In water engineering, understanding inorganic chemistry principles is essential for managing water quality effectively. Water sources often contain inorganic constituents like minerals, metals, and salts that influence water chemistry and treatment processes. Kenyan water engineers must grasp these principles to control contamination, prevent scaling and corrosion, and ensure compliance with national water quality standards. This chapter introduces fundamental inorganic chemistry concepts relevant to water quality management in Kenya’s diverse water supply systems.

1.1 Inorganic chemistry

Inorganic chemistry deals with compounds that generally do not contain carbon-hydrogen bonds and includes minerals, salts, metals, and non-metals commonly found in water. For water engineers, it is critical to understand how inorganic substances interact in water to predict behavior such as solubility, reactivity, and potential health impacts. Kenya’s water sources, including rivers, boreholes, and reservoirs, often contain inorganic ions like calcium, magnesium, iron, and chloride, which affect water hardness, taste, and suitability for domestic or industrial use.

1.1.1 Periodic table

The periodic table is the arrangement of all known chemical elements in order of increasing atomic number, in rows called periods and columns called groups, so that elements with similar properties fall in the same column. It is based on the periodic law: the properties of elements are a periodic function of their atomic numbers. The modern table has 118 elements in 7 periods and 18 groups, with metals on the left, non-metals on the right and metalloids along a stepped line between them. For a water engineer, the table explains why sodium and potassium salts are always soluble, why calcium and magnesium cause hardness, and why chlorine and fluorine are so reactive. Example: a water quality laboratory at a county water company uses the position of iron and manganese among the transition metals to explain why both can be removed by oxidation and filtration.

Development of the periodic table

• Döbereiner's triads: Early grouping of elements in threes with similar properties, such as calcium, strontium and barium.
• Newlands' law of octaves: Elements arranged by atomic mass showed similar properties every eighth element, but the pattern failed beyond calcium.
• Mendeleev's table (1869): Arranged elements by atomic mass and similar properties, leaving gaps for undiscovered elements and correctly predicting their properties.
• Moseley's work (1913): Showed that atomic number, not atomic mass, is the true basis of the order, removing the anomalies in Mendeleev's table.
• Modern long form: Arranges elements by atomic number and electronic configuration into s, p, d and f blocks.

Main features of the modern periodic table

• Periods: Seven horizontal rows; the period number equals the number of occupied electron shells.
• Groups: Eighteen vertical columns; elements in a group have the same number of outer electrons and similar chemical properties.
• Metals, non-metals and metalloids: About three quarters of elements are metals on the left and centre; non-metals are on the upper right; metalloids such as silicon and arsenic lie between them.
• Blocks: The table is divided into s, p, d and f blocks according to the subshell receiving the last electron.
• Lanthanides and actinides: The f-block elements are placed below the main table to keep it a manageable width.

1.1.2 Elements

An element is a pure substance made of only one kind of atom, all having the same atomic number, which cannot be broken down into simpler substances by chemical means. Each element has a name and a symbol of one or two letters, the first capital and the second small, such as Ca for calcium and Cl for chlorine. Elements combine to form the compounds found in water, such as calcium carbonate, sodium chloride and aluminium sulphate. Example: a borehole drilling contractor's water analysis reports the elements calcium, magnesium, iron and fluorine, each of which affects how the water must be treated.

Classes of elements

Elements of importance in water engineering

• Calcium and magnesium: Their dissolved ions cause water hardness, scale in pipes and boilers, and high soap use.
• Iron and manganese: Cause reddish or black staining and metallic taste, common in borehole water.
• Chlorine: Used as a disinfectant in the form of chlorine gas, sodium hypochlorite or calcium hypochlorite.
• Fluorine: Occurs as fluoride, which in excess causes dental and skeletal fluorosis, especially in parts of the Rift Valley.
• Aluminium: Used in alum (aluminium sulphate) as a coagulant in water treatment.
• Nitrogen: Occurs as nitrate and ammonia from fertilisers and sewage, indicating pollution.

1.1.3 Atomic structure

An atom is the smallest particle of an element that can take part in a chemical reaction. It consists of a tiny, dense nucleus containing protons and neutrons, surrounded by electrons arranged in energy levels or shells. The number of protons decides which element the atom is, while the electrons, especially those in the outer shell, decide how it reacts. Example: a chemistry trainer explains that a sodium atom becomes the Na⁺ ion found in salty borehole water by losing its single outer electron.

