Comprehensive Study Guide: Water, Solutions, and Water Treatment

Chemical Composition, Structure, and Physical Properties of Water

• Chemical Identity & Structure

  • Chemical Formula: $H_2O$

  • Chemical Name: Dihydrogen oxide

  • Molecular Mass: $1 \times 2 + 16 \times 1 = 18\,\text{amu}$

  • Bonding: Water consists of two hydrogen atoms covalently bonded to a single oxygen atom. The oxygen atom carries a partial negative charge, while each hydrogen atom carries a partial positive charge, giving rise to a polar molecular arrangement.

  • Natural Abundance: Over $70\%$ of the Earth's surface is covered by water. However, fresh water accounts for only $2\text{--}5\%$ of the total water on Earth, which includes frozen fresh water located in glaciers and polar ice caps.

  • States of Matter: Water exists in three physical states: solid (ice), liquid (water), and gas (water vapour). It exists naturally in both free and chemically combined states.

• Biological Significance

  • Water is a primary constituent of all living organisms, accounting for approximately $70\%$ of human body weight.

  • It is vital for the continuance and maintenance of plant, animal, and human life, driving basic biochemical and physiological processes.

• Physical Properties of Water

  • Nature: Pure water is clear, transparent, colourless, odourless, and tasteless. Any taste present in ordinary water is due to dissolved gases and solid impurities.

  • Boiling Point:

  • Pure water boils at $100^\circ\text{C}$ under standard atmospheric pressure ($760\,\text{mm Hg}$).

  • Boiling occurs when the vapour pressure of the liquid equals atmospheric pressure. Increased pressure raises the boiling point, whereas decreased pressure lowers it.

  • Application (Pressure Cooker): A pressure cooker increases internal pressure, raising the boiling point of water above $100^\circ\text{C}$, allowing food to cook faster.

  • Hills/High Altitudes: At high altitudes, low atmospheric pressure reduces the boiling point of water (e.g., $70^\circ\text{C}$ at the summit of Mount Everest). Food does not cook thoroughly at lower boiling temperatures, making pressure cookers necessary.

  • Impurities: The presence of dissolved impurities increases the boiling point of water.

  • Freezing Point and Melting Point:

  • Pure water freezes into ice at $0^\circ\text{C}$ under $1\,\text{atm}$ of pressure.

  • Increasing the pressure or adding dissolved impurities decreases the freezing point of water.

  • Density & Expansion:

  • Maximum Density: Water achieves its maximum density of $1\,\text{g\,cm}^{-3}$ (or $1000\,\text{kg\,m}^{-3}$) at $4^\circ\text{C}$, at which point its volume is at a minimum.

  • Freezing Expansion: Water expands upon freezing, such that $92$ volumes of liquid water expand to form $100$ volumes of ice.

  • Relative Density of Ice: Ice has a relative density of $0.92$, enabling it to float on liquid water.

  • Anomalous Expansion of Water:

  • Liquids generally contract continuously when cooled. Water contracts when cooled until it reaches $4^\circ\text{C}$. Below $4^\circ\text{C}$, it anomalously expands until it freezes into ice at $0^\circ\text{C}$.

  • Ecological Significance: In cold climates, surface water freezes first. Because ice is less dense than water and acts as a poor conductor of heat, liquid water at $4^\circ\text{C}$ remains underneath the surface layer of ice, preserving marine and aquatic life.

  • Latent Heat of Fusion of Ice:

  • Definition: The heat energy required to transform a substance from solid to liquid state at its melting point without a change in temperature.

  • Specific Value: $336\,\text{J\,g}^{-1}$ (or $80\,\text{cal\,g}^{-1}$).

  • An identical quantity of heat ($336\,\text{J\,g}^{-1}$) is liberated when $1\,\text{g}$ of water solidifies into $1\,\text{g}$ of ice at $0^\circ\text{C}$.

  • High specific latent heat of solidification prevents lakes and rivers from freezing suddenly during freezing temperatures.

  • Latent Heat of Vaporization of Water:

  • Definition: The energy required to convert water into water vapour at its boiling point without any change in temperature.

  • Specific Value: $2268\,\text{J\,g}^{-1}$ (or $540\,\text{cal\,g}^{-1}$).

  • Body Cooling Mechanism: Evaporation of sweat absorbs heat energy from the human body ($2268\,\text{J\,g}^{-1}$), producing a cooling effect.

  • Condensation Heat: Steam at $100^\circ\text{C}$ releases $2268\,\text{J\,g}^{-1}$ upon condensing into water at $100^\circ\text{C}$, making burns caused by steam far more severe than burns caused by boiling water at the same temperature.

  • Torrential Rain: The sudden condensation of atmospheric water vapour releases massive amounts of latent heat, contributing to severe atmospheric storm violence.

  • Specific Heat Capacity:

  • Definition: The heat energy absorbed by $1\,\text{g}$ of a substance to raise its temperature by $1^\circ\text{C}$.

  • Specific Value: Water absorbs $4.2\,\text{J}$ (or $1\,\text{cal}$) per gram per degree Celsius ($4.2\,\text{J\,g}^{-1}\,^\circ\text{C}^{-1}$).

  • Applications: Water is used as a coolant in automobile radiators and desert coolers. Large bodies of water moderate climate, making surrounding regions cooler in summer and warmer in winter, driving land and sea breezes.

Universal Solvent Properties and Dissolved Substances

• Solvent Characteristics

  • Solvent: A medium capable of dissolving other substances.

  • Water is termed the universal solvent because it dissolves a wider variety of substances (solids, liquids, and gases) than any other known liquid.

