happen by deposition of material eroded elsewhere by wind or water, or by mixing of the soil with less weathered material (such as the addition of limestone to acid soils);
- Addition of water containing dissolved bicarbonates (as occurs when irrigating with high-bicarbonate waters).
The accumulation of alkalinity in a soil (as carbonates and bicarbonates of Na, K, Ca and Mg) occurs when there is insufficient water flowing through the soils to leach soluble salts. This may be due to arid conditions, or poor internal soil drainage; in these situations most of the water that enters the soil is transpired (taken up by plants) or evaporates, rather than flowing through the soil.[13]
The soil pH usually increases when the total alkalinity increases, but the balance of the added cations also has a marked effect on the soil pH. For example, increasing the amount of sodium in an alkaline soil tends to induce dissolution of calcium carbonate, which increases the pH. Calcareous soils may vary in pH from 7.0 to 9.5, depending on the degree to which Ca2+ or Na+ dominate the soluble cations.[13]
—
Treatment of High pH Soil
Fertilizers and chelates can be added to soil to increase concentrations of plant nutrients. It is important to note that addition of phosphate fertilizer alone will further reduce the availability of other nutrients.
Lowering the pH of alkaline soils, or acidifying the soil, is an option. Elemental sulfur can be added to soil as it forms sulfuric acid when it reacts with water and oxygen in the presence of sulfur-oxidizing bacteria. Iron and aluminum compounds can be added to soil, as they cause the release of hydrogen when they react with water. Sulfuric acid may also be added directly.
Additions of appreciable amounts of organic matter will help to acidify the soil as microbes decompose the material, releasing CO2 which then forms carbonic acid. Organic acids are also released during humus decomposition. Peat and peat moss are highly acidic forms of organic matter but can be costly.
Application of acidifying fertilizers, such as ammonium sulfate [(NH4)2SO4,], can help lower soil pH. Ammonium is nitrified by soil bacteria into nitrate and hydrogen ions.
Soils naturally containing carbonates, or lime [Calcium Carbonate CaCO3.] , are very difficult to acidify, and it may take years before a significant change in soil pH is seen. Even then, the carbonatic parent material will continue to weather, producing more soluble carbonate and buffering the soil solution pH.
Many plants can tolerate pH values between 7 and 8, and some actually thrive at these higher pH values. Choosing plants that grow well in mildly alkaline soils can be selected. This is the most reasonable “treatment” option for soils that have developed from carbonatic parent material.
Vegetable garden plants such as asparagus, beets, cabbage, cauliflower, celery, carrots, lettuce, parsley and spinach grow well in soils whose pH is between 7 and 8.
—————-
A chemical change or chemical reaction is a process in which one or more pure substances are converted into one or more different pure substances. Chemical changes lead to the formation of substances that help grow our food, make our lives more productive, cure our heartburn, and much, much more. For example, nitric acid, HNO3, which is used to make fertilizers and explosives, is formed in the chemical reaction of the gases ammonia, NH3, and oxygen, O2. Silicon dioxide, SiO2, reacts with carbon, C, at high temperature to yield silicon, Si—which can be used to make computers—and carbon monoxide, CO. An antacid tablet might contain calcium carbonate, CaCO3, which combines with the hydrochloric acid in your stomach to yield calcium chloride, CaCl2, water, and carbon dioxide. The chemical equations for these three chemical reactions are below.
Once you know how to read these chemical equations, they will tell you many details about the reactions that take place.
Chemical changes lead to the formation of substances that help grow our food, make our lives more productive, and cure our heartburn.
Interpreting a Chemical Equation
In chemical reactions, atoms are rearranged and regrouped through the breaking and making of chemical bonds. For example, when hydrogen gas, H2(g), is burned in the presence of gaseous oxygen, O2(g), a new substance, liquid water, H2O(l), forms. The covalent bonds within the H2 molecules and O2 molecules break, and new covalent bonds form between oxygen atoms and hydrogen atoms (Figure 7.1).
chemical equation is a shorthand description of a chemical reaction. The following equation describes the burning of hydrogen gas to form liquid water.
2H2(g) + O2(g) → 2H2O(l)
Chemical equations give the following information about chemical reactions.
- Chemical equations show the formulas for the substances that take part in the reaction. The formulas on the left side of the arrow represent the reactants, the substances that change in the reaction. The formulas on the right side of the arrow represent the products, the substances that are formed in the reaction. If there are more than one reactant or more than one product, they are separated by plus signs. The arrow separating the reactants from the products can be read as “goes to” or “yields” or “produces.”
- The physical states of the reactants and products are provided in the equation. A (g) following a formula tells us the substance is a gas. Solids are described with (s). Liquids are described with (l). When a substance is dissolved in water, it is described with (aq) for aqueous, which means “mixed with water.”
- The relative numbers of particles of each reactant and product are indicated by numbers placed in front of the formulas. These numbers are called coefficients. An equation containing correct coefficients is called a balanced equation. For example, the 2’s in front of H2 and H2O in the equation we saw above are coefficients. If a formula in a balanced equation has no stated coefficient, its coefficient is understood to be 1, as is the case for oxygen in the equation above (Figure 7.2).
