Water Revision Notes

Water detection is essential in many chemical processes and practical applications. Two common chemical tests use anhydrous cobalt(II) chloride and anhydrous copper(II) sulfate as indicators to confirm the presence of water. These tests rely on the distinct colour changes that occur when these compounds absorb water molecules and form hydrated compounds.

• Anhydrous means the compound contains no water molecules in its crystal structure.
• Hydrated compounds contain water molecules chemically bound within their crystal lattice.
• Colour changes occur due to changes in the coordination environment of the metal ions when water molecules bind.

Anhydrous cobalt(II) chloride is a blue crystalline solid. When it comes into contact with water, it forms hydrated cobalt(II) chloride, which is pink. This colour change is a clear visual indicator of moisture or water presence.

The chemical reaction can be represented as: CoCl₂ (blue, anhydrous) + 6H₂O → CoCl₂·6H₂O (pink, hydrated)

This reaction is reversible; heating the hydrated form drives off water, returning it to the blue anhydrous form.

• Initial colour: blue (anhydrous CoCl₂)
• Colour after hydration: pink (CoCl₂·6H₂O)
• Test method: expose cobalt(II) chloride paper or solid to suspected water source and observe colour change
• Safety: cobalt compounds are toxic; handle with care and avoid ingestion or inhalation

Anhydrous copper(II) sulfate is a white or pale grey powder. Upon contact with water, it forms hydrated copper(II) sulfate pentahydrate, which is bright blue.

The chemical reaction is: CuSO₄ (white, anhydrous) + 5H₂O → CuSO₄·5H₂O (blue, hydrated)

This test is commonly used to detect water in various samples, including hydrated salts or moisture in gases.

• Initial colour: white (anhydrous CuSO₄)
• Colour after hydration: blue (CuSO₄·5H₂O)
• Test method: add anhydrous copper(II) sulfate powder to the sample and observe any colour change
• Safety: copper compounds can be harmful; use gloves and avoid inhaling dust

To perform these tests, place a small amount of the anhydrous compound on a watch glass or in a test tube. Add the sample suspected to contain water or expose the compound to moist air.

Observe the colour change carefully: cobalt(II) chloride changes from blue to pink, and copper(II) sulfate changes from white to blue.

Repeat the test to confirm results and ensure reliability. Record observations precisely, noting the time taken for the colour change and any partial changes indicating moisture rather than liquid water.

• Use clean, dry apparatus to avoid contamination.
• Control variables such as temperature and sample amount for consistent results.
• Repeat tests to confirm presence of water.
• Dispose of chemicals according to safety guidelines.

Both tests are qualitative and rely on distinct colour changes to indicate water presence. Cobalt(II) chloride is often used in humidity indicators, while copper(II) sulfate is widely used in laboratory water detection.

In exams, students should be able to describe the colour changes, write balanced equations for hydration, and explain the practical steps involved.

Understanding the reversible nature of these reactions is useful for explaining drying and rehydration processes.

• Cobalt(II) chloride: blue to pink; forms hexahydrate
• Copper(II) sulfate: white to blue; forms pentahydrate
• Both tests confirm water presence visually and are simple to perform
• Balanced equations are essential for full marks

Pure water has a melting point of exactly 0°C and a boiling point of exactly 100°C at standard atmospheric pressure (1 atm or 101 kPa). These fixed temperatures are physical constants for pure water and result from the energy required to change its state: melting (solid to liquid) and boiling (liquid to gas).

The melting point is the temperature at which ice and liquid water coexist in equilibrium. The boiling point is the temperature at which water vapour pressure equals atmospheric pressure, causing rapid vaporisation throughout the liquid.

Any deviation from these temperatures suggests the presence of dissolved substances or impurities, which alter the physical properties of water.

• Melting point of pure water = 0°C at 1 atm.
• Boiling point of pure water = 100°C at 1 atm.
• Melting and boiling points are characteristic physical constants for pure substances.

Impurities dissolved in water disrupt the regular arrangement of water molecules, affecting the energy required for phase changes. This causes melting point depression and boiling point elevation.

Melting point depression means the water freezes at a temperature below 0°C. This is because impurities interfere with ice crystal formation, requiring lower temperatures to solidify.