Key terms

• Atomic number (Z): The number of protons in the nucleus, which also equals the number of electrons in a neutral atom.
• Mass number (A): The total number of protons and neutrons in the nucleus.
• Number of neutrons: Found as A − Z.
• Isotopes: Atoms of the same element with the same atomic number but different mass numbers, such as chlorine-35 and chlorine-37.
• Relative atomic mass: The average mass of an element's atoms, taking account of the abundance of its isotopes, compared with one twelfth of a carbon-12 atom.
• Ions: Atoms that have gained electrons (negative ions) or lost electrons (positive ions).

Worked examples

Example 1: Find the number of protons, neutrons and electrons in an atom of chlorine-35 (Z = 17).

Protons = Z = 17

Neutrons = A − Z = 35 − 17 = 18

Electrons = 17

Answer: 17 protons, 18 neutrons, 17 electrons

Example 2: How many electrons are in a sodium ion, Na⁺ (Z = 11)?

Electrons = 11 − 1

Answer: 10 electrons

Example 3: Chlorine is 75 percent chlorine-35 and 25 percent chlorine-37. Find its relative atomic mass.

Relative atomic mass = (35 × 75 + 37 × 25) ÷ 100

= (2,625 + 925) ÷ 100 = 3,550 ÷ 100

Answer: 35.5

1.1.4 Electronic configuration of the first 30 elements.

Electronic configuration is the arrangement of electrons in the shells and subshells of an atom. It explains an element's position in the periodic table, the ions it forms and how it reacts. Electrons fill the lowest available energy levels first: the first shell holds up to 2 electrons, the second up to 8, and the third up to 18, although for the first 20 elements it is written as filling to 8 before the fourth shell starts. Example: knowing that iron is [Ar] 3d⁶ 4s² helps a treatment plant chemist explain why iron forms both Fe²⁺ and Fe³⁺ ions in borehole water.

Rules for filling electrons

• Aufbau principle: Electrons enter the lowest energy subshell available, in the order 1s, 2s, 2p, 3s, 3p, 4s, 3d.
• Pauli exclusion principle: Each orbital holds at most two electrons, with opposite spins.
• Hund's rule: Electrons occupy orbitals of the same subshell singly before pairing up.
• Subshell capacities: s holds 2, p holds 6 and d holds 10 electrons.
• Exceptions: Chromium and copper take one electron from 4s to give a half-full (3d⁵) or full (3d¹⁰) d subshell, which is more stable.

Electronic configuration of the first 30 elements

Procedure for writing an electronic configuration

1. Find the atomic number: Look up Z, which gives the number of electrons in the neutral atom.
2. Adjust for ions: Subtract one electron for each positive charge or add one for each negative charge.
3. Follow the filling order: Place electrons in 1s, 2s, 2p, 3s, 3p, 4s, then 3d.
4. Respect subshell capacities: Fill s to 2, p to 6 and d to 10 before moving on.
5. Apply the exceptions: Write chromium as 3d⁵ 4s¹ and copper as 3d¹⁰ 4s¹.
6. Remove 4s electrons first for transition metal ions: When a transition metal forms an ion, it loses 4s electrons before 3d electrons.
7. Check the total: Add up the superscripts to confirm they equal the number of electrons.

Worked example

Write the configuration of iron (Z = 26) and of the Fe³⁺ ion.

Fe: 26 electrons = 1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d⁶ = [Ar] 3d⁶ 4s²

Fe³⁺: 26 − 3 = 23 electrons; remove the two 4s electrons, then one 3d electron

Answer: Fe is [Ar] 3d⁶ 4s² and Fe³⁺ is [Ar] 3d⁵

1.1.5 Groups, periods and blocks.

Groups, periods and blocks are the three ways the periodic table is divided. A group is a vertical column of elements with the same number of outer electrons and similar chemical properties; a period is a horizontal row of elements with the same number of occupied electron shells; and a block is a region named after the subshell (s, p, d or f) that receives the last electron. Knowing an element's group, period and block allows its properties to be predicted from its position. Example: a trainee at a water treatment works predicts that bromine, below chlorine in group 17, will also act as a disinfectant, though less strongly.

Main groups and families

The four blocks

• s-block: Groups 1 and 2, plus helium; the last electron enters an s subshell; reactive metals.
• p-block: Groups 13 to 18; the last electron enters a p subshell; includes non-metals, metalloids and some metals.
• d-block: Groups 3 to 12, the transition metals; the last electron enters a d subshell.
• f-block: Lanthanides and actinides; the last electron enters an f subshell.
• Locating an element: The number of shells gives the period, the outer electrons give the group of main-group elements, and the last subshell filled gives the block.

Worked example

Find the period, group and block of sulphur (Z = 16).