  • High Dielectric Constant: Water possesses a high dielectric constant, which weakens the electrostatic attractive forces between positive and negative ions of inorganic/electrovalent compounds, enabling their dissolution.

  • Polar Nature: Partial negative charge on oxygen and partial positive charges on hydrogens allow water molecules to interact with and bind to ionic and polar compounds.

  • Slight Solubilization of Glass: Materials generally regarded as insoluble dissolve in trace amounts. Water stored in sealed glass bottles over prolonged periods dissolves trace glass, producing visible etchings on the inner surface.

• Experimental Demonstration of Dissolved Solids

  • Procedure: Tap water is placed on a clean watch glass positioned over a beaker of boiling water. The water on the watch glass is evaporated to dryness.

  • Observation: Concentric rings of solid residue remain on the watch glass after total liquid evaporation.

  • Comparison: Tap water, river water, and well water contain dissolved solids and form concentric rings. Rainwater and distilled water contain no dissolved solids and leave no residue.

• Biological and Practical Roles of Dissolved Salts

  • Essential Minerals: Provide necessary minerals required for plant growth and human health.

  • Palatability: Dissolved mineral salts give drinking water its pleasant taste.

Air Dissolved in Water and its Biological Significance

• Composition of Air Dissolved in Water

  • Natural water bodies contain dissolved atmospheric gases. Oxygen is significantly more soluble in water than nitrogen.

  • Dissolved Air Composition: $33\%$ oxygen, $66\%$ nitrogen, and $1\%$ carbon dioxide (compared to ordinary atmospheric air which contains $21\%$ oxygen and $78\%$ nitrogen).

  • Expressing Gas Solubility: Quantified as the volume of gas dissolved per unit volume of water at a given temperature (e.g., $0.02$ volume of $H_2$ per unit volume of water at $0^\circ\text{C}$; $506$ volumes of $HCl$ gas per unit volume of water at $0^\circ\text{C}$).

  • Expulsion: Boiling water completely expels dissolved gases. Consequently, boiled or distilled water tastes flat.

• Experimental Collection of Dissolved Gases

  • Procedure: A round-bottom flask completely filled with tap water is fitted with a delivery tube leading beneath a water-filled graduated tube inverted over a trough. Heating the flask drives off dissolved gas bubbles, which collect in the graduated tube by downward displacement of water.

• Biological Significance of Dissolved Gases

  • Respiration of Aquatic Organisms: Marine life utilizes oxygen dissolved in water for respiration ($1\,\text{dm}^3$ or $1\,\text{L}$ of water contains approximately $40\,\text{cm}^3$ of dissolved oxygen).

  • Photosynthesis: Aquatic plants consume dissolved carbon dioxide for photosynthesis:     $6CO_2(g) + 12H_2O(l) \xrightarrow{\text{sunlight, chlorophyll}} C_6H_{12}O_6(aq) + 6O_2(g) + 6H_2O(l)$

  • Shell Formation: Dissolved carbon dioxide reacts with underwater limestone ($CaCO_3$) to produce soluble calcium bicarbonate:     $CaCO_3(s) + CO_2(g) + H_2O(l) \rightarrow Ca(HCO_3)_2(aq)$     Marine organisms (snails, oysters) extract calcium carbonate from calcium bicarbonate to build their protective shells.

Classification of Matter: Mixtures, True Solutions, Suspensions, and Colloids

• Mixtures

  • Combination of two or more substances (elements or compounds) physically mixed in any arbitrary proportion without chemical bonding.

  • Homogeneous Mixture: Constituents are uniformly distributed with no visible boundaries of separation (e.g., air, sugar solution).

  • Heterogeneous Mixture: Constituents are non-uniformly distributed with clear physical boundaries of separation (e.g., oil and water, salt and sand, sand and water).

• True Solutions

  • Homogeneous mixture composed of a solute dissolved in a solvent.

  • Components: Solute (minor component being dissolved) and Solvent (major component acting as the dissolving medium).

  • Terminology based on component count: Binary ($2$ components), Ternary ($3$ components), Quaternary ($4$ components).

  • Phase Rule: If components exist in different physical states, the solvent is the component in the same physical state as the resulting solution. If both components share the same physical state, the major component is the solvent.

  • Solid Solutions: Homogeneous mixtures of solids (e.g., brass containing $70\%$ copper and $30\%$ zinc; jewellery gold containing gold mixed with copper or silver).

  • Particle Size: Solute particle size is less than $1\,\text{nm}$ ($< 10^{-9}\,\text{m}$).

  • Characteristics:

  • Homogeneous, clear, and transparent.

  • Solute particles do not settle under gravity.

  • Cannot be separated by ordinary filtration or ultrafiltration.

  • Does not scatter light (no Tyndall effect).

• Non-Aqueous Solutions

  • Solutions in which solvents other than water are used.

  • Non-aqueous solvents include alcohol, petrol, ether, benzene, carbon disulphide ($CS_2$), liquid ammonia, acetone, turpentine oil, and carbon tetrachloride ($CCl_4$).

  • Common Solute-Solvent Combinations:

  • Chlorophyll dissolved in Methylated spirit

  • Iodine dissolved in Alcohol (termed Tincture of Iodine, used as an antiseptic)

  • Paraffin wax dissolved in Turpentine oil

  • Sulphur dissolved in Carbon disulphide

  • Paint pigments dissolved in Turpentine oil

  • Phosphorus dissolved in Carbon disulphide

  • Nail polish pigments dissolved in Acetone

  • Grease oils dissolved in Petrol

  • Bromine dissolved in Carbon tetrachloride

• Suspensions

  • Heterogeneous mixture containing insoluble solid particles spread throughout a liquid without dissolving.

  • Particle Size: Solute particles are greater than $1000\,\text{nm}$ ($> 10^{-6}\,\text{m}$).