Boiling point elevation means the water boils at a temperature above 100°C. The dissolved substances reduce the vapour pressure of the liquid, so a higher temperature is needed to reach atmospheric pressure and boil the water.

• Impurities lower the melting point below 0°C (melting point depression).
• Impurities raise the boiling point above 100°C (boiling point elevation).
• The magnitude of change depends on the amount and nature of impurities.

Apparatus needed includes a small test tube or capillary tube, a thermometer or temperature probe, a cooling bath (ice-salt mixture), and a water sample.

Procedure: Place a small amount of the water sample in the test tube or capillary tube. Cool the sample in an ice-salt bath to ensure it is fully solidified (below 0°C). Then slowly warm the sample while monitoring the temperature carefully.

Observation: Record the temperature at which the ice begins to melt and remains constant during melting. This temperature is the melting point of the sample. Repeat the measurement several times to ensure accuracy and calculate an average value.

• Use a thermometer accurate to ±0.1°C.
• Cool sample below 0°C before warming.
• Record temperature when ice starts to melt and remains steady.
• Repeat measurements for reliability.
• Compare measured melting point to 0°C.

Apparatus includes a boiling tube or conical flask, a thermometer or temperature probe, a heat source (Bunsen burner or electric heater), and the water sample.

Procedure: Place the water sample in the boiling tube and insert the thermometer so the bulb is immersed in the liquid but not touching the container sides or bottom. Heat the water gently and steadily.

Observation: Record the temperature at which the water starts to boil and remains constant during vigorous bubbling. This temperature is the boiling point of the sample. Repeat the measurement to improve accuracy and calculate an average.

• Ensure thermometer bulb is immersed in liquid, not touching container.
• Heat water gently to avoid superheating.
• Record temperature when boiling starts and remains constant.
• Repeat measurements and average results.
• Compare measured boiling point to 100°C.

If the measured melting point is close to 0°C and the boiling point close to 100°C (within experimental error), the water sample is considered pure or very nearly pure.

If the melting point is significantly below 0°C and/or the boiling point significantly above 100°C, impurities are present. The greater the deviation, the higher the impurity level.

Common impurities include dissolved salts, minerals, or organic substances, which cause freezing point depression and boiling point elevation due to colligative properties of solutions (depend on solute particle concentration).

• Melting point near 0°C and boiling point near 100°C indicate pure water.
• Lower melting point and higher boiling point indicate impurities.
• Degree of deviation correlates with impurity concentration.
• Use repeated measurements and controlled conditions for reliable assessment.

Tap water is sourced from rivers, lakes, or underground reservoirs and treated to make it safe for drinking. However, it contains dissolved minerals such as calcium (Ca²⁺), magnesium (Mg²⁺), and chloride (Cl⁻) ions, as well as trace amounts of other chemicals like fluoride, chlorine compounds, and sometimes organic matter. These dissolved substances are called impurities because they are not pure water (H₂O).

The presence of these impurities means tap water is not chemically pure and can vary in composition depending on the source and treatment processes. For example, hard water contains higher concentrations of calcium and magnesium ions, which can affect chemical reactions.

• Tap water contains dissolved minerals and chemicals.
• Impurities vary with source and treatment.
• Common ions include Ca²⁺, Mg²⁺, Cl⁻, and residual chlorine.

Distilled water is produced by boiling water to create steam and then condensing the steam back into liquid water. This process removes most dissolved solids, minerals, and many chemical impurities because these substances do not vaporise with the water.

The result is water that is nearly pure H₂O, free from dissolved salts and many other contaminants. Distilled water is colourless, odourless, and tasteless, and its purity is essential for sensitive chemical experiments where impurities could cause unwanted side reactions or inaccurate results.

• Distillation removes dissolved solids and many impurities.
• Steam is condensed to yield pure water.
• Distilled water is nearly 100% pure H₂O.

Impurities in tap water can interfere with chemical reactions in several ways. For example, dissolved ions may react with reagents, forming precipitates or altering reaction pathways. This can lead to incorrect observations or unexpected products.

For instance, if tap water containing calcium ions is used in a reaction to prepare a soluble salt, calcium may form insoluble compounds, contaminating the product. Additionally, impurities can affect the pH or conductivity of solutions, influencing reaction rates and equilibria.

Therefore, using impure water can reduce the accuracy and reproducibility of experimental results.

