This chapter explores the application of physical chemistry principles specifically tailored for horticulture professionals in Kenya. Understanding these principles is crucial for managing soil chemistry, nutrient availability, and agrochemical interactions that influence crop health and productivity. The knowledge of acid-base chemistry, in particular, aids in optimizing soil pH, preparing fertilizers and pesticides, and ensuring safe handling of chemicals in farm management. This foundation equips horticulturists with the skills to make informed decisions that enhance crop yield and environmental sustainability.

1.1 Acids and bases

Acid and base chemistry forms the cornerstone of many horticultural practices in Kenya, from soil amendment to pest control. Acids and bases influence nutrient solubility and availability, affecting plant growth and health. Furthermore, many agrochemicals used in horticulture, such as herbicides and fungicides, are formulated based on acid-base reactions. Understanding their properties and interactions enables precise application and prevents damage to crops or the environment.

1.1.1 Definition of terms

Acids, bases and salts are the three linked groups of compounds met in every soil test, fertiliser mix and spray tank. Acids and bases are characterised by their ability to donate or accept protons, and a salt is what remains when an acid and a base react. Knowing these definitions allows a horticulturist to read a fertiliser label or a soil report and predict how the product will behave in the soil.

1.1.1.1 Acids

An acid is a substance that releases hydrogen ions (H⁺) when dissolved in water, giving a solution with a pH below 7. In a horticultural store, hydrochloric, sulphuric, nitric and phosphoric acids are the usual examples, and weak organic acids such as citric and acetic acid are found in plant sap, compost and vinegar.

• Proton donors: Acids release H⁺ ions in water, increasing the acidity of the solution. For example, sulphuric acid added to alkaline irrigation water releases H⁺ and lowers its pH.
• Hydronium formation: In water a free H⁺ ion attaches to a water molecule to form the hydronium ion (H₃O⁺). This is why acid solutions are said to contain H⁺ even though the bare ion does not exist alone.
• Electron pair acceptors: According to Lewis theory, acids accept electron pairs during chemical reactions. This is important in the complexation of micronutrients such as iron and manganese in the soil.
• Characteristic pH: Solutions with pH less than 7 are acidic. Low pH increases the solubility of micronutrients such as iron and manganese, and in very acidic soils it can raise them to toxic levels.
• Common examples: Hydrochloric acid (HCl), sulphuric acid (H₂SO₄), nitric acid (HNO₃), phosphoric acid (H₃PO₄), acetic acid (CH₃COOH) and citric acid are the acids most often met in horticulture. Their names also decide the names of the salts they form.

1.1.1.2 Bases

A base is a substance that accepts hydrogen ions or, in Arrhenius terms, releases hydroxide ions (OH⁻) in water, giving a pH above 7. Bases used in horticulture include quicklime (calcium oxide, CaO), hydrated lime (calcium hydroxide, Ca(OH)₂), agricultural lime (mainly calcium carbonate, CaCO₃), sodium hydroxide and ammonia.

• Proton acceptors: Bases accept hydrogen ions (H⁺) in aqueous solutions, reducing acidity and raising pH. This is the action that makes lime effective for neutralising acidic soils.
• Electron pair donors: According to Lewis theory, bases donate electron pairs to form bonds. This affects soil chemistry and the way nutrients bind to soil particles.
• Characteristic pH: Solutions with pH greater than 7 are basic or alkaline. In strongly alkaline soils the availability of nutrients such as phosphorus, iron and zinc is reduced.
• Alkalis: A base that dissolves in water is called an alkali. Sodium hydroxide and potassium hydroxide dissolve readily, while calcium hydroxide is only slightly soluble.
• Common examples: Metal oxides, hydroxides, carbonates and ammonia are the usual bases in a horticultural setting. Examples include CaO, Ca(OH)₂, CaCO₃, NaOH and NH₃.

1.1.1.3 Salts

A salt is an ionic compound formed when the hydrogen ions of an acid are replaced by metal ions or by the ammonium ion (NH₄⁺). Most fertilisers that a horticulturist applies are salts, for example potassium nitrate, ammonium sulphate and calcium nitrate. The general reaction is acid + base → salt + water, for example:

HCl + NaOH → NaCl + H₂O

The full study of salt properties, classification, uses and preparation is covered in the topic on salts.

