Chapter 8 — Water

What is water?
Water is a simple chemical substance with the formula H2O that exists in nature in several forms: liquid, solid (ice) and vapour (steam). It is found almost everywhere on Earth and is essential for life and many human activities. In pure form water is colourless, tasteless and odourless, but natural water usually contains dissolved substances and suspended particles that change its properties.

Importance of water
Water supports life at many levels. Plants require water for photosynthesis and transport of nutrients; animals and humans use water for digestion, temperature regulation and waste removal. Water is also important for hygiene, food preparation, cooking, cleaning and industrial processes such as cooling and manufacturing. Fresh water availability therefore affects health, agriculture and economies.

Sources of water
Sources commonly used by people include surface water (rivers, lakes, reservoirs), groundwater (wells and boreholes tapping aquifers), rainwater (collected from roofs), springs and water supplied by municipal treatment plants. Each source has different typical impurities: surface water may have suspended soil and microbes; groundwater may contain dissolved minerals like calcium and magnesium; rainwater picks up gases and dust from the air. Some communities also use treated wastewater (recycled water) for irrigation or industrial use after suitable treatment.

Choice of source and protection
Choosing a water source depends on availability, quality and intended use. Protecting catchment areas, preventing sewage discharge into rivers, and maintaining wells are vital to keep the supply safe. Understanding the source helps decide the treatment needed: for example, groundwater high in dissolved salts may need demineralisation, while river water with suspended solids requires coagulation and filtration.

Practical note
Students should learn to identify the common water sources around them and note visible differences in colour, smell and clarity. Observing seasonal changes, such as lower river levels in summer or muddy water after heavy rains, builds practical awareness of how natural events affect water availability and quality.

Basic composition
A water molecule contains two hydrogen atoms and one oxygen atom, written H2O. The atoms are joined by covalent bonds where electrons are shared. Each hydrogen shares one electron with oxygen, while oxygen shares one electron with each hydrogen.

Electron pairs and molecular shape
Oxygen has six electrons in its outer shell; by sharing two electrons with two hydrogens, it completes an octet. In addition to the two bonding pairs, oxygen has two lone pairs of electrons. These four pairs (two bonding, two lone) repel one another and arrange themselves to minimise repulsion. This gives a bent or V-shaped geometry rather than a straight line. The bond angle between the O–H bonds is about 104.5°, which is slightly less than the ideal tetrahedral angle because lone pairs repel more strongly than bonding pairs.

Polarity of the molecule
Oxygen is more electronegative than hydrogen, so the shared electrons spend more time closer to oxygen. This uneven electron distribution makes the oxygen end slightly negative (δ-) and the hydrogen ends slightly positive (δ+). Because of the bent shape, these partial charges do not cancel out, so the water molecule is polar. The polarity is essential: it explains why water interacts strongly with charged or polar substances and why it is a good solvent for many salts and polar compounds.

Hydrogen bonding and consequences
Because of polarity, the positive hydrogen of one molecule attracts the negative oxygen of a nearby molecule. This attraction, called a hydrogen bond, is weaker than a covalent bond but strong enough to hold molecules together in a network. Each water molecule can form hydrogen bonds with several neighbours. Hydrogen bonding explains many unusual properties of water: relatively high boiling and melting points, high surface tension, cohesion and adhesion, and why ice has a lower density than liquid water due to an open hydrogen-bonded structure.

Visual models
Students should practice drawing the V-shaped water molecule, mark the partial charges (δ+ and δ-) and show dashed lines for hydrogen bonds between molecules. This visual model links small-scale structure to large-scale behaviour of water.

States and transition points
Water commonly exists as solid (ice), liquid (water) and gas (vapour). At standard atmospheric pressure, water freezes at 0°C and boils at 100°C. These temperatures are higher than for many small molecules of similar size because water molecules are strongly attracted to each other by hydrogen bonds.

High specific heat
Water has a high specific heat capacity, meaning it can absorb a large amount of heat with only a small change in temperature. This property helps moderate climates near large water bodies and stabilises temperatures in living organisms, allowing them to maintain steady internal conditions despite temperature changes outside.