Configuration: 2.8.6, or [Ne] 3s² 3p⁴

Number of shells = 3, so period 3

Outer electrons = 6, so group 16 (VI)

Last electron enters 3p, so p-block

Answer: period 3, group 16 (VI), p-block

1.1.6 Properties of Elements

The properties of elements change in regular patterns across periods and down groups, because they depend on the number of protons and the arrangement of electrons. These patterns, called periodic trends, allow chemists to predict how an element will behave without testing it. Properties are divided into physical properties (1.1.7), which can be observed without changing the substance, and chemical properties (1.1.8), which show how it reacts. Example: a laboratory technician at a bottled water plant predicts that potassium will react more violently with water than sodium because reactivity of group 1 metals increases down the group.

Periodic trends

Key terms

• Atomic radius: Half the distance between the nuclei of two bonded atoms of the same element.
• Ionisation energy: The energy needed to remove one mole of electrons from one mole of gaseous atoms.
• Electronegativity: The ability of an atom to attract the shared electrons in a bond; fluorine is the most electronegative element.
• Metallic character: How readily an element loses electrons to form positive ions.
• Electron affinity: The energy released when an atom gains an electron to form a negative ion.

1.1.7 Physical

Physical properties are characteristics that can be observed or measured without changing the substance into a new one, such as state, melting and boiling points, density, appearance and conductivity. They differ widely between metals and non-metals and change in a regular way across the periodic table. Engineers use them to choose materials, for example selecting metals that conduct heat for heat exchangers or that resist corrosion for pipes. Example: a pipe fabrication contractor chooses copper for hot water pipework because it conducts heat well, resists corrosion and can be bent easily.

Comparison of physical properties of metals and non-metals

Physical trends in the periodic table

• Melting points across period 3: Rise from sodium to silicon, which has a giant covalent structure, then fall sharply for the molecular non-metals phosphorus, sulphur, chlorine and argon.
• Melting points down group 1: Decrease from lithium to caesium as the metallic bonding weakens with larger atoms.
• Melting and boiling points down group 17: Increase from fluorine (gas) to chlorine (gas), bromine (liquid) and iodine (solid) as the molecules become larger.
• Density: Generally increases down a group as atoms become heavier.
• Electrical conductivity: High for metals on the left, falling to almost zero for the non-metals on the right of each period.

1.1.8 chemical

Chemical properties describe how an element reacts with other substances to form new substances, such as its reactions with oxygen, water, acids and other elements. They depend mainly on the number of outer electrons: metals tend to lose electrons to form positive ions, while non-metals tend to gain or share electrons. Chemical properties explain corrosion of pipes, the hardness of water, the action of disinfectants and the choice of treatment chemicals. Example: a dairy processor's boiler suffers scale because calcium and magnesium in the feed water form insoluble carbonates when heated.

Chemical properties of metals

• Reaction with oxygen: Metals form basic oxides, for example 2Mg + O₂ → 2MgO; iron rusts in moist air, corroding steel pipes.
• Reaction with water: Reactive metals form hydroxides and hydrogen, for example 2Na + 2H₂O → 2NaOH + H₂ and Ca + 2H₂O → Ca(OH)₂ + H₂.
• Reaction with acids: Metals above hydrogen in the reactivity series displace hydrogen, for example Mg + 2HCl → MgCl₂ + H₂.
• Formation of positive ions: Metals lose electrons, for example Ca → Ca²⁺ + 2e⁻.
• Displacement reactions: A more reactive metal displaces a less reactive one from its salt, for example Zn + CuSO₄ → ZnSO₄ + Cu.

Chemical properties of non-metals

• Reaction with oxygen: Non-metals form acidic oxides, for example S + O₂ → SO₂, which dissolves in water to form an acidic solution.
• Formation of negative ions: Non-metals gain electrons, for example Cl₂ + 2e⁻ → 2Cl⁻.
• Reaction of halogens with water: Chlorine reacts with water to form hypochlorous acid, the active disinfectant: Cl₂ + H₂O ⇌ HCl + HOCl.
• Displacement among halogens: A more reactive halogen displaces a less reactive one, for example Cl₂ + 2KBr → 2KCl + Br₂.
• Reaction with metals: Non-metals combine with metals to form salts, for example 2Na + Cl₂ → 2NaCl.

Reactivity series of common metals

From most to least reactive: potassium, sodium, calcium, magnesium, aluminium, zinc, iron, lead, (hydrogen), copper, silver, gold. Metals high in the series corrode easily and cannot be used unprotected for water pipes, while copper, low in the series, resists corrosion and is widely used in plumbing.