• Ionic compounds: A salt is made of positive ions (cations) and negative ions (anions) held together in a crystal. Potassium nitrate, for instance, contains K⁺ and NO₃⁻ ions.
• Origin from an acid and a base: The metal or ammonium part comes from the base and the negative part comes from the acid. Ammonium sulphate comes from ammonia and sulphuric acid.
• Naming from the parent acid: Hydrochloric acid gives chlorides, sulphuric acid gives sulphates, nitric acid gives nitrates and phosphoric acid gives phosphates. A label reading "calcium nitrate" therefore tells you the acid involved was nitric acid.
• Dissolving into ions: Many salts dissolve in water and separate into ions that plant roots can absorb. This is why soluble salts are used in nutrient solutions for crops grown in the greenhouse.
• Link to horticulture: Nitrates, sulphates, phosphates and chlorides of potassium, calcium, magnesium and ammonium supply the major plant nutrients. They are the end products of the acid-base reactions described in this topic.

1.1.2 Characteristics of acids and bases

Understanding the distinct characteristics of acids and bases is vital for horticulturists when handling chemicals, diagnosing soil issues, and applying amendments. These characteristics are the observable and measurable behaviours used to identify an acid or a base in the laboratory or in the field. For example, a strip of litmus paper dipped into a soil-water mixture can show quickly whether the soil is acidic or alkaline.

Characteristics of acids

• Sour taste: Acids taste sour, as seen in lemon juice (citric acid) and vinegar (acetic acid). Laboratory and agrochemical acids must never be tasted because they are corrosive and poisonous.
• Corrosive nature: Strong acids can corrode metals and damage skin and plant tissues if improperly applied, so gloves, goggles and correct dilution are required.
• Reactivity with metals: Acids react with metals such as zinc or iron to release hydrogen gas. For example, acid left in a mild steel sprayer or tank corrodes it.
• Indicator response: Acids turn blue litmus paper red and give a red or pink colour with methyl orange. This is the basis of simple soil acidity tests.
• Conductivity: Acids conduct electricity in solution because they ionise to form mobile ions. A conductivity meter in a nutrient solution responds to such ions.

Characteristics of bases

• Bitter taste: Bases taste bitter, though tasting agrochemicals is unsafe and must never be done.
• Slippery feel: Strong alkalis such as sodium hydroxide feel slippery because they react with the oils in skin to form soap. This is also why they cause severe burns and must not be touched.
• Reactivity with oils and fats: Bases react with oils and fats to form soaps (saponification). Insecticidal soaps used against soft-bodied pests are made from fatty acid salts formed in this way.
• Indicator response: Bases turn red litmus paper blue and give a pink colour with phenolphthalein. This helps in rapid screening of soil and water pH.
• Conductivity: Dissolved bases also conduct electricity because they release ions. In soil, this contributes to the electrolyte balance of the soil solution.

The pH scale and neutralisation

The pH scale gives a measurable value for the acidity or alkalinity described above, and neutralisation is the reaction that corrects it. These ideas allow a horticulturist to explain soil test results and calculate treatments. Example: a soil with a pH of 4.5 is more acidic than one with pH 6.5 by a factor of 100 in hydrogen ion concentration.

• pH scale: The pH scale runs from 0 to 14 at 25 °C, with 7 as neutral, and it is defined by pH = −log₁₀[H⁺], where [H⁺] is in mol/L. Each unit change on the scale is a tenfold change in H⁺ concentration.
• Bronsted-Lowry action in soil: Acids donate protons and bases accept them, so ammonia in the soil reacts as NH₃ + H⁺ → NH₄⁺. This explains acid-base behaviour even where no hydroxide ion is present.
• Neutralisation with a hydroxide: An acid and a base react to form a salt and water, for example Ca(OH)₂ + 2HCl → CaCl₂ + 2H₂O.
• Neutralisation in soil: Lime neutralises soil acidity as CaCO₃ + 2H⁺ → Ca²⁺ + H₂O + CO₂. The calcium released also becomes available to the crop.

Worked example: pH of a dilute hydrochloric acid solution used for soil extraction.

Formula: pH = −log₁₀[H⁺]

Substitution: [H⁺] = 0.001 mol/L (HCl is a strong acid, so [H⁺] equals the acid concentration)

Working: pH = −log₁₀(0.001) = −(−3)

Result: pH = 3

1.1.3 Classification of acids and bases

The classification of acids and bases helps horticulturists select appropriate chemicals for soil and crop management based on their strength, origin and behaviour. Strength is the most important basis, because it decides how much of the substance ionises in water and so how fast the pH changes. Other classifications, such as origin and number of protons, explain why one acid is used in fertiliser manufacture and another in the compost heap.

1.1.3.1 Strong acids and bases

Strong acids and bases ionise or dissociate completely in water, so every molecule or formula unit contributes ions to the solution. A small amount of strong acid therefore changes pH sharply and is dangerous to skin, crops and equipment. In a horticultural laboratory, dilute hydrochloric acid is used to extract nutrients from soil samples, and sodium hydroxide solution is used to clean tools.