Surface tension and cohesion
Hydrogen bonding between water molecules makes them stick together (cohesion), producing a high surface tension. This causes water to form rounded droplets and supports small insects on the surface. Adhesion, the attraction between water and other materials, helps water climb up narrow tubes (capillarity), which is vital for moving water through plant vessels.

Density behaviour and ice
Water shows an unusual density pattern: it reaches maximum density at around 4°C and becomes less dense as it cools further to form ice. Ice has an open crystal structure stabilized by hydrogen bonds that keeps molecules further apart than in liquid water. As a result, ice floats on water. This has important ecological consequences: lakes freeze on the surface in winter, insulating water below and allowing aquatic life to survive.

Solvent properties
Because of its polarity, water dissolves many ionic and polar substances by surrounding and separating ions or molecules. It dissolves salts, sugars and some gases, making it an excellent medium for chemical reactions in nature and in living cells. However, non-polar substances like oils do not dissolve well in water.

Practical observations
Students can observe boiling and evaporation differences, test floating versus sinking of ice, and perform simple experiments showing surface tension (floating a paperclip) or capillary rise in thin tubes or between glass slides to connect these physical properties to everyday experiences.

Amphoteric behaviour
Water behaves both as an acid and a base. It can donate a proton (H+) to form OH- or accept a proton to form H3O+ (often written simply as H+). This ability to act in both roles makes water chemically versatile and central to acid-base chemistry. Pure water self-ionises slightly into equal concentrations of H+ and OH- ions.

Reaction with reactive metals
Some highly reactive metals like sodium and potassium react vigorously with water to produce a metal hydroxide and hydrogen gas. These reactions release heat and can be violent. Less reactive metals such as iron react slowly in the presence of oxygen and water to form rust (hydrated iron oxides), a process that damages metal structures.

Hydrolysis and solvation
Water participates in hydrolysis reactions where it breaks chemical bonds by adding H+ and OH- to different parts of a molecule. Water molecules also solvate ions: the polar water molecules surround and stabilise ions in solution, enabling ionic compounds to dissolve and conduct electricity when dissolved.

Oxidation and reduction roles
Water can act as an oxidant or reductant in different chemical reactions. In electrochemical processes, water on electrolysis decomposes into hydrogen and oxygen gases. In redox reactions in nature, water can donate or accept oxygen or hydrogen atoms depending on the reacting partners.

Electrolysis of water
Electrolysis involves passing an electric current through water containing a small amount of electrolyte. At the cathode hydrogen ions gain electrons to form hydrogen gas, and at the anode hydroxide ions lose electrons to form oxygen gas and water. Electrolysis shows the composition of water and is the basis for producing hydrogen as a fuel in modern technology.

Practical implications
These chemical properties affect how water interacts with materials (corrosion), how it behaves in reactions (hydrolysis), and how we treat water (adding disinfectants, adjusting pH). Understanding these behaviours helps in metal protection, water treatment and many laboratory processes.

Overview of stages
The water cycle, also called the hydrological cycle, describes the continuous movement of water between the Earth's surface and the atmosphere. The major stages are evaporation (liquid water becomes vapour from oceans, lakes and soil), transpiration (plants release water vapour from leaves), condensation (water vapour cools into tiny droplets forming clouds), precipitation (droplets fall as rain, snow or hail), infiltration (water soaks into the ground), and runoff (water flows across the surface into rivers and lakes).

Energy and driving forces
The sun is the primary energy source that powers evaporation and transpiration. Temperature differences, wind and topography determine where vapour moves and where condensation and precipitation occur. Gravity drives runoff and the flow of groundwater toward lower elevations.

Storage and travel
Water is temporarily stored in various reservoirs: oceans contain most of the Earth's water; freshwater is stored in ice caps, glaciers, lakes and aquifers beneath the ground. Groundwater moves slowly through porous rocks and soils, providing water to wells and springs. Surface runoff collects in streams and rivers which return water to the sea, completing the cycle.

Human influences
Human activities alter the natural cycle. Deforestation reduces transpiration and can change local rainfall patterns. Urbanisation increases runoff and reduces infiltration due to paved surfaces, which can raise flood risks. Pollution enters rivers and groundwater, moving contaminants through the cycle to affect wider areas. Water extraction for irrigation and industry can lower groundwater levels, affecting springs and surface water flows.

