Water and Hardness Chemistry Notes

General Properties and Composition of Water

• Water is a colourless and odourless liquid.
• Chemical Composition: It consists of hydrogen and oxygen in a ratio of $2:1$ ($H_2O$).
• Solvent Properties: It is known as the "universal solvent" because it is capable of dissolving a wide variety of substances.
• Purity in Nature: Water rarely occurs in its pure form in nature due to its ability to dissolve minerals and other substances it comes into contact with.

Sources of Water

• Natural Sources:
  • Rain Water: Known as the purest natural form of water, essentially acting as a natural form of distilled water.
  • Spring Water.
  • Well Water.
  • River Water.
  • Lake Water.
  • Sea Water.
• Impurities in Natural Sources: Waters from sources other than rain contain significant amounts of dissolved air, soil particles, mineral salts, bacteria, and organic remains.
• Treated Water (Prepared for Special Purposes):
  • Distilled Water: Chemically pure water prepared by condensing steam or passing water through an ion-exchange resin column.
  • Pipe-borne Water: Prepared for use in townships.
  • Chlorinated Water: Specifically prepared for swimming pools to ensure hygiene.

Applications of Distilled Water

• Distilled water is utilized in specific settings where chemical purity is required:
  • Laboratory use: Essential for preparing reagents and for sensitive analytical work.
  • Industrial processes: Used in various manufacturing and chemical operations.
  • Pharmaceuticals: Used in the preparation of drugs.
  • Automotive: Used in car batteries to prevent mineral buildup.

Unusual Physical Properties of Water

• High Melting and Boiling Points:
  • Water possesses melting and boiling points that are unexpectedly high for a molecule of its size and weight.
  • Cause: This is attributed to the presence of hydrogen bonds which link water molecules together strongly.
• Surface Tension: The same hydrogen bonds responsible for high boiling points also result in high surface tension.
• Density of Solid Phase (Ice):
  • Ice is less dense than liquid water, allowing it to float. This is because ice has an open hydrogen-bonded hexagonal lattice structure.
  • Molecular Arrangement: Each oxygen atom is surrounded tetrahedrally by four hydrogen atoms. Two atoms are covalently bonded to the oxygen within the $H_2O$ molecule (closer), and two others are held by hydrogen bonds (farther away).
  • Three-Dimensional Structure: The tetrahedral angles create an open three-dimensional space within the structure, reducing density.
• Large Heat Capacity:
  • Definition: Heat capacity is the amount of heat required to raise the temperature of a substance by one degree.
  • Coolant Properties: Water absorbs a significant amount of heat before its temperature rises. Consequently, it is used as a coolant in industrial settings and car radiators.
  • Climate Regulation: The high specific heat index of water regulates air temperature changes. This explains why seasonal temperature shifts are gradual, particularly near oceans.
  • Oceanic Influence: Oceans and lakes regulate temperatures for nearby towns and cities. Because water takes longer to heat up and cool down than land masses, coastal cities experience less extreme temperature fluctuations than inland cities.
• Solubility and Polarity:
  • Water is both a polar substance and a hydrogen-bonding molecule, making it the premier solvent biologically and industrially.

General Uses of Water

• Domestic Purposes: Essential for drinking, washing, and bathing.
• Irrigation: Used in agriculture to water crops.
• Hydro-electric Power: Used for generating electricity.
• Transportation: Water bodies provide routes for trade and travel.
• Biological Processes: Used by green plants for photosynthesis.
• Industrial Purposes: Used as a raw material, solvent, and coolant.

Chemical Tests for Water

• Cobalt (II) Chloride Test:
  • Adding water to blue anhydrous cobalt (II) chloride turns it pink.
  • Equation: $CoCl_2 + 6H_2O \rightarrow CoCl_2.6H_2O$
  • Transition: Blue (Anhydrous) to Pink (Hydrated).
• Copper (II) Tetraoxosulphate (VI) Test:
  • Adding water to white anhydrous copper (II) tetraoxosulphate (VI) turns it blue.
  • Equation: $CuSO_4(s) + 5H_2O \rightarrow CuSO_4.5H_2O$
  • Transition: White (Anhydrous) to Blue (Hydrated).

Treatment of Municipal and Township Water

• Goal: To make water germ-free and fit for human consumption.
• Process Steps:
  1. Coagulation/Flocculation: Untreated water enters large settling tanks. Lime and chemicals like potash alum ($KAl(SO_4)_2$) or sodium aluminate (III) ($NaAlO_2$) are added to clump fine particles together.
  2. Sedimentation: Particles settle at the bottom of the tanks.
  3. Filtration: The water passes through a filter bed to remove remaining fine particles of dirt.
  4. Disinfection/Chemical Treatment:
     • Chlorine: Added to kill harmful germs.
     • Iodine: Added to prevent goitre.
     • Fluorine: Added to prevent tooth decay.
  5. Storage and Distribution: The treated water is stored in a reservoir and distributed via underground pipes.

