Inorganic and organic chemistry principles are fundamental to agricultural science and extension services in Kenya. Understanding the nature and behaviour of acids and bases enables agricultural professionals to manage soil fertility, pesticide formulation, and nutrient availability effectively. This chapter applies physical chemistry principles to practical agricultural contexts, equipping extension officers and farmers with knowledge to optimize crop production and protect environmental health.

1.1 Acid and Bases

Acids and bases play a vital role in soil chemistry, plant nutrition, and agrochemical application. Many soil amendments, fertilisers, and pest control substances involve acidic or basic compounds. Grasping their definitions, properties, classification, and preparation methods helps agricultural professionals make informed decisions about soil treatment and crop management.

1.1.1 Definition of terms

The terms acid and base are foundational in chemistry, describing substances with distinct chemical behaviours that influence agricultural environments.

Acid definition

An acid is a substance that increases the concentration of hydrogen ions (H⁺) when dissolved in water, making the solution acidic. In soil science, acids can lower soil pH, affecting nutrient solubility and microbial activity. Example: acid rain containing sulfuric acid can damage crops by altering soil chemistry.

Base definition

A base is a substance that increases the concentration of hydroxide ions (OH⁻) in an aqueous solution, resulting in alkalinity. A base that dissolves in water is called an alkali. Bases neutralise acids, raising soil pH and influencing nutrient availability. Example: lime (calcium carbonate) is applied to acidic soils to increase pH and improve crop yields.

Arrhenius and Bronsted-Lowry definitions

The Arrhenius definition states that acids produce H⁺ ions in water, and bases produce OH⁻ ions. The Bronsted-Lowry theory broadens this by defining acids as proton donors and bases as proton acceptors, so it also applies where water is not the solvent. Example: when ammonia (NH₃) from a fertiliser dissolves in soil water, it accepts a proton from water and forms the ammonium ion (NH₄⁺), so it behaves as a Bronsted-Lowry base.

pH concept

pH measures the acidity or alkalinity of a solution, calculated as the negative logarithm (base 10) of the hydrogen ion concentration in mol/L. At 25 °C, a pH of 7 is neutral, values below 7 are acidic and values above 7 are alkaline. Soil pH influences nutrient solubility; for instance, phosphorus availability decreases in highly acidic or alkaline soils, impacting fertiliser effectiveness.

Worked example: a dilute hydrochloric acid solution used to rinse a soil sample bottle has [H⁺] = 0.001 mol/L.

pH = -log[H⁺]

pH = -log(0.001)

pH = -(-3)

pH = 3

The solution is acidic, with a pH of 3.

At 25 °C, pH + pOH = 14, so a solution with pH 3 has pOH 11. Each unit of pH change represents a tenfold change in hydrogen ion concentration.

1.1.2 Characteristics of acid and bases

Identifying acids and bases by their properties is essential for safe handling and effective use in agriculture. In a laboratory or farm store, chemicals must never be tasted or touched to identify them; indicators and pH meters are used instead.

Physical properties of acids

Acids often have a sour taste and can be corrosive. They turn blue litmus paper red and conduct electricity in solution because they form ions. Example: citric acid in fruits contributes to their sourness, and concentrated sulfuric acid used in fertiliser manufacture requires careful handling because of its corrosive nature.

Physical properties of bases

Bases usually taste bitter and feel slippery because they react with the oils in skin to form a soapy layer. They turn red litmus paper blue and also conduct electricity in solution. Example: lime, a base used in soil treatment, is alkaline and can irritate skin and eyes if not handled with care.

Chemical reactivity of acids

Acids react with reactive metals to produce hydrogen gas and a salt, and with carbonates to release carbon dioxide.

Zn + 2HCl → ZnCl₂ + H₂

CaCO₃ + 2HCl → CaCl₂ + H₂O + CO₂

This reactivity is useful for removing carbonate scale from water tanks and drip irrigation lines, where a dilute acid dissolves the deposit and fizzes as carbon dioxide escapes. It also explains why acid fertilisers and sprays corrode unprotected steel equipment.

Chemical reactivity of bases

Bases react with acids to form a salt and water, a neutralisation process important in adjusting soil pH.

Ca(OH)₂ + 2HCl → CaCl₂ + 2H₂O

CaCO₃ + 2H⁺ → Ca²⁺ + H₂O + CO₂

Example: applying lime to an acidic maize field neutralises excess hydrogen ions, improving conditions for root growth and nutrient uptake. Strong alkalis such as sodium hydroxide also react with fats and oils, which is why they are used in cleaning.

Indicators of acids and bases

Indicators such as litmus, phenolphthalein, and methyl orange change colour depending on acidity or alkalinity. Extension workers use these indicators, or universal indicator paper, to test soil extracts and water pH in the field for immediate assessment.

1.1.3 Classification of acids and bases

Acids and bases are classified mainly by how completely they dissociate (ionise) in water, which determines their strength. Strength is not the same as concentration: a concentrated solution of a weak acid can still be far less reactive than a dilute solution of a strong acid. Strength influences how an acid or base is handled and which agricultural task it suits.