Importance for water management
Understanding the cycle helps water resource planning: where to build reservoirs, how land use affects recharge, and how pollution can travel. Measures such as protecting catchment areas, planting trees, promoting groundwater recharge and recycling water help maintain sustainable supplies. Students should learn to trace a drop of water through the cycle and consider how local actions can alter its journey.

Pure water
Pure water is chemically H2O without dissolved salts, gases or suspended particles. In practice, absolute purity is rare and is achieved only in laboratories or by special treatments like distillation and deionisation. Pure water is a poor conductor of electricity because it lacks ions that carry charge.

Potable water
Potable water is water safe for drinking. It may contain small amounts of dissolved minerals that are not harmful and in some cases beneficial, but must be free from dangerous microbes, toxic chemicals and unacceptable levels of dissolved solids. Potability is judged by tests for turbidity, taste, odour, pH, presence of bacteria, and concentrations of certain ions like nitrate, fluoride and heavy metals.

Hard water
Hard water contains significant concentrations of calcium (Ca2+) and magnesium (Mg2+) ions, usually from dissolving minerals in rocks such as limestone. Hardness affects household tasks: it reduces soap lathering, forms scum with soap, and causes scale deposits in kettles, boilers and pipes. Hard water can also affect industrial processes and reduce efficiency of heating equipment.

Soft water
Soft water has negligible amounts of calcium and magnesium ions. Rainwater is naturally soft although it may dissolve small amounts of gases like CO2. Soft water allows soap to lather easily, prevents scale formation and is preferred for some industrial uses. However, very soft water (for example water with high sodium from ion-exchange softening) may taste salty and is not always ideal for people on low-sodium diets.

Practical distinctions and uses
For drinking, potable water should balance mineral content: some dissolved minerals are desirable for taste and nutrition, but excessive hardness is inconvenient. For laundry and boilers, soft water reduces soap usage and scale. Knowing which type you have helps choose treatment: boiling helps remove temporary hardness, whereas permanent hardness needs chemical or ion-exchange methods.

What causes hardness?
Water hardness is caused mainly by dissolved salts of calcium and magnesium. These ions come from the dissolution of minerals in rocks and soils as water flows through them. Common sources are calcium carbonate, calcium bicarbonate, magnesium bicarbonate, calcium sulphate and magnesium sulphate. Natural groundwater in areas with limestone or dolomite often has high hardness.

Temporary and permanent hardness
Temporary hardness is usually due to bicarbonates of calcium and magnesium. It is called temporary because it can be removed by boiling: heating causes bicarbonates to decompose into insoluble carbonates which precipitate out. Permanent hardness is caused by non-bicarbonate salts such as sulphates and chlorides; these do not precipitate on boiling and need chemical treatment or ion-exchange to be removed.

Chemical changes on boiling
When calcium bicarbonate is heated, it decomposes to form calcium carbonate (a white precipitate), carbon dioxide and water. This reaction removes the calcium from the water and softens it temporarily. Laboratory tests use these changes to demonstrate the difference between temporary and permanent hardness.

Measuring hardness
Hardness is commonly expressed as milligrams per litre (mg/L) of calcium carbonate (CaCO3), also called parts per million (ppm). Simple household indications include a soap test: hard water produces little lather and forms scum. Laboratory methods provide accurate measurements: titration with EDTA is a standard procedure to determine total hardness precisely. Total Dissolved Solids (TDS) and electrical conductivity measurements give an idea of the amount of dissolved ions but do not distinguish calcium and magnesium specifically.

Effects and significance
Hard water causes scaling in kettles and boilers, reduces efficiency of detergents, and can affect taste. In industrial boilers, scale can cause overheating and damage. Knowing the type and amount of hardness helps select treatment: boiling or adding sodium carbonate for temporary hardness, and ion-exchange or lime-soda processes for permanent hardness on a larger scale.

