Water and Solution
Contents
- Introduction to Water
- Physical Properties of Water
- Chemical Properties of Water
- Uses of Water
- Solutions – Types & Classification
- Solubility & Factors Affecting It
- Molarity & Molality
- True Solution, Suspension & Colloid
- Crystallisation & Water of Crystallisation
- Hygroscopic, Efflorescent & Deliquescent Substances
Introduction to Water
Apart from air, water is the most abundant substance on Earth. Henry Cavendish in 1784 burnt a mixture of hydrogen and oxygen and proved that water is not an element but a compound.
Water is the only substance on Earth that naturally exists in all three states — solid (ice), liquid (water) and gas (steam). Water can be changed from one state to another by addition or removal of heat.
Check Out- CBSE Class 9 Science Notes
Physical Properties of Water
Odourless, tasteless, transparent liquid. Appears bluish in thick layers. Taste is due to dissolved salts.
0°C at normal atmospheric pressure. Decreases with increased pressure or dissolved impurities.
100°C at normal atmospheric pressure. Elevated by dissolved impurities or increased pressure.
Maximum density 1 g/cc at 4°C. Shows anomalous expansion: expands on cooling between 0°C–4°C instead of contracting.
Pure water is a poor conductor of heat and electricity. Electrical conductivity is due to dissolved salts.
1 cal/g = 4.2 J/g. Highest specific heat among all substances.
Dissolves many substances due to its high dielectric constant (reduces attractive forces between ions of inorganic compounds).
Latent heat of fusion = 80 cal/g. Latent heat of vaporisation = 540 cal/g.
Chemical Properties of Water
Water reacts with metals, non-metals, metallic and non-metallic oxides, and even acts as a catalyst in some reactions.
Reactions with Metals
Potassium (cold water / moisture)
- Vigorous and exothermic; burns with lilac flame; water becomes alkaline.
Sodium (cold water)
- Floats as silvery globules; burns with golden yellow flame; less vigorous than K.
Calcium (cold water)
- Sinks; water turns milky and alkaline; no fire.
Magnesium (boiling water / steam)
- Burns brilliantly with white light; produces white ash of MgO; hydrogen liberated.
Zinc (steam over red-hot zinc)
- Yellow ZnO formed (becomes white on cooling).
Aluminium (steam)
- Al₂O₃ coating protects the metal from further reaction.
Iron (steam over red-hot iron)
- Brown ferroso-ferric oxide formed; reversible in a closed container.
Reactions with Non-Metals
Carbon (superheated steam over red-hot coke)
Chlorine
Reactions with Oxides
Metallic oxides + water → bases:
Non-metallic oxides + water → acids:
Catalytic Property
Water catalyses some reactions. For example, perfectly dry H₂ and Cl₂ do not react even in sunlight — a few drops of water catalyse the reaction. Similarly, combustion of phosphorus requires trace amounts of moisture.
Uses of Water
- Domestic (≈8%): Cooking, washing, cleaning, flushing.
- Industrial (≈22%): Chemical labs, pharmaceutical industries, steam for electricity generation, cooling systems (high specific heat).
- Agriculture (≈70%): Irrigation; transport of minerals and nutrients in plants; dissolving fertilisers/insecticides for spraying.
- Other: Production of common salt (saline water); production of MgO (magnesia) for refractories.
Solutions
A solution is a homogeneous mixture of two or more substances. The component in the same physical state as the solution (or present in greater quantity) is the solvent; the other is the solute. Two-component solutions are called binary solutions.
Classification by Solvent/Solute States
| Solvent | Solute | Example |
|---|---|---|
| Solid | Solid | Metal alloys |
| Liquid | Hydrated crystalline salts | |
| Gas | H₂ adsorbed on platinum | |
| Liquid | Solid | Common salt in water |
| Liquid | Petrol in kerosene | |
| Gas | Aerated water / soft drinks | |
| Gas | Gas | Air |
Aqueous solutions use water as solvent. Non-aqueous solutions use other solvents (alcohol, petrol, ether, benzene).
Based on Solute Proportion
- Dilute solution: Small amount of solute in a given mass of solvent.
- Concentrated solution: Relatively large amount of solute.
- Saturated solution: Maximum solute dissolved at a given temperature; no more can dissolve.
- Unsaturated solution: Less solute than maximum possible; can dissolve more.
- Supersaturated solution: Holds more solute than the saturated solution at the same temperature (metastable; slight disturbance causes excess solute to precipitate).
Solubility
The maximum mass (in grams) of a solute that can dissolve in 100 g of solvent at a given temperature.
Factors Affecting Solubility
- Nature of solute: Ionic/polar compounds dissolve more in polar solvents (water); non-polar in non-polar solvents (benzene).