Hardness of Water

• Definition: Hard water is water that does not form lather easily with soap.
• Chemical Cause: Caused by dissolved salts such as $CaSO_4$, $MgSO_4$, $Ca(HCO_3)_2$, $CaCl_2$, and $MgCl_2$.
• Reaction with Soap:
  • Soap (sodium octadecanoate) reacts with the dissolved calcium or magnesium salts.
  • This forms an insoluble scum (calcium or magnesium octadecanoate).
  • The scum sticks to clothes and is difficult to rinse.
• Interaction with Detergents: Detergents are not affected by hardness because they form soluble salts with calcium and magnesium.
• Origin: Water becomes hard by dissolving minerals like Gypsum ($CaSO_4.2H_2O$) or Limestone ($CaCO_3$) as it flows over soil.
  • Dissolution of Limestone: While $CaCO_3$ is insoluble in pure water, water containing dissolved carbon dioxide ($CO_2$) can dissolve it:
  • $CaCO_3(s) + CO_2(g) + H_2O(l) \rightleftharpoons Ca(HCO_3)_2(aq)$

Types of Hardness and Methods of Removal

• Temporary Hardness:

  • Cause: Dissolved calcium hydrogen trioxocarbonate (IV), $Ca(HCO_3)_2$.
  • Removal by Boiling: Heating causes decomposition into insoluble $CaCO_3$.
  • Equation: $Ca(HCO_3)_2(aq) \rightarrow CaCO_3(s) + CO_2(g) + H_2O(l)$
  • Removal by Slaked Lime ($Ca(OH)_2$): Adding slaked lime precipitates the calcium ions.
  • Equation: $Ca(HCO_3)_2(aq) + Ca(OH)_2(s) \rightarrow 2CaCO_3(s) + 2H_2O(l)$

• Permanent Hardness:

  • Cause: Dissolved calcium or magnesium tetroxosulphates (VI) and chlorides ($CaSO_4$, $MgSO_4$, $CaCl_2$, $MgCl_2$).
  • Removal by Washing Soda ($Na_2CO_3$): Adds sodium carbonate to precipitate calcium as a solid.
  • Equation: $Na_2CO_3(aq) + CaSO_4(aq) \rightarrow CaCO_3(s) + Na_2SO_4(aq)$
  • Removal by Permutit or Zeolite (Ion-Exchange):
    • The process uses an ion-exchange resin containing sodium aluminium trioxosilicate (IV).
    • Hard water passes through the resin, and sodium ions ($Na^+$) in the complex salt are replaced by calcium ($Ca^{2+}$) and magnesium ($Mg^{2+}$) ions.
    • Equation: $CaSO_4(aq) + 2NaAl(SiO_3)_2(s) \rightarrow CaAl_2(SiO_3)_4(s) + Na_2SO_4(aq)$
    • Regeneration: When used up, it is regenerated by passing a strong solution of common salt ($NaCl$) through the resin ($2NaCl + CaAl_2(SiO_3)_4(s) \rightarrow 2NaAl(SiO_3)_2(s) + CaCl_2$).

Consequences and Effects of Water Hardness

• Furring: The decomposition of $Ca(HCO_3)_2$ during boiling causes a gradual deposition of a white "fur" of $CaCO_3$ on the inner surfaces of kettles and boilers.
• Stalagmites and Stalactites: As temporarily hard water drips through cave roofs, heat decomposes the $Ca(HCO_3)_2$. $CaCO_3$ deposits on the roof (stalactites) and floor (stalagmites).
• Advantages of Hard Water:
  1. Better taste due to dissolved minerals.
  2. Calcium salts assist in building strong teeth and bones in animals.
  3. Provides materials ($CaCO_3$) for shells of snails and crabs.
  4. Safety in Lead Pipes: Hard water does not dissolve lead. Soft water can dissolve lead, leading to lead poisoning if lead pipes are used.
• Disadvantages of Hard Water:
  1. Soap Waste: Requires more soap to achieve lather.
  2. Boiler Damage: Causes furring in industrial and domestic boilers.
  3. Industrial Interference: Cannot be used in dyeing and tanning because the salts interfere with the chemical processes.

Questions & Discussion

• Q: State five sources of water.
  • A: Rain water, spring water, well water, river water, and sea water.
• Q: Highlight three uses of treated water.
  • A: Analytical lab work (distilled), chlorinated water for swimming pools, and pipe-borne water for towns.
• Q: Define hard water.
  • A: Hard water is water that does not form lather readily/easily with soap due to dissolved calcium and magnesium salts.
• Q: Explain three methods of removing permanent hardness of water.
  • A: Addition of washing soda ($Na_2CO_3$) to precipitate calcium/magnesium, addition of permutit/zeolite for ion-exchange, and the use of specialized ion-exchange resins.
• Q: State the role of Sand Bed, Alum, and Chlorine in water treatment.
  • A: Sand Bed: Filtration of fine dirt. Alum: Coagulation/flocculation of particles. Chlorine: Disinfection (killing germs).
• Q: Explain why water is a universal solvent.
  • A: Because it is a polar substance with hydrogen bonding capabilities, allowing it to dissolve a wide range of ionic and polar substances.
• Q: Why is it bad to build with limestone in an environment polluted by $SO_4$ or $SO_2$?
  • A: This relates to the dissolution of limestone. Acidic pollutants like $SO_2$ react with moisture to form acid rain, which dissolves the limestone ($CaCO_3$) into soluble forms, causing structural degradation.