• Impurities may react with chemicals, causing side reactions.
• They can form precipitates or alter product purity.
• Impurities affect pH and reaction conditions.

Using distilled water minimises the risk of contamination and ensures that chemical reactions proceed as expected. This is crucial for obtaining reliable, reproducible results, especially in quantitative experiments such as titrations or preparing standard solutions.

Distilled water prevents interference from extraneous ions or substances, allowing chemists to attribute observed changes solely to the intended reactants.

In practical work, distilled water is used for dissolving reagents, rinsing apparatus, and preparing solutions to maintain experimental integrity.

• Distilled water ensures reliable and reproducible results.
• Prevents contamination and interference in reactions.
• Essential for accurate quantitative and qualitative experiments.

Always use distilled water rather than tap water in chemical experiments to avoid errors caused by impurities. When preparing solutions, rinse glassware with distilled water to prevent contamination.

If tap water must be used, be aware of its potential impurities and how they might affect your results. However, distilled water remains the preferred choice for precision and accuracy in chemistry.

Remember that distilled water is not sterile; for microbiological work, further purification may be necessary, but for general chemistry, its chemical purity is the key advantage.

• Use distilled water for all chemical preparations and rinsing.
• Avoid tap water to prevent impurity interference.
• Distilled water is chemically pure but not necessarily sterile.

Dissolved oxygen (O₂) is oxygen gas present in water, essential for the respiration of aquatic organisms such as fish and invertebrates. It enters water mainly by diffusion from the atmosphere and as a by-product of photosynthesis by aquatic plants and algae.

The concentration of dissolved oxygen varies with temperature, pressure, and biological activity. Colder water holds more dissolved oxygen than warmer water.

Adequate dissolved oxygen levels are crucial for maintaining healthy aquatic ecosystems. Low oxygen levels (hypoxia) can cause stress or death to aquatic life.

• Dissolved oxygen is O₂ gas dissolved in water.
• Sources include atmospheric diffusion and photosynthesis.
• Essential for aquatic organism respiration.
• Concentration depends on temperature and biological activity.

Natural water often contains metal ions dissolved from rocks and soils. Common metal compounds include calcium carbonate (CaCO₃), magnesium sulfate (MgSO₄), and iron compounds such as iron(III) oxide (Fe₂O₃).

Some metal ions are beneficial in trace amounts, providing essential minerals. However, toxic metals like lead (Pb²⁺), mercury (Hg²⁺), and cadmium (Cd²⁺) can contaminate water through industrial pollution, posing health risks.

Metal ions in water can be present as dissolved ions or as suspended particles, affecting water hardness and colour.

• Metal ions dissolve from natural sources and pollution.
• Examples: Ca²⁺, Mg²⁺, Fe³⁺, Pb²⁺, Hg²⁺.
• Beneficial metals provide minerals; toxic metals cause harm.
• Metal compounds affect water hardness and colour.

Plastics enter water bodies mainly through littering and improper waste disposal. They do not dissolve but break down into microplastics, which can be ingested by aquatic organisms, causing physical harm and chemical contamination.

Sewage contains organic waste and harmful microbes such as bacteria, viruses, and parasites. It enters water through untreated or partially treated wastewater discharge, leading to waterborne diseases and oxygen depletion due to microbial decomposition.

Both plastics and sewage degrade water quality and threaten aquatic life and human health.

• Plastics pollute water as debris and microplastics.
• Sewage contains organic waste and harmful microbes.
• Sewage causes disease and oxygen depletion.
• Both reduce water quality and harm ecosystems.

Nitrates (NO₃⁻) and phosphates (PO₄³⁻) enter water bodies mainly from agricultural fertilisers and detergents.

Excess nitrates and phosphates cause eutrophication, a process where nutrient enrichment leads to excessive algae growth. When algae die and decompose, oxygen is consumed, causing deoxygenation harmful to aquatic life.

Nitrate ions can also contaminate drinking water, posing health risks such as methemoglobinemia (blue baby syndrome) in infants.

• Nitrates and phosphates originate from fertilisers and detergents.
• Cause eutrophication and oxygen depletion in water.
• Excess nitrates pose health risks in drinking water.
• Control of nutrient runoff is important for water quality.