  • Example: Bubbling carbon dioxide through lime water ($Ca(OH)_2$) produces a milky precipitate of calcium carbonate:     $Ca(OH)_2(aq) + CO_2(g) \rightarrow CaCO_3(s) + H_2O(l)$

  • Other Examples: Muddy water, chalk-water mixture, paints, liquid medicines.

  • Characteristics:

  • Heterogeneous and opaque.

  • Solute particles visible to the naked eye.

  • Particles settle at the bottom under gravity when left undisturbed.

  • Can be separated easily using ordinary filter paper.

Types of Colloidal Systems, Properties, and Applications

• Colloidal Solutions (Sols)

  • Heterogeneous system composed of microscopically insoluble particles dispersed throughout a dispersion medium.

  • Particle Size: Ranges from $1\,\text{nm}$ to $1000\,\text{nm}$ ($10^{-9}\,\text{m}$ to $10^{-6}\,\text{m}$).

  • Consists of two phases: Dispersed Phase (component present in smaller proportion) and Dispersion Medium (component present in larger proportion).

• Types of Colloidal Solutions

  • Emulsion: Dispersed Phase = Liquid | Dispersion Medium = Liquid | Examples: Cod liver oil, liquid tonics, milk.

  • Gel: Dispersed Phase = Liquid | Dispersion Medium = Solid | Examples: Butter, cheese, jelly, curd.

  • Liquid Aerosol: Dispersed Phase = Liquid | Dispersion Medium = Gas | Examples: Cloud, fog, mist.

  • Foam: Dispersed Phase = Gas | Dispersion Medium = Liquid | Examples: Soda water, soap water.

  • Solid Foam: Dispersed Phase = Gas | Dispersion Medium = Solid | Examples: Rubber, bread, pumice stone.

  • Solid Sol: Dispersed Phase = Solid | Dispersion Medium = Solid | Examples: Ruby glass, alloys.

  • Sol: Dispersed Phase = Solid | Dispersion Medium = Liquid | Examples: Paints, gold sol, starch in water.

  • Solid Aerosol: Dispersed Phase = Solid | Dispersion Medium = Gas | Examples: Dust storm, smoke, fumes, vehicular exhaust.

• Key Phenomenon & Properties of Colloids

  • Tyndall Effect: The scattering of a beam of light by colloidal particles, rendering the light path visible. Observed when light enters a dark room through a small hole (scattered by dust/smoke) or when sunlight passes through a forest canopy (scattered by mist droplets).

  • Brownian Motion: Continuous, rapid zig-zag movement of colloidal particles within the dispersion medium, first observed by Robert Brown in 1828. Prevents particles from settling.

  • Electrical Charge: Colloidal particles carry similar electric charges, generating electrostatic repulsion that maintains dispersion stability.

  • Separation: Colloidal particles pass through standard filter paper but are retained by semipermeable ultrafilters (e.g., animal membranes). They can be separated using centrifugation.

  • Emulsions & Emulsifiers: Liquid-liquid colloidal systems (moisturizing cream, cold cream, milk). Emulsifiers (such as proteins) stabilize immiscible liquid mixtures.

• Applications of Colloids

  • Coagulation: The process of neutralizing charges on colloidal particles by adding oppositely charged ions, leading to particle aggregation and precipitation.

  • Blood Clotting with Alum: Blood is a colloidal solution containing charged haemoglobin particles. Applying alum (a hydrated double sulphate salt) supplies ions that neutralize particle charges, coagulating the blood and stopping bleeding.

  • Cleansing Action of Soap: Soap forms an emulsion between water and oily dirt, entrapping dirt particles in micelles.

  • Blue Colour of Sky: Fine dust particles, air molecules, and water vapour act as colloids, scattering shorter blue wavelengths of sunlight more intensely than longer wavelengths.

  • Delta Formation: River water carries suspended colloidal particles (clay, silt). Upon contact with saline sea water containing dissolved electrolytes, charge neutralization occurs, leading to coagulation and accumulation of sediment at the river mouth.

• Comprehensive Comparison Table

  • Nature: Suspension = Heterogeneous | Colloidal Solution = Heterogeneous (appears homogeneous) | True Solution = Homogeneous

  • Particle Size: Suspension = $> 1000\,\text{nm}$ ($> 10^{-6}\,\text{m}$) | Colloidal Solution = $1\text{--}1000\,\text{nm}$ ($10^{-9}\text{--}10^{-6}\,\text{m}$) | True Solution = $< 1\,\text{nm}$ ($< 10^{-9}\,\text{m}$)

  • Ordinary Filtration: Suspension = Separable | Colloidal Solution = Not separable | True Solution = Not separable

  • Ultrafiltration: Suspension = Separable | Colloidal Solution = Separable | True Solution = Not separable

  • Settling Behavior: Suspension = Settles under gravity | Colloidal Solution = Settles only on centrifugation | True Solution = Does not settle

  • Appearance: Suspension = Opaque | Colloidal Solution = Translucent/Generally transparent | True Solution = Clear & Transparent

  • Tyndall Effect: Suspension = Exhibits | Colloidal Solution = Exhibits | True Solution = Does not exhibit

  • Particle Diffusion: Suspension = Does not diffuse | Colloidal Solution = Diffuses slowly | True Solution = Diffuses rapidly

  • Brownian Motion: Suspension = May show | Colloidal Solution = Shows distinct motion | True Solution = Negligible

Solution Concentration, Saturated Solutions, and Solubility

• Saturation States

  • Unsaturated Solution: A solution that can dissolve additional solute at a given temperature.

  • Saturated Solution: A solution that cannot dissolve any more solute at a specific temperature.