Strong acids:

• Hydrochloric acid (HCl): It ionises completely to H⁺ and Cl⁻ and is used in dilute form in some soil testing procedures.
• Sulphuric acid (H₂SO₄): It is a strong diprotic acid used in making phosphatic and ammonium sulphate fertilisers.
• Nitric acid (HNO₃): It ionises completely and is used in the manufacture of nitrate fertilisers.
• Low pH at low concentration: A 0.1 mol/L solution of a strong monoprotic acid has a pH of about 1. Even dilute solutions are therefore hazardous.
• Handling: Strong acids need protective clothing and must be diluted by adding the acid slowly to water, never the reverse.

Strong bases:

• Sodium hydroxide (NaOH): It dissociates completely into Na⁺ and OH⁻ and is used in alkaline detergents and some chemical formulations.
• Potassium hydroxide (KOH): It is also fully dissociated and is used in making potassium soaps.
• Calcium hydroxide (Ca(OH)₂): It dissociates fully but is only slightly soluble, so its solutions are limited in strength. It is used as hydrated lime in soil and in some fungicide mixtures.
• High pH at low concentration: A 0.1 mol/L solution of NaOH has a pH of about 13.
• Handling: Strong bases burn skin and eyes, so goggles, gloves and an apron must be worn when preparing solutions.

1.1.3.2 Weak acids and bases

Weak acids and bases ionise or dissociate only partially in water, so most of the molecules remain unchanged and an equilibrium is set up. Their solutions change pH gradually and are safer to handle, which is why they are used in foods, composts and natural pest control. Example: acetic acid in vinegar, CH₃COOH ⇌ CH₃COO⁻ + H⁺, is sometimes used as a contact herbicide on young weeds.

Weak acids and bases in horticulture:

• Acetic acid: Found in vinegar and used for weed control and cleaning.
• Citric and oxalic acids: Organic acids released by roots and microbes that help dissolve soil nutrients.
• Carbonic acid: Formed when carbon dioxide dissolves in rain and soil water, slowly acidifying the soil.
• Ammonia (NH₃): A weak base, NH₃ + H₂O ⇌ NH₄⁺ + OH⁻, and the starting material for ammonium and urea fertilisers.
• Humic acids: Weak acids from decayed organic matter that improve nutrient holding in the soil.

Classification by origin and composition

• Organic acids: These contain carbon and hydrogen, for example citric, oxalic and acetic acids. They occur naturally in plants and decaying matter and influence soil organic matter processes.
• Inorganic acids: These are mineral acids that do not contain carbon in the acidic group, such as sulphuric, nitric and hydrochloric acids. Sulphuric acid is used in the production of phosphate fertilisers.
• Organic bases: These contain carbon, such as the alkaloids found in some plants, which may influence pest resistance.
• Inorganic bases: These include calcium hydroxide, sodium hydroxide and ammonia. Lime products are widely used to neutralise acidic soils.
• Arrhenius and Bronsted-Lowry bases: Arrhenius bases increase OH⁻ concentration in water, while Bronsted-Lowry bases are any substances that accept protons. Ammonia is a Bronsted-Lowry base even though it contains no OH⁻ group.

Classification by number of protons donated

• Monoprotic acids: These donate one proton per molecule, for example HCl, HNO₃ and CH₃COOH.
• Diprotic acids: These donate two protons per molecule, for example sulphuric acid (H₂SO₄) and carbonic acid (H₂CO₃).
• Triprotic acids: These donate three protons per molecule, for example phosphoric acid (H₃PO₄). The step-wise release of protons from phosphoric acid affects its buffering effect in nutrient solutions.
• Buffering effect: Polyprotic acids and their salts can resist pH change over a wider range than monoprotic acids. This is important in soils and fertigation tanks.
• Bases by number of hydroxide groups: Bases such as NaOH release one OH⁻ per unit, while Ca(OH)₂ releases two. This is taken into account when calculating the amount of lime needed to neutralise a given acid.

1.1.4 Uses of acids and bases

Acids and bases have diverse applications in horticulture, from improving soil conditions to enhancing plant protection measures. The right choice depends on the starting pH, the crop and the strength of the chemical. Example: a farmer growing vegetables on a strongly acidic soil applies lime, while a nursery with alkaline borehole water may acidify the water before fertigation.