Water naturally contains dissolved oxygen (O₂), which originates from the atmosphere and photosynthesis by aquatic plants and algae. This oxygen is crucial for the survival of aerobic aquatic organisms such as fish, invertebrates, and aerobic bacteria.

Dissolved oxygen enables cellular respiration, a process where organisms use oxygen to convert glucose into energy (ATP), carbon dioxide, and water.

The concentration of dissolved oxygen is affected by temperature, pressure, and the presence of other substances. Cooler water holds more oxygen than warmer water.

• Dissolved oxygen supports aerobic respiration in aquatic organisms.
• Sources include atmospheric diffusion and photosynthesis.
• Essential for maintaining healthy aquatic ecosystems.

Water often contains dissolved metal ions from natural mineral deposits or soil erosion. These metal ions include calcium (Ca²⁺), magnesium (Mg²⁺), iron (Fe²⁺/Fe³⁺), and zinc (Zn²⁺), which are essential minerals for both aquatic life and humans.

Calcium and magnesium ions contribute to water hardness and are important for bone formation and enzyme function in organisms.

Iron ions are vital for oxygen transport in fish and other aquatic animals, as well as for plant chlorophyll synthesis.

• Calcium and magnesium ions contribute to hardness and nutrition.
• Iron ions support oxygen transport and photosynthesis.
• Metal ions are micronutrients necessary for biological processes.

Metal ions in water exist as hydrated cations, for example, Ca²⁺(aq) and Mg²⁺(aq). These ions can react with other ions to form compounds such as calcium carbonate (CaCO₃) and magnesium carbonate (MgCO₃), which may precipitate as scale.

Iron can exist in two oxidation states: Fe²⁺ (ferrous) and Fe³⁺ (ferric). The balance between these forms depends on the water's oxygen content and pH.

The presence of these ions can be detected by specific chemical tests, such as adding sodium hydroxide solution to form characteristic coloured precipitates (e.g., Fe(OH)₃ is reddish-brown).

• Metal ions are hydrated in solution.
• Calcium and magnesium form carbonates contributing to hardness.
• Iron exists as Fe²⁺ or Fe³⁺ depending on conditions.

Dissolved oxygen is vital for aerobic respiration in aquatic organisms; low oxygen levels (hypoxia) can cause fish kills and ecosystem collapse.

Essential minerals from metal ions support growth, reproduction, and metabolic functions in aquatic life and contribute to human dietary mineral intake through drinking water.

Maintaining appropriate levels of these substances is important; deficiencies can harm organisms, while excesses may cause water quality issues.

• Dissolved oxygen prevents hypoxia and supports biodiversity.
• Mineral ions are necessary nutrients for life.
• Balanced concentrations maintain ecosystem and human health.

Dissolved oxygen concentration is measured using probes or chemical titration methods such as the Winkler method.

Concentrations of metal ions can be determined by chemical analysis techniques including atomic absorption spectroscopy or colourimetric tests.

Regular monitoring helps manage water quality to ensure it supports aquatic life and is safe for human consumption.

• Dissolved oxygen measured by probes or Winkler titration.
• Metal ions analysed by spectroscopy or chemical tests.
• Monitoring ensures water remains beneficial and safe.

Certain metals dissolve in water forming toxic metal compounds that can harm aquatic organisms. Examples include lead(II) ions (Pb²⁺), mercury(II) ions (Hg²⁺), and cadmium(II) ions (Cd²⁺). These metal ions can enter water from industrial discharge, mining, or corrosion of pipes.

Toxic metal ions interfere with biological processes in aquatic animals and plants, causing poisoning and death. For example, lead ions can inhibit enzyme activity and damage fish nervous systems.

Metal ions in water are often present as hydrated ions, e.g., Pb²⁺(aq). Their toxicity depends on concentration and exposure time. Water contaminated with toxic metals is unsafe for drinking and aquatic life.

• Common toxic metals: lead (Pb²⁺), mercury (Hg²⁺), cadmium (Cd²⁺).
• Sources: industrial waste, mining runoff, corroded plumbing.
• Toxic effects: enzyme inhibition, nervous system damage, bioaccumulation.

Plastics enter water bodies through littering, improper waste disposal, and breakdown of larger plastic items. They do not dissolve but persist as solid particles or microplastics.

Aquatic animals can ingest plastic particles, mistaking them for food, leading to physical harm such as internal injury, blockage of digestive tracts, and starvation.