  • Preparation: Dissolving nitre ($KNO_3$) in $100\,\text{g}$ of water gram-by-gram with continuous stirring until undissolved solute remains at the bottom.

  • Effect of Cooling: Lowering the temperature of a saturated solution decreases solubility, causing excess solute to separate as crystals.

  • Effect of Heating: Heating a saturated solution makes it unsaturated, enabling it to dissolve more solute.

  • Supersaturated Solution: A solution that holds more dissolved solute than theoretically possible at a given temperature.

  • Preparation: Dissolving excess solute in a heated saturated solution and cooling it slowly without mechanical disturbance.

  • Metastability: Highly unstable; mechanical shaking, stirring, scratching the vessel wall, or adding a seed crystal precipitates excess solute, reverting it to a saturated solution.

• Expressing Concentration

  • Mass Percent (used for solid solute in liquid solvent):     $\text{Mass Percent} = \frac{\text{Mass of Solute}}{\text{Mass of Solution}} \times 100 = \frac{\text{Mass of Solute}}{\text{Mass of Solute} + \text{Mass of Solvent}} \times 100$

  • Example: $10\,\text{g}$ $NaCl$ in $90\,\text{g}$ water gives a $\frac{10}{10 + 90} \times 100 = 10\%$ mass solution.

  • Example: $50\,\text{g}$ sugar in $2.45\,\text{kg}$ ($2450\,\text{g}$) water yields $\frac{50}{50 + 2450} \times 100 = 2\%$ mass solution.

  • Volume Percent (used for liquid solute in liquid solvent):     $\text{Volume Percent} = \frac{\text{Volume of Solute}}{\text{Volume of Solution}} \times 100 = \frac{\text{Volume of Solute}}{\text{Volume of Solute} + \text{Volume of Solvent}} \times 100$

  • Example: $2.5\,\text{L}$ alcohol in $10\,\text{L}$ total solution yields $\frac{2.5}{10} \times 100 = 25\%$ volume solution.

• Definition and Quantitative Expressions of Solubility

  • Solubility: The maximum mass of solute in grams that dissolves in $100\,\text{g}$ of solvent to form a saturated solution at a specified temperature.     $\text{Solubility} = \frac{\text{Mass of Solute}}{\text{Mass of Solvent}} \times 100$

  • Examples at $40^\circ\text{C}$: $100\,\text{g}$ of water dissolves $36.5\,\text{g}$ of $NaCl$ and $63\,\text{g}$ of $KNO_3$.

  • Solubility Classifications:

  • Insoluble: Trace solubility ($< 0.000015\,\text{g}$ per $100\,\text{g}$ water for silver chloride, $AgCl$).

  • Sparingly Soluble: Low solubility ($0.17\,\text{g}$ per $100\,\text{g}$ water for calcium hydroxide, $Ca(OH)_2$).

  • Soluble: High solubility (e.g., sodium chloride, $NaCl$).

Solubility Curves and Experimental Determination

• Factors Influencing Dissolution Rate

  • Particle Size: Smaller solute particles present a greater surface area to solvent molecules, increasing dissolution rate.

  • Stirring: Continuously brings fresh solvent into contact with solute surfaces.

  • Temperature: Increases kinetic energy, usually enhancing dissolution of solids in liquids.

• Solubility Curves

  • A graphical plot displaying variations in solubility (y-axis, in grams of solute per $100\,\text{g}$ solvent) against temperature (x-axis, in $^\circ\text{C}$).

  • Behavioral Patterns:

  • Substantial Increase: $KNO_3$, $NaNO_3$, $KBr$ show dramatic solubility increases with rising temperature.

    • $KNO_3$ solubility: $63\,\text{g}$ at $40^\circ\text{C}$, $84\,\text{g}$ at $50^\circ\text{C}$, $108\,\text{g}$ at $60^\circ\text{C}$.

    • $NaNO_3$ solubility: $105\,\text{g}$ at $40^\circ\text{C}$, $114\,\text{g}$ at $50^\circ\text{C}$, $124\,\text{g}$ at $60^\circ\text{C}$.

  • Slight Increase: $NaCl$ shows only a minor solubility increase over wide temperature ranges.

  • Solubility Decrease: Hydrated calcium sulphate ($CaSO_4 \cdot 2H_2O$), anhydrous calcium sulphate ($CaSO_4$), ceric sulphate ($Ce_2(SO_4)_3$), and slaked lime ($Ca(OH)_2$) exhibit decreasing solubility with rising temperature due to exothermic dissolution processes.

  • Anomalous Solubility (Glauber's Salt, $Na_2SO_4 \cdot 10H_2O$): Solubility increases up to $32.8^\circ\text{C}$ as hydrated salt, then decreases above $32.8^\circ\text{C}$ as it dehydrates into anhydrous $Na_2SO_4$.

  • Thermodynamic Basis:

  • Endothermic dissolution: Solubility increases with rising temperature ($KNO_3$, $NaNO_3$).

  • Exothermic dissolution: Solubility decreases with rising temperature ($CaSO_4$, $Ca(OH)_2$).

• Experimental Determination of Solubility

  1. Measure mass of clean, dry evaporating dish = $M\,\text{g}$.

  2. Add saturated solution and weigh dish = $M_1\,\text{g}$.

  3. Evaporate solution to dryness and weigh residue + dish = $M_2\,\text{g}$.

  4. Mass of saturated solution = $(M_1 - M)\,\text{g}$.

  5. Mass of dry solute residue = $(M_2 - M)\,\text{g}$.

  6. Mass of solvent = $(M_1 - M) - (M_2 - M) = (M_1 - M_2)\,\text{g}$.

  7. Formula:      $\text{Solubility} = \frac{M_2 - M}{M_1 - M_2} \times 100$

Gases in Liquids: Factors Affecting Solubility and Henry's Law

• Pressure Effect (Henry's Law)

  • Henry's Law: At constant temperature, the mass of a gas dissolved by a fixed volume of liquid is directly proportional to the pressure applied to the surface of the liquid.