1.1.3.1 Strong acids and bases

Strong acids and bases dissociate almost completely in water, giving a high concentration of H⁺ or OH⁻ ions. They are highly corrosive, change pH sharply and must be diluted and handled with full protective equipment.

Strong acids:

• Hydrochloric acid (HCl): ionises fully into H⁺ and Cl⁻ ions and is used in laboratory soil tests and cleaning.
• Sulfuric acid (H₂SO₄): a strong acid and powerful dehydrating agent, used in fertiliser manufacture.
• Nitric acid (HNO₃): a strong acid used to make nitrogen fertilisers such as ammonium nitrate.

Strong bases:

• Sodium hydroxide (NaOH): dissociates fully into Na⁺ and OH⁻ ions and is used for cleaning dairy and poultry equipment.
• Potassium hydroxide (KOH): dissociates fully and is used in making soft soaps and some liquid cleaners.
• Calcium hydroxide (Ca(OH)₂): only slightly soluble, but the portion that dissolves dissociates fully; it is the slaked lime used in soil treatment and in Bordeaux mixture.

Example: 0.1 mol/L HCl has a pH of 1 because it gives 0.1 mol/L of H⁺, so a few drops on bare skin or a metal sprayer part cause rapid damage.

1.1.3.2 Weak acids and bases

Weak acids and bases dissociate only partially in water, so an equilibrium exists between the unionised molecules and their ions. They give a lower concentration of H⁺ or OH⁻ than a strong acid or base of the same concentration, and they are generally safer to handle.

Weak acids:

• Ethanoic (acetic) acid: the acid in vinegar; only a small fraction of its molecules ionise. It can act as a natural herbicide in organic farming.
• Citric acid: found in citrus fruits and used in food preservation.
• Carbonic acid (H₂CO₃): formed when carbon dioxide dissolves in rainwater and soil water, making them slightly acidic.
• Phosphoric acid (H₃PO₄): a weak acid used to make phosphate fertilisers.
• Lactic acid: produced during silage fermentation, where it lowers pH and preserves the forage.

Weak bases:

• Ammonia (NH₃): accepts protons from water to form NH₄⁺ and OH⁻ ions, and is a nitrogen source in fertilisers.
• Ammonium hydroxide solution: the name often given to ammonia dissolved in water.
• Organic amines: compounds such as methylamine, which are weakly basic.
• Sodium carbonate and sodium hydrogencarbonate solutions: these salts give weakly alkaline solutions, used in mild cleaning of equipment.

The ionisation of ethanoic acid is written with a reversible arrow:

CH₃COOH ⇌ CH₃COO⁻ + H⁺

Organic and inorganic acids and bases

Acids and bases can also be grouped by origin. Organic compounds contain carbon and are often produced by living organisms, while inorganic compounds are mostly mineral or industrially made.

• Organic and inorganic acids: organic acids contain carbon and are commonly found in plants, such as citric and oxalic acids, influencing plant metabolism and soil interactions. Inorganic (mineral) acids, like sulfuric acid, are industrially produced and used in fertiliser manufacture.
• Organic and inorganic bases: organic bases include compounds like alkaloids derived from plants, whereas inorganic bases include minerals like calcium hydroxide used in soil amendment. Both types have roles in crop protection and soil health.

Monobasic, dibasic and tribasic acids

Acids are also classified by the number of hydrogen ions they can donate per molecule. Monobasic acids donate one H⁺ ion (for example hydrochloric acid, HCl), dibasic acids donate two (for example sulfuric acid, H₂SO₄, or sulfurous acid, H₂SO₃), and tribasic acids donate three (for example phosphoric acid, H₃PO₄). Phosphoric acid is crucial in producing phosphate fertilisers.

1.1.4 Uses of acids and bases

Acids and bases have diverse applications in agriculture, influencing soil management, pest control, and post-harvest treatment.

Soil pH adjustment

Bases such as agricultural lime are applied to acidic soils to raise pH, enhancing nutrient availability and crop growth. Example: acidic soils in tea and coffee growing highlands benefit from lime application, because it reduces aluminium toxicity and makes phosphorus more available. Where soils are too alkaline, acidifying materials such as elemental sulfur are used to lower pH over time.

Fertiliser production

Acids like sulfuric, nitric and phosphoric acids are raw materials in manufacturing fertilisers such as ammonium sulfate, ammonium nitrate and superphosphate. Example: ammonia (a base) reacts with sulfuric acid to form ammonium sulfate, a nitrogen fertiliser for maize and wheat.

2NH₃ + H₂SO₄ → (NH₄)₂SO₄

Pest and disease control

Some acids and bases are used in formulating pesticides and herbicides. Example: acetic acid can act as a natural herbicide in organic farming, while slaked lime is mixed with copper sulfate to make Bordeaux mixture, a fungicide used on vegetable and fruit crops.

Food preservation

Organic acids like citric, lactic and acetic acids are used to preserve fruits and vegetables after harvest. They lower pH to a level at which many spoilage microbes cannot grow, reducing spoilage and extending shelf life for market distribution. Example: pickling vegetables in vinegar and fermenting silage both rely on acidity.