Plastics can also entangle animals, restricting movement and causing injury or death. The persistence of plastics in water makes them a long-term hazard to ecosystems.

• Plastics do not dissolve but persist as debris or microplastics.
• Ingestion causes internal injury and starvation in aquatic animals.
• Entanglement restricts movement and can cause death.

Sewage contains human and animal waste, which introduces harmful microbes such as bacteria, viruses, and parasites into water bodies.

These microbes can cause diseases in humans and animals, including cholera, typhoid, and dysentery.

The presence of sewage increases the biological oxygen demand (BOD) as microbes decompose organic matter, reducing oxygen available for aquatic life.

• Sewage introduces pathogenic bacteria, viruses, and parasites.
• Causes waterborne diseases affecting humans and animals.
• Increases biological oxygen demand, reducing dissolved oxygen levels.

Nitrates (NO₃⁻) and phosphates (PO₄³⁻) enter water mainly from fertilisers used in agriculture and detergents.

Excess nitrates and phosphates promote rapid growth of algae and aquatic plants, a process called eutrophication.

Although detailed eutrophication is beyond this scope, the key effect is that decomposition of excess plant material consumes dissolved oxygen, leading to deoxygenation of water and harm to aquatic organisms.

• Sources: fertilisers and detergents.
• Cause excessive growth of algae and aquatic plants.
• Decomposition of excess biomass reduces dissolved oxygen, harming aquatic life.

Sedimentation is the initial stage in treating domestic water supplies. Water from natural sources often contains suspended solids such as sand, soil particles, and organic debris. In sedimentation, water is held in large tanks or reservoirs where it is allowed to stand undisturbed for several hours. Due to gravity, the heavier solid particles settle to the bottom as sediment.

This process reduces the turbidity (cloudiness) of the water and prepares it for further purification steps. Sedimentation is a physical separation method and does not remove dissolved substances or microorganisms.

Key points: Larger solids settle out by gravity; clear water remains above the sediment layer; sediment is periodically removed from the tank.

• Water is held in a tank to allow solids to settle.
• Heavier particles settle at the bottom forming sediment.
• Clearer water is collected from the top for further treatment.

After sedimentation, some smaller suspended particles remain in the water. Filtration is used to remove these finer particles. Water passes through layers of filter media, typically sand, gravel, and sometimes charcoal.

The filter media trap and hold suspended solids by physical straining and adsorption. This process further clarifies the water by removing particles too small to settle by sedimentation alone.

Filtration does not remove dissolved chemicals or microorganisms effectively but significantly improves water clarity and quality before chemical treatment stages.

• Water passes through sand and gravel filters.
• Smaller suspended solids are trapped in the filter media.
• Filtered water is clearer and ready for chemical treatment.

Even after filtration, water may contain dissolved organic compounds or chlorine by-products that cause unpleasant tastes and odours. Activated carbon is used to adsorb these substances.

Activated carbon has a high surface area and porous structure, allowing it to adsorb organic molecules and chlorine compounds from the water. This improves the water's taste and smell, making it more acceptable for drinking.

This process does not remove suspended solids or kill microbes but targets chemical impurities responsible for taste and odour problems.

• Activated carbon adsorbs dissolved organic compounds and chlorine by-products.
• Improves taste and odour of treated water.
• Activated carbon filters are used after sedimentation and filtration.

Chlorination is the final and critical step in domestic water treatment. Chlorine, usually added as chlorine gas (Cl₂), sodium hypochlorite (NaOCl), or calcium hypochlorite (Ca(OCl)₂), is introduced into the water to kill pathogenic microorganisms such as bacteria, viruses, and protozoa.

Chlorine acts as a strong oxidising agent, disrupting microbial cell walls and enzymes, leading to their death. This disinfection process ensures the water is microbiologically safe to drink.

The amount of chlorine added is carefully controlled to be sufficient to kill microbes but low enough to avoid harmful effects or unpleasant taste. Residual chlorine remains in the water to prevent recontamination during distribution.

• Chlorine kills harmful microbes by oxidation.
• Common chlorine sources: Cl₂ gas, NaOCl, Ca(OCl)₂.
• Residual chlorine prevents microbial regrowth in pipes.
• Chlorination ensures microbiological safety of drinking water.