  • Carbonated Beverages: Soda water is bottled under high pressure to dissolve high volumes of carbon dioxide ($CO_2$). Opening the cap reduces pressure, causing dissolved $CO_2$ gas to escape rapidly as effervescence.

• Temperature Effect

  • Gas solubility in liquids decreases continuously with increasing temperature.

  • Boiling Water: Heating water expels dissolved gases, leaving boiled water flat and tasteless.

  • Chilled Beverages: Cold soft drinks retain higher amounts of dissolved gas, giving off less effervescence until warmed.

• Summary of Solute Dissolution Dependencies

  • Solid Solutes: Solubility generally increases with temperature; pressure has no effect.

  • Gaseous Solutes: Solubility decreases with temperature; pressure significantly increases solubility.

Crystals, Crystallisation, and Water of Crystallisation

• Crystals and Geometrical Shapes

  • Crystal: A homogeneous solid exhibiting a definite geometrical shape bounded by smooth, symmetrically arranged plane surfaces meeting at sharp edges.

  • Geometrical Forms:

  • Cubic: Sodium chloride ($NaCl$)

  • Rhombohedric: Copper sulphate ($CuSO_4$)

  • Octahedral: Iron (II) sulphate ($FeSO_4$)

  • Prismatic: Potassium nitrate ($KNO_3$)

• Crystallisation Methods

  • Process of obtaining crystals from a hot saturated solution upon cooling.

  • Laboratory Methods:

  1. Cooling a hot saturated solution.

  2. Slow evaporation of a saturated solution.

  3. Cooling a molten/fused mass.

  4. Sublimation.

  • Seeding: Inducing rapid crystal growth by suspending a well-formed "seed crystal" into a dust-free, cooled saturated solution using a thread.

  • Trace Impurities in Gemstones: Ruby (red colour due to trace mercury oxide), Sapphire, Diamond.

• Water of Crystallisation (Water of Hydration)

  • Definition: A fixed amount of water molecules held in loose chemical combination with a salt, forming an integral part of its crystalline structure.

  • Thermal Decomposition: Heating hydrated crystals above $100^\circ\text{C}$ expels water of crystallisation as steam, leaving behind an amorphous, non-crystalline anhydrous residue.

• Key Hydrated Salts Table

  • Washing soda crystals: Sodium carbonate decahydrate | Formula: $Na_2CO_3 \cdot 10H_2O$ | Colour: White

  • Epsom salt: Magnesium sulphate heptahydrate | Formula: $MgSO_4 \cdot 7H_2O$ | Colour: White

  • Potash alum: Hydrated potassium aluminium sulphate | Formula: $K_2SO_4 \cdot Al_2(SO_4)_3 \cdot 24H_2O$ | Colour: White

  • Glauber's salt: Sodium sulphate decahydrate | Formula: $Na_2SO_4 \cdot 10H_2O$ | Colour: White

  • Blue vitriol: Copper (II) sulphate pentahydrate | Formula: $CuSO_4 \cdot 5H_2O$ | Colour: Blue

  • Green vitriol: Iron (II) sulphate heptahydrate | Formula: $FeSO_4 \cdot 7H_2O$ | Colour: Bluish green

  • Plaster of Paris: Calcium sulphate semihydrate | Formula: $CaSO_4 \cdot \frac{1}{2}H_2O$ | Colour: White

  • White vitriol: Zinc sulphate heptahydrate | Formula: $ZnSO_4 \cdot 7H_2O$ | Colour: White

  • Gypsum: Hydrated calcium sulphate | Formula: $CaSO_4 \cdot 2H_2O$ | Colour: Colourless

• Reversible Heating Reaction of Blue Vitriol:   $CuSO_4 \cdot 5H_2O(s) \underset{\text{Cool (+ water)}}{\overset{\text{Heat (- water)}}{\rightleftharpoons}} CuSO_4(s) + 5H_2O(g)$   (Blue hydrated crystals revert to white anhydrous powder upon heating, and turn blue again upon adding water drops).

• Calculating Percentage of Water of Crystallisation

  • Mass of hydrated crystals = $a\,\text{g}$

  • Mass of dry residue after heating at $101^\circ\text{C}$ = $b\,\text{g}$

  • Mass of water driven off = $(a - b)\,\text{g}$   $\%\,\text{Water of Crystallisation} = \frac{a - b}{a} \times 100$

Hydrated vs. Anhydrous Substances and Water Elimination

• Definitions

  • Hydrated Substances: Compounds containing chemically combined water molecules of crystallisation ($Na_2CO_3 \cdot 10H_2O$, $CuSO_4 \cdot 5H_2O$).

  • Anhydrous Substances: Compounds containing no water molecules ($NaCl$, $KNO_3$, gaseous $HCl$).

• Methods for Preparing Anhydrous Salts from Hydrated Salts

  • Direct heating.

  • Heating in hot, dry air.

  • Heating under vacuum.

  • Exposure to desiccants/dehydrating agents (e.g., warm concentrated sulphuric acid).

Efflorescence, Deliquescence, and Hygroscopy

• Efflorescence

  • Definition: Phenomenon where a hydrated crystalline salt loses part or all of its water of crystallisation when exposed to air, forming a powdery coating.

  • Physical Condition: Occurs when the vapour pressure of the hydrated salt exceeds the atmospheric vapour pressure.