- Nature of solvent: Solvents with high dielectric constants dissolve polar/ionic compounds better.
- Temperature: Solubility of most solids increases with temperature. Solubility of gases in liquids decreases with temperature.
Rate of Dissolution (Factors)
- Particle size: Smaller particles → greater surface area → faster dissolution.
- Agitation: Stirring improves solute–solvent contact → faster dissolution.
- Temperature: Higher temperature generally increases rate of dissolution for solids.
Effect of Pressure on Gases
Henry's Law: At constant temperature, increased pressure on the surface of a liquid increases the solubility of a gas in that liquid.
Molarity & Molality
Molarity (M)
The number of moles of solute dissolved in one litre of solution. Units: mol/L (abbreviated M). Example: 0.1 M means 0.1 mol of solute per litre of solution.
Molality (m)
The number of moles of solute divided by the kilograms of solvent. Unlike molarity, molality does not change with temperature (since mass is used, not volume).
True Solution, Suspension & Colloidal Solution
| Property | True Solution | Colloidal Solution | Suspension |
|---|---|---|---|
| Particle size | < 1 nm | 1–1000 nm | > 1000 nm |
| Nature | Homogeneous | Heterogeneous | Heterogeneous |
| Filter paper | Passes through | Passes through | Does not pass |
| Parchment paper | Passes through | Does not pass | Does not pass |
| Visibility | Not visible (naked eye) | Not visible (needs ultra-microscope) | Visible to naked eye |
| Tyndall effect | No | Yes | May or may not |
| Appearance | Transparent | Translucent | Opaque |
Crystallisation & Water of Crystallisation
When a saturated solution is allowed to cool, the solute settles with a regular arrangement of particles called a crystalline solid. The smallest repeating unit is the unit cell.
Crystal Systems
| Crystal System | Examples |
|---|---|
| Cubic | NaCl, CsCl |
| Hexagonal | Graphite, ZnO |
| Tetragonal | SnO₂, TiO₂ |
| Rhombohedral | CaCO₃, HgS |
| Orthorhombic | Rhombic sulphur |
| Monoclinic | Monoclinic sulphur, PbCrO₂ |
| Triclinic | K₂Cr₂O₇, CuSO₄·5H₂O |
Water of Crystallisation
A fixed number of water molecules that combine with a crystal and are necessary for maintaining its crystalline properties. Examples: green vitriol (FeSO₄·7H₂O), blue vitriol (CuSO₄·5H₂O), washing soda (Na₂CO₃·10H₂O).
Hydrated salts contain water of crystallisation; anhydrous salts have lost it. Loss of water often changes colour — e.g., blue CuSO₄·5H₂O turns white when heated.
Common Hydrous Substances
| Common Name | Salt |
|---|---|
| Blue vitriol | CuSO₄·5H₂O |
| Green vitriol | FeSO₄·7H₂O |
| Epsom salt | MgSO₄·7H₂O |
| Washing soda (crystal) | Na₂CO₃·10H₂O |
| Washing soda (powder) | Na₂CO₃·H₂O |
| Glauber's salt | Na₂SO₄·10H₂O |
| White vitriol | ZnSO₄·7H₂O |
| Potash alum | KAl(SO₄)₂·12H₂O |
| Gypsum | CaSO₄·2H₂O |
| Plaster of Paris | CaSO₄·½H₂O |
Hygroscopic, Efflorescent & Deliquescent Substances
Efflorescence
The property of certain hydrous crystals to lose their water of crystallisation when exposed to air, crumbling into a powder. Higher air temperature → more efflorescence.
Examples: Washing soda (Na₂CO₃·10H₂O → Na₂CO₃·H₂O + 9H₂O), Glauber's salt, Epsom salt.
Hygroscopic Substances
Substances that absorb water vapour from the air without dissolving in it. Used as drying agents in the laboratory.
| Hygroscopic Substances | |
|---|---|
| Concentrated H₂SO₄ | Anhydrous CaCl₂ |
| CaO (quicklime) | Solid NaOH / KOH |
| P₂O₅ | Silica gel |
Deliquescence
Certain hygroscopic substances absorb so much moisture that they dissolve in it, forming a saturated solution. Occurs when the vapour pressure of the salt is much lower than atmospheric vapour pressure. Minimised in dry conditions (where efflorescence is maximised).
Examples: NaOH, KOH, MgCl₂, ZnCl₂, CaCl₂, FeCl₃, Zn(NO₃)₂, Cu(NO₃)₂.
Drying Agents
Since deliquescent and hygroscopic substances have affinity for water, they can be used as drying agents. Commonly used: anhydrous CaCl₂, quicklime (CaO), and concentrated H₂SO₄.