Cleaning and sanitation

Bases such as sodium hydroxide are used in cleaning farm equipment and storage facilities because they break down fats, oils and proteins. This helps prevent microbial contamination in milking parlours and feed stores, supporting food safety and animal health. Acid cleaners are also used to remove milk stone and mineral scale from dairy equipment.

1.1.5 Preparation of acids and bases

Understanding how acids and bases are prepared is essential for their safe and effective use in agricultural processes and extension services.

Laboratory preparation of acids

Acids can be prepared by dissolving acidic (non-metal) oxides in water, by displacing a volatile acid from its salt using a less volatile acid, or by diluting a concentrated stock.

CO₂ + H₂O ⇌ H₂CO₃

SO₃ + H₂O → H₂SO₄

NaCl + H₂SO₄ → NaHSO₄ + HCl

In agricultural laboratories, dilute acids are usually prepared from concentrated stock for soil and water testing.

Equipment for preparing a dilute acid:

Procedure for preparing 500 mL of 0.1 mol/L hydrochloric acid from concentrated acid of about 12 mol/L:

1. Put on goggles, gloves and an apron, and work in a well-ventilated area or fume cupboard.
2. Calculate the volume of concentrated acid needed using C₁V₁ = C₂V₂, so V₁ = (0.1 × 500) ÷ 12 = 4.17 mL.
3. Pour about 300 mL of distilled water into the volumetric flask first, because acid must always be added to water and never water to acid.
4. Measure 4.17 mL of concentrated acid in a small measuring cylinder, handling it slowly.
5. Add the acid slowly down the inside of the flask while swirling gently, and allow any heat to dissipate.
6. Rinse the cylinder into the flask with distilled water, then top up to the 500 mL mark using the wash bottle.
7. Stopper the flask, invert several times to mix, and label it with the name, concentration, date and hazard warning.

Laboratory preparation of bases

Bases are commonly prepared by dissolving metal oxides or hydroxides in water, or by reacting very reactive metals with water.

CaO + H₂O → Ca(OH)₂

Na₂O + H₂O → 2NaOH

2Na + 2H₂O → 2NaOH + H₂

Calcium hydroxide, known as slaked lime, is produced by adding water to calcium oxide (quicklime), a reaction that releases a lot of heat. It is used extensively in soil treatment and fungicide preparation. Ammonia solution is prepared by dissolving ammonia gas in water.

Industrial production methods

Large-scale acid production involves processes such as the Contact Process for sulfuric acid. In this process sulfur dioxide is oxidised to sulfur trioxide over a vanadium(V) oxide catalyst; the sulfur trioxide is absorbed in concentrated sulfuric acid to form oleum, which is then diluted with water to give the acid.

2SO₂ + O₂ ⇌ 2SO₃

The Haber Process makes ammonia from nitrogen and hydrogen, and the ammonia is the starting material for nitrogen fertilisers.

N₂ + 3H₂ ⇌ 2NH₃

Sodium hydroxide is made industrially by the electrolysis of concentrated sodium chloride solution, and sodium carbonate, which gives an alkaline solution, is made by the Solvay Process.

Preparation from salt hydrolysis

Hydrolysis is the reaction of the ions of a salt with water. Salts of a weak base and a strong acid, such as ammonium sulfate, give slightly acidic solutions, while salts of a weak acid and a strong base, such as sodium carbonate, give alkaline solutions.

NH₄⁺ + H₂O ⇌ NH₃ + H₃O⁺

CO₃²⁻ + H₂O ⇌ HCO₃⁻ + OH⁻

Example: repeated use of ammonium-based fertilisers gradually acidifies soil, which is why liming may be needed.

Safety and handling in preparation

Preparation of acids and bases requires strict safety measures, including use of protective gear and proper ventilation. Extension officers must train farmers on safe handling to prevent chemical burns and environmental contamination.

• Protective equipment: wear goggles, gloves, an apron and closed shoes whenever stock chemicals are handled.
• Dilution order: always add acid to water, and add solid alkalis such as sodium hydroxide slowly to water, stirring, because both release heat.
• Labelling: label every container clearly with its contents, concentration and hazard.
• Storage: store acids and bases separately, away from each other, from metals and from children and animals.
• Spills: neutralise small acid spills with sodium hydrogencarbonate or lime, and small alkali spills with a weak acid such as vinegar, then rinse with plenty of water.
• First aid: rinse affected skin or eyes with running water for at least 15 minutes and seek medical help.

Chapter Summary

This chapter explored the fundamental concepts of acids and bases, beginning with their definitions and inherent characteristics, followed by their classification into strong and weak categories based on their ionization properties. The uses and methods of preparation for both acids and bases were discussed, highlighting their practical significance. The chapter then examined salts, defining their nature and describing their distinctive properties, classifications, and common applications. Preparation techniques for salts were also outlined to provide a comprehensive understanding. Further, the chapter addressed the application of ionic and chemical equilibrium principles, emphasizing their role in maintaining system stability. The properties of gases were analyzed with respect to their behavior under various conditions, and finally, practical procedures for preparing acids, bases, and salts were detailed to reinforce theoretical knowledge through laboratory practice.