  • Minimized in humid conditions; accelerated by higher temperatures.

  • Equations:

  • Washing soda: $Na_2CO_3 \cdot 10H_2O(s) \rightarrow Na_2CO_3 \cdot H_2O(s) + 9H_2O(g)$

  • Glauber's salt: $Na_2SO_4 \cdot 10H_2O(s) \rightarrow Na_2SO_4(s) + 10H_2O(g)$

  • Epsom salt: $MgSO_4 \cdot 7H_2O(s) \rightarrow MgSO_4 \cdot H_2O(s) + 6H_2O(g)$

• Deliquescence

  • Definition: Phenomenon where a substance absorbs moisture from atmosphere upon exposure, dissolves in the absorbed moisture, and forms a saturated liquid solution.

  • Physical Condition: Occurs when the vapour pressure of the salt's saturated solution is lower than atmospheric vapour pressure.

  • Examples: Caustic soda ($NaOH$), Caustic potash ($KOH$), Calcium chloride ($CaCl_2$), Magnesium chloride ($MgCl_2$), Ferric chloride ($FeCl_3$).

  • Impurities in Common Salt: Pure sodium chloride ($NaCl$) is non-deliquescent. Common table salt turns wet/sticky during rainy seasons because it contains trace deliquescent impurities ($MgCl_2$, $CaCl_2$).

  • Purification Method: Dry $HCl$ gas is passed through a saturated table salt solution. Pure $NaCl$ precipitates out, which is filtered and washed with water and alcohol.

• Hygroscopy

  • Definition: Substances that absorb atmospheric moisture upon exposure without dissolving or forming solutions.

  • Examples: Concentrated sulphuric acid ($H_2SO_4$), Phosphorus pentoxide ($P_2O_5$), Quicklime ($CaO$), Silica gel.

Drying Agents vs. Dehydrating Agents

• Definitions & Differences

  • Drying (Desiccating) Agents:

  • Remove physical moisture/water vapour from other substances without chemical reaction.

  • Involve physical changes.

  • Used to dry acidic/neutral gases ($Cl_2$, $SO_2$, $HCl$) and maintain dry conditions inside desiccators.

  • Examples: Quicklime ($CaO$), fused $CaCl_2$, anhydrous $ZnCl_2$, $P_2O_5$, $MgSO_4$, dry $Na_2SO_4$, conc. $H_2SO_4$.

  • Gas Specifics: Acidic gases ($HCl$) are dried using concentrated $H_2SO_4$. Basic gases ($NH_3$) are dried using Quicklime ($CaO$) because $CaO$ does not react chemically with ammonia.

  • Dehydrating Agents:

  • Remove chemically combined hydrogen and oxygen elements in a $2:1$ ratio (elements of water) from compounds.

  • Involve chemical changes, generating entirely new substances.

  • Used in organic synthesis (e.g., preparing carbon monoxide, sugar charcoal).

  • Examples: Concentrated sulphuric acid ($H_2SO_4$), Alumina ($Al_2O_3$), Phosphorus pentoxide ($P_2O_5$).

  • Dehydration Chemical Equations:

  • Blue Vitriol: $CuSO_4 \cdot 5H_2O \xrightarrow{\text{Conc. } H_2SO_4} CuSO_4 + 5H_2O$

  • Cane Sugar: $C_{12}H_{22}O_{11} \xrightarrow{\text{Conc. } H_2SO_4} 12C + 11H_2O$

  • Formic Acid: $HCOOH \xrightarrow{\text{Conc. } H_2SO_4} CO + H_2O$

Hard Water and Soft Water: Causes, Types, and Implications

• Soft Water vs. Hard Water

  • Soft Water: Readily forms rich lather with soap (e.g., rainwater, distilled water).

  • Hard Water: Does not readily form lather with soap; forms insoluble scum/curd instead (e.g., river water, well water, spring water).

• Causes of Hardness

  • Presence of dissolved bicarbonates, chlorides, or sulphates of calcium ($Ca^{2+}$) and magnesium ($Mg^{2+}$).

  • Formation Mechanism: Rainwater containing dissolved atmospheric $CO_2$ flows over limestone or dolomite beds ($CaCO_3$, $MgCO_3$), converting them into soluble bicarbonates:     $CaCO_3(s) + CO_2(g) + H_2O(l) \rightarrow Ca(HCO_3)_2(aq)$     $MgCO_3(s) + CO_2(g) + H_2O(l) \rightarrow Mg(HCO_3)_2(aq)$

  • Gypsum dissolving: Water flowing over gypsum beds ($CaSO_4 \cdot 2H_2O$) dissolves small amounts of $CaSO_4$.

• Stalactites and Stalagmites

  • Dripping water containing dissolved $Ca(HCO_3)_2$ inside limestone caves experiences warmth and pressure drop, decomposing into solid $CaCO_3$:     $Ca(HCO_3)_2(aq) \rightarrow CaCO_3(s) + H_2O(l) + CO_2(g)$

  • Stalactite: Conical pillar growing downward from the cave ceiling.

  • Stalagmite: Conical pillar growing upward from the cave floor.

  • Growth rates vary from less than $1\,\text{cm}$ to over $100\,\text{cm}$ per year.

• Types of Hardness

  • Temporary Hardness: Caused by dissolved hydrogen carbonates of calcium and magnesium [$Ca(HCO_3)_2$, $Mg(HCO_3)_2$]. Can be removed completely by boiling.

  • Permanent Hardness: Caused by dissolved chlorides and sulphates of calcium and magnesium ($CaCl_2$, $MgCl_2$, $CaSO_4$, $MgSO_4$). Cannot be removed by boiling.

• Advantages of Hard Water

  • Tastes pleasant; used in preparing beverages and wines.