- If special conditions are necessary for a reaction to take place, they are often specified above the arrow. Some examples of special conditions are electric current, high temperature, high pressure, and light.
Cations
Soil fertility testing is a valuable tool to make informed nutrient management decisions. Three main philosophies exist to interpret soil test reports and provide appropriate fertilizer recommendations. Those include (i) sufficiency level of available nutrients (SLAN), (ii) buildup and maintenance, and (iii) basic cation saturation ratio (BCSR) concepts (Black, 1993). The SLAN approach works on the principle that there are certain critical levels of individual nutrients. If a soil tests above the critical level, the crop will not likely respond to fertilization but if the soil tests below the critical level, the crop will respond to fertilization (Eckert, 1987). The buildup-and-maintenance approach calls for a gradual buildup of soil nutrient levels above the critical levels over time, and then to maintain these levels by replacing the amounts of each nutrient removed by the crop at harvest (Olson et al., 1987; Black, 1993; Voss, 1998). The focus here is to always maintain the soil fertility status at a high level with constant fertilizer applications, so that yields are sustained. The third approach is the BCSR, which recommends an optimum or ideal calcium (Ca), magnesium (Mg), and potassium (K) saturation ratio on the soil exchange complex to achieve maximum crop yields (McLean, 1977; Voss, 1998). With BCSR, fertilizer recommendations are made to adjust the cation saturation ratios to an optimum or ideal level, irrespective of actual soil nutrient levels. A soil with such an optimal saturation ratio of base cations is considered to be a balanced soil and the practice of adding fertilizer/amendments to achieve a desired ratio is called “soil balancing.” Generally, the SLAN approach recommends fertilizing based on plant needs, the buildup-and-maintenance approach focuses on fertilizing the soil, and BCSR targets on countering the mineral imbalances in soil (Black, 1993; Eckert, 1987). In contrast to sufficiency level and buildup-and-maintenance philosophies, BCSR focuses only on Ca, Mg, and K and does not directly relate to the availability of nitrogen (N), phosphorus (P), sulfur (S), or micronutrients (Eckert, 1987).
There is an apparent discrepancy in using different philosophies for soil fertilizer recommendation programs. However, over the past three decades, ongoing research by soil fertility scientists on soil testing, interpretation, and calibration has resulted in the adoption of SLAN and buildup-and-maintenance philosophies or a hybrid of these two, as a standard fertilizer recommendation practice by land-grant universities. On the other hand, some commercial soil-testing laboratories employ BCSR and buildup-and-maintenance approaches in their lime and fertilizer recommendation programs (Voss, 1998). To date, there is little published research that substantiates the BCSR theory and the concept of a balanced soil for maximizing yields, and only a few studies have tested its efficacy. Still, some agronomists, consultants, commercial soil-testing labs, and farmers strongly subscribe to this practice and continue to use it to guide their soil management decisions and nutrient recommendations. This review (i) presents a brief history of the BCSR theory, (ii) provides an overview of the research that has been conducted on BCSR and crop production, and (iii) identifies knowledge gaps that need attention.
What is basic cation saturation?
Soil nutrients can exist in many forms, but the nutrients that plants take up are mostly positively (cations) or negatively (anions) charged ions. Cations and anions are available in soil solution or on the soil exchange sites. Those that are in soil solution are readily available to plants. Cation exchange sites are negatively charged surfaces of clay and organic matter that attract and hold cations in the soil. Soil tests measure the sum of these exchange sites and report them as cation exchange capacity (CEC). The CEC is a defining feature of soils, and the greater the CEC of a soil, the more cations it can hold (Hazelton and Murphy, 2007). Cations are generally classified as “basic” and “acidic” based on their influence on soil pH through various soil reactions. Basic cations (also called as nonacidic cations) include calcium (Ca2+), magnesium (Mg2+), potassium (K+), and sodium (Na+). Acidic cations consist of hydrogen (H+) and aluminum (Al3+). Figure 1 shows different basic and acidic cations on soil exchange sites. Base saturation indicates the proportion of these basic cations that occupy the soil exchange sites (CEC). In other words, if a soil has a 50% base saturation of Ca, then Ca occupies 50% of the exchange sites. Figure 2 gives a pictorial representation of the base saturations of Ca2+ at approximately 60%, Mg2+ at approximately 15%, and K+ at approximately 5%, making up to approximately 80% of the total CEC of the soil. Remaining sites (approximately 20% of CEC) could be occupied by other basic cations, such as Na+, or acidic cations such as H+ and Al3+. However, BCSR is primarily concerned with the percent saturation of only Ca2+, Mg2+, and K+ ions on the exchange sites.
Refrences
Chapter 6: More on Chemical Compounds
Chapter 7: An Introduction to Chemical Reactions
Chapter 8: Acids, Bases, and Acid-Base Reactions
Lab 9 – Soil pH and Soil Testing – Crop and Soil Science
Historical Perspective of Soil Balancing Theory and Identifying Knowledge Gaps: A Review