Uses of acids

• Soil and water pH adjustment: Acids such as sulphuric acid are used to lower the pH of alkaline soils and irrigation water, improving nutrient uptake for crops such as vegetables and flowers. The quantity must be calculated carefully to avoid over-acidification.
• Fertiliser production: Acids are integral in manufacturing fertilisers, for example sulphuric acid in phosphatic fertilisers and nitric acid in nitrate fertilisers.
• Pesticide formulation: Certain herbicides and fungicides use acidic components to maintain stability and effectiveness.
• Post-harvest treatment: Dilute organic acid washes can reduce microbial spoilage on fruits and vegetables and help prolong freshness.
• Cleaning agents: Acidic solutions remove mineral scale and deposits from irrigation pipes, drippers and tanks, ensuring efficient water delivery.
• Soil testing: Dilute acids are used in the laboratory to extract nutrients from soil samples for analysis.

Uses of bases

• Soil amendment: Agricultural lime and hydrated lime raise the pH of acidic soils, improving nutrient availability and crop productivity for crops such as vegetables and fruit trees. Crops that prefer acidic soil, such as tea, are not limed to a high pH.
• Nutrient supply: Ammonia, a weak base, is the starting material for ammonium and urea fertilisers that supply nitrogen, which is essential for plant growth.
• Neutralising spills: Bases such as lime or sodium carbonate are used to neutralise acid spills before cleaning, preventing damage to crops and soil.
• Cleaning and sanitation: Alkaline detergents remove oils, fats and organic residues from tools, trays and storage facilities, reducing disease spread.
• Buffering the soil: Lime and other carbonates neutralise acids added by fertilisers and rain, helping the soil resist rapid pH change.
• Fungicide and pesticide preparation: Hydrated lime is a component of some copper-based fungicide mixtures and soap-based insecticides are made using alkalis.

1.1.5 Preparation of acids and bases

The preparation of acids and bases involves chemical reactions that horticulturists must understand to safely handle and apply these substances in farming operations. Most horticultural enterprises buy concentrated chemicals and prepare working solutions, so dilution is the most frequent preparation task. Understanding how acids and bases are made also explains where fertilisers and lime products come from. Practical preparation is carried out in the practical topic on preparing acids and bases.

Preparation of acids

• Reaction of non-metal oxides with water: Sulphur trioxide reacts with water to form sulphuric acid, SO₃ + H₂O → H₂SO₄, a process relevant in the fertiliser industry. Carbon dioxide dissolving in water forms weak carbonic acid in the same way.
• Direct combination of elements: Hydrogen and chlorine gases combine to give hydrogen chloride, H₂ + Cl₂ → 2HCl. The gas is dissolved in water to give hydrochloric acid, used in soil testing.
• Reaction of a salt with a strong acid: Sodium chloride reacts with concentrated sulphuric acid to release hydrogen chloride gas, NaCl + H₂SO₄ → NaHSO₄ + HCl. Dissolving the gas in water gives hydrochloric acid.
• Dilution of concentrated acids: Concentrated acids are diluted with water to safe working concentrations for field or laboratory use. Always add the acid slowly to water while stirring, never water to acid, because the heat released can make the liquid spit.
• Biological production: Organic acids such as citric and acetic acid are produced by microbial fermentation, and they also form during composting, influencing compost quality.

Preparation of bases

• Reaction of metal oxides with water: Calcium oxide (quicklime) reacts with water to form calcium hydroxide (hydrated lime), CaO + H₂O → Ca(OH)₂, with release of much heat. This is commonly used in soil amendment and fungicide preparation.
• Reaction of reactive metals with water: Sodium reacts with water to form sodium hydroxide and hydrogen gas, 2Na + 2H₂O → 2NaOH + H₂. It is a laboratory method only because the reaction is violent.
• Electrolysis of salt solutions: Sodium hydroxide is produced industrially by electrolysis of concentrated sodium chloride solution (brine), giving NaOH, chlorine and hydrogen. It supplies bases for agricultural chemicals and detergents.
• Ammonia synthesis: Industrial processes combine nitrogen and hydrogen, N₂ + 3H₂ ⇌ 2NH₃, to produce ammonia gas, which is the starting material for ammonium and urea fertilisers or is dissolved in water to make an ammonia solution.
• Dilution of concentrated bases: Concentrated bases are diluted to safe concentrations for cleaning and soil treatment. Solid alkalis are added slowly to water with stirring because dissolving them releases heat.

Chapter Summary

This chapter explores the fundamental principles of acids and bases, beginning with their definitions, characteristics, and classifications into strong and weak categories. It also examines the roles and uses of acids and bases in various chemical processes, alongside methods for their preparation. The study then extends to salts, detailing their properties, classification, practical applications, and preparation techniques. A significant focus is placed on the application of ionic and chemical equilibrium concepts to understand reactions in solution. The chapter further addresses the properties of gases and their relevance in physical chemistry. Practical procedures for preparing acids, bases, and salts are also covered, providing a comprehensive foundation in applying physical chemistry principles across these key areas.