  • Supplies calcium and magnesium essential for bone and tooth growth.

  • Checks lead poisoning: Forms an insoluble protective layer of lead sulphate/calcium sulphate along the interior walls of lead pipes.

• Disadvantages of Hard Water

  • Boiler Scale: Deposits thick scale ($CaCO_3$, $CaSO_4$) inside copper steam tubes. Scale is a poor thermal conductor, reducing heating efficiency, wasting fuel, narrowing tube diameter, and triggering boiler explosions.

  • Furring of Kettles: Forms white carbonate crusts inside tea kettles.

  • Plumbing & Appliance Damage: Corrodes pipes, taps, water heaters, and washing machines.

  • Soap Wastage: Soap (sodium stearate, $NaSt$ or $C_{17}H_{35}COONa$) reacts with hard water ions to form insoluble scum ($CaSt_2$, $MgSt_2$), wasting soap and staining fabrics:

  • Temporary Hardness Reactions:       $2NaSt(aq) + Ca(HCO_3)_2(aq) \rightarrow CaSt_2(s)\downarrow + 2NaHCO_3(aq)$       $2NaSt(aq) + Mg(HCO_3)_2(aq) \rightarrow MgSt_2(s)\downarrow + 2NaHCO_3(aq)$

  • Permanent Hardness Reactions:       $2NaSt(aq) + CaCl_2(aq) \rightarrow CaSt_2(s)\downarrow + 2NaCl(aq)$       $2NaSt(aq) + MgCl_2(aq) \rightarrow MgSt_2(s)\downarrow + 2NaCl(aq)$       $2NaSt(aq) + CaSO_4(aq) \rightarrow CaSt_2(s)\downarrow + 2Na_2SO_4(aq)$       $2NaSt(aq) + MgSO_4(aq) \rightarrow MgSt_2(s)\downarrow + 2Na_2SO_4(aq)$

Cleansing Action of Soap and Detergents

• Synthetic Detergents

  • Manufactured by sulphonating higher alkenes with sulphuric acid and converting them into sodium salts using caustic soda ($NaOH$).

  • Unaffected by water hardness because their calcium and magnesium salts are water-soluble, preventing scum formation ("soapless soaps").

• Mechanism of Cleansing (Micelle Formation)

  • Detergent molecules form spherical clusters termed micelles.

  • Hydrocarbon Tails: Hydrophobic (water-repelling) tails stick inward, attaching directly to oily dirt particles.

  • Ionic Heads: Hydrophilic (water-attracting) negatively charged heads face outward in water.

  • Agitation breaks oil and dirt into small suspended globules surrounded by micelles. Negative charge repulsions between outer heads prevent coalescence, allowing dirt to rinse away.

Methods of Softening Hard Water

• Removal of Temporary Hardness

  1. Boiling: Converts soluble bicarbonates into insoluble carbonates:      $Ca(HCO_3)_2(aq) \xrightarrow{\text{Boil}} CaCO_3(s)\downarrow + H_2O(l) + CO_2(g)\uparrow$      $Mg(HCO_3)_2(aq) \xrightarrow{\text{Boil}} MgCO_3(s)\downarrow + H_2O(l) + CO_2(g)\uparrow$

  2. Clark's Process (Addition of Slaked Lime): Calculated quantity of slaked lime [$Ca(OH)_2$] is added in mixing tanks fitted with revolving paddles:      $Ca(HCO_3)_2(aq) + Ca(OH)_2(aq) \rightarrow 2CaCO_3(s)\downarrow + 2H_2O(l)$      $Mg(HCO_3)_2(aq) + Ca(OH)_2(aq) \rightarrow MgCO_3(s)\downarrow + CaCO_3(s)\downarrow + 2H_2O(l)$

  3. Addition of Washing Soda ($Na_2CO_3$):      $Ca(HCO_3)_2(aq) + Na_2CO_3(aq) \rightarrow CaCO_3(s)\downarrow + 2NaHCO_3(aq)$      $Mg(HCO_3)_2(aq) + Na_2CO_3(aq) \rightarrow MgCO_3(s)\downarrow + 2NaHCO_3(aq)$

• Removal of Permanent Hardness

  1. Addition of Washing Soda / Soda Ash ($Na_2CO_3$):      $CaSO_4(aq) + Na_2CO_3(aq) \rightarrow CaCO_3(s)\downarrow + Na_2SO_4(aq)$      $CaCl_2(aq) + Na_2CO_3(aq) \rightarrow CaCO_3(s)\downarrow + 2NaCl(aq)$      If magnesium sulphate is present, slaked lime is added along with soda ash:      $MgSO_4(aq) + Na_2CO_3(aq) + Ca(OH)_2(aq) \rightarrow Mg(OH)_2(s)\downarrow + CaCO_3(s)\downarrow + Na_2SO_4(aq)$

• Removal of Combined Hardness (Large Scale)

  1. Permutit Process (Zeolite Process):

  • Permutit is an artificial zeolite chemically defined as hydrated sodium aluminium orthosilicate ($Na_2Al_2Si_2O_8 \cdot xH_2O$, abbreviated $Na_2P$).

  • Hard water percolates through a cylinder filled with permutit lumps. $Na^+$ ions exchange with $Ca^{2+}$ and $Mg^{2+}$ ions, yielding soft water:        $Na_2P + Ca^{2+} \rightarrow CaP + 2Na^+$

  • Regeneration: Exhausted permutit ($CaP$) is regenerated by running concentrated brine ($NaCl$) solution over it:        $CaP + 2NaCl \rightarrow Na_2P + CaCl_2$

  1. Synthetic Ion Exchange Resins (Deionisation):

  • Water passes sequentially through Cation Exchange Resins (replaces all cations $Na^+$, $Ca^{2+}$, $Mg^{2+}$ with $H^+$) and Anion Exchange Resins (replaces all anions $Cl^-$, $SO_4^{2-}$, $HCO_3^-$ with $OH^-$).