Nature of oxides across period 3

1.1.9 The Scope of Inorganic Chemistry in Water Engineering

Inorganic chemistry in water engineering primarily focuses on the study of mineral ions, metal contaminants, and their chemical behavior in aqueous environments. This knowledge aids in identifying water quality parameters such as hardness, alkalinity, and the presence of toxic metals. For instance, understanding the chemistry of iron and manganese helps in designing removal processes for borehole water in parts of Kenya where these metals are prevalent.

Key Aspects of Inorganic Chemistry Applications

• Mineral Ion Identification: Determining ions like calcium, magnesium, sodium, and potassium that influence water hardness and scaling tendencies.
• Metal Contaminant Control: Understanding metals such as lead, arsenic, and chromium to prevent toxic exposure in drinking water.
• Chemical Reactions in Water: Studying precipitation, complexation, and redox reactions that affect water treatment efficiency.
• pH and Alkalinity Management: Controlling acidity and buffering capacity to maintain water stability and prevent corrosion.
• Water Quality Standard Compliance: Applying inorganic chemistry to meet Kenya Bureau of Standards (KEBS) regulations for potable water.

1.1.10 Properties of Inorganic Compounds Relevant to Water Quality

Inorganic compounds in water exhibit properties such as solubility, ionic charge, and reactivity that influence their behavior and treatment. For example, calcium carbonate solubility affects scaling in pipes and boilers, while the presence of chloride ions can accelerate corrosion in metal infrastructure. Water engineers must evaluate these properties to select appropriate treatment methods.

Properties Impacting Water Treatment

• Solubility: Determines whether a compound remains dissolved or precipitates, influencing turbidity and scaling.
• Ionic Charge and Size: Affect mobility and interaction with other constituents, impacting coagulation and ion exchange processes.
• Reactivity: Governs chemical transformations such as oxidation of iron or reduction of nitrate.
• Stability: Some compounds remain stable under varying pH and temperature, affecting treatment consistency.
• Toxicity: Certain inorganic substances, like arsenic, pose health risks even at low concentrations, requiring careful monitoring.

1.1.11 Role of Inorganic Chemistry in Water Treatment Processes

Water treatment relies heavily on inorganic chemistry principles to remove or neutralize contaminants. Processes such as coagulation, flocculation, precipitation, and ion exchange are designed based on the chemical nature of inorganic substances in water. For example, lime softening exploits the precipitation of calcium and magnesium salts to reduce hardness in hard water supplies.

Chemical Principles Guiding Treatment

• Precipitation Reactions: Used to remove hardness by converting dissolved ions into insoluble solids.
• Oxidation-Reduction Reactions: Employed to convert soluble iron and manganese into insoluble oxides that can be filtered out.
• Adsorption and Ion Exchange: Techniques that rely on ionic interactions to remove heavy metals and other contaminants.
• pH Adjustment: Critical for optimizing chemical reactions and ensuring effective disinfection.
• Complexation: Formation of complexes with organic or inorganic ligands to stabilize or remove contaminants.

1.1.12 Inorganic Contaminants and Their Impact on Water Quality

Inorganic contaminants originate from natural sources or human activities such as mining, agriculture, and industrial discharge. Their presence affects water aesthetics, safety, and usability. For example, nitrate from fertilizer runoff in intensively farmed areas threatens drinking water quality, requiring engineers to implement targeted remediation.

Common Inorganic Contaminants in Kenyan Water Sources

• Nitrates and Nitrites: Cause methemoglobinemia ("blue baby syndrome") and indicate agricultural pollution.
• Heavy Metals (Lead, Arsenic, Mercury): Toxic even at trace levels, causing chronic health issues.
• Fluoride: Naturally occurring in some regions such as parts of the Rift Valley, where excessive levels cause dental and skeletal fluorosis.
• Chlorides and Sulfates: Affect taste and corrosivity, impacting water distribution systems.
• Calcium and Magnesium: Major contributors to water hardness, influencing soap usage and scaling.

Chapter Summary

This chapter explored fundamental inorganic chemistry principles essential for understanding water quality. It began by examining chemical bonding and structures, highlighting the differences between metallic, covalent, and ionic bonds and their properties. The chapter also addressed hazardous chemicals, emphasizing their identification and safe management. It introduced units of concentration and the importance of accurate measurement in chemical analysis. Students learned to apply nomenclature rules and write chemical formulae for inorganic compounds, alongside the mole concept and preparation of aqueous solutions. Stoichiometry was covered to enable calculation of reactants and products in chemical reactions. The chapter further discussed chemical reagents and techniques for solution preparation. Finally, it outlined the Global Hazardous System for chemical classification and labelling, including best practices for storage, packaging, transportation, handling, and disposal of hazardous wastes, with a focus on personal safety and accident prevention.