Boiling to remove temporary hardness
Boiling is the simplest household method to remove temporary hardness caused by bicarbonates. Heating water causes calcium and magnesium bicarbonates to decompose into their carbonate forms which are insoluble and precipitate as solid particles. These can be removed by filtration. Boiling is effective for small amounts of water and is easy to perform at home, though it uses energy and cannot remove permanent hardness.

Using washing soda (sodium carbonate)
Adding sodium carbonate to hard water converts soluble calcium and magnesium salts into insoluble carbonates which precipitate and can be filtered out. This reaction is used in laundry and cleaning: washing soda helps remove hardness ions so soap can lather better. The chemical reactions produce soluble sodium salts in the water and insoluble carbonate precipitates, simplifying removal. Care must be taken to add correct amounts and remove precipitates.

Ion-exchange softening
Ion-exchange softening is commonly used in domestic softeners and water treatment plants. Water is passed through a bed of resin beads charged with sodium ions. The resin exchanges its sodium ions for calcium and magnesium ions in the water, thus removing the hardness. The softener must be regenerated periodically by flushing with concentrated sodium chloride solution, which restores sodium ions on the resin and washes away displaced calcium and magnesium. Ion-exchange works for both temporary and permanent hardness but increases sodium content of water.

Lime-soda and other large-scale methods
For large volumes in municipal treatment, lime (calcium hydroxide) and soda (sodium carbonate) processes precipitate hardness-causing ions as carbonates and hydroxides. Chemical precipitation, settling and filtration are used in sequence. Reverse osmosis and demineralisation are more intensive options that remove most dissolved ions and produce very soft or pure water, used where high purity is needed.

Choosing a method
Choice depends on the type of hardness, volume of water and cost. Boiling and washing soda are practical for households, while ion-exchange and lime-soda are suited for continuous or industrial treatment. Understanding how each method works helps pick the right approach for domestic or industrial needs.

Kinds of impurities
Natural and human activities introduce different impurities into water. Suspended solids such as silt, clay and organic debris make water cloudy. Dissolved inorganic salts like calcium, magnesium, sodium, chloride and sulphate change hardness and taste. Dissolved gases (oxygen, carbon dioxide) and organic matter (decomposing plants) affect odour and biological activity. Microorganisms including bacteria, viruses and protozoa cause diseases. Industrial chemicals, pesticides and heavy metals are harmful at low concentrations. Knowing the type of impurity guides suitable treatment.

Simple home and school tests
Many basic checks can indicate water quality. Visual inspection for colour and turbidity is the simplest: suspended particles scatter light and make water cloudy. Smell and taste tests (carefully) can reveal organic contamination or high dissolved salts; however, absence of taste or smell does not mean safety. pH paper or universal indicator paper shows acidity or alkalinity. A conductivity or TDS meter indicates total dissolved solids and gives a quick idea of dissolved ions. These are useful screening tools but not definitive for contamination.

Specific tests for hardness and microbes
Hard water is detected by soap tests: shaking measured volumes of water and soap solution shows reduced lather and scum formation in hard water. Boiling tests demonstrate temporary hardness if precipitate forms on heating. Microbial contamination requires laboratory methods: presence-absence tests, colony counts or membrane filtration detect coliform bacteria as indicators of faecal pollution. Simple household disinfectants cannot replace proper laboratory testing for pathogens.

Chemical contaminants and indicators
Some chemicals need specific tests: nitrate testing kits for fertiliser pollution, fluoride test strips, and kits for heavy metals like lead. Colourimetric test kits and portable sensors allow field testing by public health workers. Advanced laboratory analysis (spectrometry, chromatography) is used when exact concentrations must be known.

Interpreting results and action
Screening tests help decide immediate actions: if turbidity is high, settle and filter before disinfection; if microbial contamination is suspected, boil or chlorinate water. Knowing the local sources and likely pollutants helps choose the right tests and treatments and prevents relying on a single test for safety.

Overview and aims
Water purification aims to remove suspended solids, reduce dissolved salts, and kill or remove harmful microorganisms and chemical pollutants to produce water suitable for its intended use. Methods range from simple household steps to complex industrial and municipal systems. Often a combination of processes is used to achieve safe drinking water.