  • The resulting $H^+$ and $OH^-$ combine to form completely deionised/distilled water.

Potable Water: Standards, Purification Processes, and Water-borne Diseases

• Definition & Quality Standards of Potable Water

  • Water that is completely safe for human consumption without immediate or long-term health risks.

  • Physical Appearance: Clear, transparent, colourless, odourless, and pleasant-tasting.

  • Bureau of Indian Standards (BIS) & WHO Guidelines:

  • pH: $6.5\text{--}8.5$

  • Total Dissolved Solids (TDS): Should not exceed $2000\,\text{mg\,L}^{-1}$

  • Total Hardness: $300\text{--}600\,\text{mg\,L}^{-1}$

  • Chlorides: $250\text{--}1000\,\text{mg\,L}^{-1}$

  • Fluoride: $1.0\text{--}1.5\,\text{mg\,L}^{-1}$

  • Iron: $0.3\text{--}1.0\,\text{mg\,L}^{-1}$

• Water Treatment Operations

  1. Coagulation and Settling: Coagulants such as alum (hydrated double sulphate salt) are added to raw water. Alum forms sticky gel-like flocs that bind bacteria, mud, and suspended impurities, causing them to settle at the bottom of sedimentation basins.

  2. Filtration: Settled water trickles through multi-layer filter beds of sand, gravel, and charcoal/coal (activated carbon) to remove remaining particulate matter.

  3. Chlorination: Addition of calculated amounts of chlorine to kill pathogenic bacteria and microorganisms.

  4. Defluoridation: Process of removing excess fluoride ($> 1.5\,\text{mg\,L}^{-1}$) to prevent skeletal and dental fluorosis. Techniques include adsorption (activated alumina/charcoal), chemical precipitation (Nalgonda alum technique), and membrane separation (Reverse Osmosis RO, Electrodialysis).

  5. Aeration: Improves taste and removes unpleasant dissolved gases/odours.

• Contamination Sources & Water-borne Diseases

  • Contamination Vectors: Improper sewage disposal, open defecation near surface water, livestock/dairy waste run-off, industrial effluents, and agricultural chemical runoff (pesticides, fertilisers).

  • Bacterial Diseases: Diarrhoea, Typhoid, Cholera, Dysentery (cause gastrointestinal pain, high fever, severe dehydration, vomiting, weakness).

  • Viral Diseases: Hepatitis A (affects liver function), Polio (attacks nervous system, causing paralysis).

  • Chemical Toxicity Hazards:

  • Chronic consumption of heavy metals (e.g., mercury, lead, arsenic) or pesticides increases cancer risk and causes irreversible kidney and liver organ damage.

  • Excess fluoride causes dental/skeletal fluorosis.

Questions & Practical Applications

• Practical Case Study: Bleeding Control with Alum

  • Scenario: An individual cuts a finger in the kitchen and applies liquid antiseptic, which fails to stop the bleeding. Applying an alum block stops the bleeding rapidly.

  • Explanation: Liquid antiseptic lacks coagulating properties. Blood is a colloidal solution of negatively charged haemoglobin particles. Alum dissolved in water yields trivalent ions ($Al^{3+}$) that neutralize the electrical charge on colloidal blood particles, inducing rapid coagulation/clotting to seal the wound.

• Numerical Practice Problems

  • Problem 1: Saturated solution made with $136\,\text{g}$ salt in $500\,\text{g}$ water at $293\,\text{K}$.

  • Solubility Calculation: $\frac{136}{500} \times 100 = 27.2\,\text{g}$ per $100\,\text{g}$ water.

  • Problem 2: $15\,\text{g}$ $NaCl$ dissolved in $285\,\text{g}$ water.

  • Mass Percent Calculation: $\frac{15}{15 + 285} \times 100 = \frac{15}{300} \times 100 = 5\%$.

  • Problem 3: $4\,\text{L}$ acetone dissolved in $90\,\text{L}$ aqueous solution.

  • Volume Percent Calculation: $\frac{4}{90} \times 100 = 4.44\%$.

  • Problem 4: Evaporating $12\,\text{g}$ saturated $KCl$ solution at $20^\circ\text{C}$ yields $3\,\text{g}$ solid residue.

  • Solubility Calculation: Mass of water = $12 - 3 = 9\,\text{g}$. Solubility = $\frac{3}{9} \times 100 = 33.3\,\text{g}$ per $100\,\text{g}$ water.

  • Problem 5: Empty dish = $50\,\text{g}$, dish + $KNO_3$ solution = $65\,\text{g}$, dish + dry residue = $54.3\,\text{g}$ at $20^\circ\text{C}$.

  • Solubility Calculation: Mass of solute = $54.3 - 50 = 4.3\,\text{g}$. Mass of water = $65 - 54.3 = 10.7\,\text{g}$. Solubility = $\frac{4.3}{10.7} \times 100 = 40.18\,\text{g}$ per $100\,\text{g}$ water.

  • Problem 6: Mass of $NaNO_3$ separating when $50\,\text{g}$ water saturated solution cools from $50^\circ\text{C}$ (solubility = $114\,\text{g}$) to $30^\circ\text{C}$ (solubility = $86\,\text{g}$).

  • Crystallisation Mass Calculation: Solute precipitated per $100\,\text{g}$ water = $114 - 86 = 28\,\text{g}$. For $50\,\text{g}$ water = $\frac{28}{2} = 14\,\text{g}$.