Sedimentation and coagulation
Coagulation adds substances such as alum to make very small suspended particles clump together into larger flocs. These settle faster in sedimentation tanks. Sedimentation is simple: water is stored undisturbed allowing heavy particles to settle to the bottom. These steps reduce turbidity and improve the efficiency of later filtration.

Filtration
Filtration removes suspended particles by passing water through porous media like sand, gravel and activated carbon. Slow sand filters rely on a microbial layer that forms on the top to remove pathogens. Activated carbon adsorbs organic molecules and improves taste and odour. Filtration is essential after sedimentation and before disinfection in many treatment plants.

Disinfection and boiling
Boiling is a reliable household method to kill disease-causing microorganisms. For larger supplies, disinfectants such as chlorine are used; chlorine is effective and leaves a residual that protects water in distribution systems. Correct dosing is important to avoid harmful by-products. Other disinfectants like chloramines or UV treatment are used depending on circumstances.

Distillation, reverse osmosis and advanced methods
Distillation boils water and condenses steam, producing water free from most dissolved salts and non-volatile impurities. Reverse osmosis (RO) forces water through a membrane that allows water molecules through but blocks many dissolved ions and molecules. Both require energy and maintenance but produce high-purity water. These methods are used where very low TDS or removal of specific contaminants is required.

Practical household treatments
For homes without central treatment, a practical sequence is settling, coarse filtration through cloth or sand, boiling or chlorination and safe storage in clean covered containers. Knowing the limits of each method helps: for example, boiling kills microbes but does not remove chemical pollutants or dissolved salts, so additional steps are needed when those are present.

Characteristics of potable water
Potable water must be free from harmful pathogens, have acceptable taste and odour, and contain chemical constituents within safe limits. Important parameters include microbial counts (absence of faecal coliforms), pH (usually near neutral), turbidity (low), Total Dissolved Solids (TDS) within recommended limits, and controlled levels of specific ions such as nitrate, fluoride and heavy metals. Standards vary by country but aim to protect public health.

Household practices to keep water safe
At home, keep water containers clean and covered to prevent recontamination. Use a clean ladle for drawing water rather than dipping cups. Regularly clean and inspect storage tanks and cisterns to remove sediment and prevent mosquito breeding. If using well water, ensure well head protection and avoid placing latrines or drainage near the source.

Safe handling and storage
After treatment (boiling or chlorination), store water in clean, covered containers and use within a recommended period. Keep storage at a cool place and avoid touching the inside of the containers. For long term storage, maintaining a residual disinfectant (as with chlorine) helps prevent regrowth of microbes.

Emergency and routine checks
In emergencies or where water quality is doubtful, boil water for at least 1–3 minutes (longer at high altitudes) to kill pathogens. Chlorination with correct dosage can disinfect larger volumes. Routine household checks include observing colour, smell and turbidity, testing pH with indicator paper, and using a TDS meter if available. For serious concerns, send samples to a laboratory for microbial and chemical analysis.

Community responsibilities
Safe potable water depends on proper sanitation and sewage management, regular maintenance of supply systems, and timely action on pollution. Households should report leaks and contamination, participate in protecting catchment areas, and follow public health advisories during contamination events. Simple personal and community actions help ensure long-term access to safe drinking water.

Sources and types of pollution
Water pollution arises when harmful substances enter water bodies. Common sources include untreated or partially treated sewage, industrial effluents containing chemicals and heavy metals, agricultural runoff carrying fertilisers and pesticides, oil spills, and solid wastes such as plastics. These pollutants can be biological (pathogens), chemical (toxic compounds) or physical (suspended solids), and each affects ecosystems and human health differently.

Effects on ecosystems and people
Pollution reduces oxygen levels in water, harming fish and other aquatic life. Nutrient pollution from fertilisers causes eutrophication, where algal blooms block sunlight and consume oxygen, creating dead zones. Toxic chemicals accumulate in food chains and can cause long-term health problems in humans and wildlife. Polluted water used for drinking or irrigation spreads disease and can contaminate crops.

Conservation and protection measures
Conserving water reduces pressure on available supplies. Measures include fixing leaks, using water-efficient appliances, adopting water-saving habits, rainwater harvesting and recycling greywater for non-drinking uses. Protecting catchment areas, preserving wetlands and forests, and maintaining natural vegetation help filter water, recharge groundwater and reduce erosion and sedimentation.

Pollution control and treatment
Sewage treatment plants remove solids and reduce biological contamination before discharge. Industries must treat effluents to remove harmful chemicals. Constructed wetlands and buffer strips near fields reduce runoff. Regulations and community monitoring enforce discharge limits and encourage cleaner production methods. Public awareness and responsible waste disposal reduce litter and chemical contamination.

Local actions with global impact
Small actions at home and in the community—proper disposal of chemicals, reduced use of pesticides, avoiding dumping into drains—help protect rivers and groundwater. Schools can run awareness campaigns, organise clean-up drives, and teach water-wise habits. Conservation of water resources and preventing pollution contribute to healthier communities and sustainable use of water resources for future generations.

Safe and useful observations
Simple experiments help students connect theory to real observations. Safety is important: use small quantities, adult supervision for heating and electricity, and follow instructions carefully. Record observations, note times and temperatures where relevant, and clean equipment after use.

Boiling and evaporation
Observe a beaker of water as it is heated. Note the formation of small bubbles at lower temperatures (dissolved gas being released) and vigorous bubbling when boiling begins. Compare boiling with evaporation at room temperature by placing a small measured volume on a watch glass and timing how long it takes to evaporate. Discuss why evaporation occurs at any temperature while boiling requires reaching the boiling point.

Testing hardness with soap
Perform a soap test: add a standard volume of soap solution to measured amounts of water samples in test tubes and shake. Hard water produces less lather and more scum compared with distilled water. For temporary hardness, boil a portion and repeat the test to observe improvement in lather after boiling due to precipitation of bicarbonates.

Filtration and sedimentation
Mix soil in water and let it stand to observe sedimentation as solids settle. Then pour the cloudy water through folded cloth or a simple sand filter made from layers of gravel and sand in a bottle. Observe clarity improvement and discuss what each layer removes (large particles, fine particles, organic matter). This models how treatment plants prepare water for disinfection.

Electrolysis demonstration (with supervision)
Set up electrolysis of water with a small battery, graphite or inert electrodes and a pinch of electrolyte (e.g., salt) under careful supervision. Collect gases at the electrodes in small test tubes and test which gas is hydrogen (burning a small sample produces a squeaky pop). Explain that the experiment decomposes water into hydrogen and oxygen and relates to chemical composition.

Recording and linking theory
Students should draw diagrams of setups, note observations and link them to concepts such as hydrogen bonding, evaporation, solubility and hardness. Simple calculations like evaporation rate or lather height can make experiments quantitative and reinforce measurement skills.

Domestic uses and hygiene
At home, water is used for drinking, cooking, bathing, washing clothes, cleaning and flushing toilets. Clean water is essential for personal hygiene to prevent diseases. Households use water also for gardening, washing vehicles and other chores. Efficient and careful use of water at home reduces bills and conserves a precious resource.

Agricultural importance
Agriculture consumes much of the world’s freshwater supply. Crops need water for germination, growth and nutrient transport. Irrigation systems—flooding, sprinkler and drip irrigation—deliver water to fields. Efficient irrigation like drip systems reduce water loss and increase yields by supplying water directly to plant roots.

Industrial and municipal uses
Industries use water for cooling machinery, processing materials, cleaning and as a raw material in many chemical processes. Municipal supplies provide water for firefighting, street cleaning and public services. Industrial wastewater must be treated before being returned to the environment to prevent pollution and protect downstream water users.

Ecological and recreational roles
Water bodies such as rivers, lakes and wetlands support biodiversity, provide habitat for fish and birds, and support food chains. People use water for recreation—swimming, boating and fishing—which depends on good water quality. Wetlands act as natural filters and flood buffers and are important to conserve.

Valuing and saving water
Understanding how much water different activities use helps students adopt conservation habits: turn taps off while brushing teeth, repair leaks, collect rainwater for gardening and use efficient appliances. Schools and communities can promote water-saving practices and teach the value of water so that future generations inherit sustainable resources. Small everyday changes add up to significant conservation at community and national levels.