Hello, and welcome to your chemistry lesson. Today, we are going to explore one of the most remarkable substances on our planet — water. We will discover why it is called the universal solvent, how it dissolves substances, what makes water hard or soft, and the fascinating properties of crystals that contain water. Let us begin.
Water, with the chemical formula H₂O, also called dihydrogen oxide, is truly extraordinary. It exists in all three states of matter — solid as ice, liquid as water, and gas as water vapour. More than seventy percent of Earth's surface is covered by water, yet only about two and a half percent is fresh water. Water is essential for all life forms. Nearly seventy percent of your body weight is water.
Pure water is colourless, odourless, and tasteless. Any taste you detect comes from dissolved gases and minerals. Under normal atmospheric pressure of seven hundred sixty millimetres of mercury, water boils at one hundred degrees Celsius and freezes at zero degrees Celsius. However, these points change with pressure and dissolved impurities. Greater pressure raises the boiling point, which is why pressure cookers work so effectively, and why water boils below one hundred degrees Celsius in the hills where atmospheric pressure is lower.
Water displays a peculiar behaviour called anomalous expansion. Most liquids contract when cooled, but water contracts only until 4 degrees Celsius. Below this temperature, it actually expands until it freezes. This is why ice floats on water. This property is vital for aquatic life — when ponds freeze in winter, ice forms at the surface while water remains liquid below, allowing fish and other organisms to survive.
Water also possesses remarkably high latent heat. The latent heat of fusion of ice is three hundred thirty-six joules per gram or eighty calories per gram, meaning substantial heat is needed to melt ice. This prevents rivers and lakes from freezing suddenly. The latent heat of vaporization is even higher at two thousand two hundred sixty-eight joules per gram or five hundred forty calories per gram. This explains why steam causes more severe burns than boiling water — it releases this enormous heat upon condensing on your skin.
Additionally, water has a high specific heat capacity of four point two joules per gram per degree Celsius, or one calorie per gram per degree Celsius. This means water absorbs large amounts of heat with only small temperature changes. This moderates climate near water bodies, creating milder winters and cooler summers, and explains why coastal cities experience less extreme temperatures than inland cities.
Now, let us understand why water is called the universal solvent. A solvent is a medium in which other substances dissolve. Water dissolves an incredibly wide range of substances — solids, liquids, and gases. This happens because water molecules are polar — the oxygen atom carries a slight negative charge while the hydrogen atoms carry slight positive charges. This polarity allows water to attract and surround different types of molecules and ions.
When a substance dissolves in water, we call the mixture a solution. The substance that dissolves is the solute, and water is the solvent. Such solutions are specifically called aqueous solutions. True solutions are homogeneous mixtures where solute particles are extremely small, about ten to the minus ten metres, making them indistinguishable from solvent molecules even under microscopes.
Natural water always contains dissolved substances. Rainwater and distilled water are nearly pure, but river water, well water, and tap water contain dissolved minerals and gases. These dissolved salts provide essential minerals for plant growth and add taste to drinking water. Dissolved air, particularly oxygen, is crucial for aquatic life — fish extract this oxygen for respiration. Carbon dioxide dissolved in water enables aquatic plants to perform photosynthesis.
Let us examine how much solute can dissolve in water. A saturated solution is defined as a solution that cannot dissolve any more of the solute at a given temperature. An unsaturated solution can still dissolve more solute. A supersaturated solution holds more solute than it theoretically should at that temperature — it is unstable and will precipitate excess solute if disturbed.
Solubility is defined as the mass of solute in grams that dissolves in 100 grams of solvent to form a saturated solution at a particular temperature. For most solids, solubility increases with temperature. Potassium nitrate shows dramatic increase — from thirty-one point six grams at twenty degrees Celsius to one hundred eight grams at sixty degrees Celsius per one hundred grams of water. Sodium chloride shows only slight increase, from about thirty-six grams at twenty degrees Celsius to about thirty-seven grams at one hundred degrees Celsius per one hundred grams of water. However, calcium sulphate and calcium hydroxide show decreasing solubility with rising temperature — this is called anomalous solubility, seen in exothermic dissolution processes.
For gases, the pattern reverses. Solubility of gases decreases with increasing temperature but increases with pressure. This is described by Henry's Law: at constant temperature, the mass of gas dissolved is directly proportional to the pressure on the liquid surface. This explains why soda water fizzes when opened — the dissolved carbon dioxide escapes as pressure drops.
Now we turn to crystals and water of crystallisation. Many salts crystallise from water with definite numbers of water molecules incorporated into their structure. This fixed amount of water associated with a compound is called water of crystallisation.
Examples include blue vitriol, CuSO₄·5H₂O, which is deep blue; Epsom salt, MgSO₄·7H₂O; washing soda, Na₂CO₃·10H₂O; Glauber's salt, Na₂SO₄·10H₂O; and plaster of Paris, CaSO₄·½H₂O. When heated above one hundred degrees Celsius, these crystals lose water and become anhydrous — without water. Blue vitriol turns white when dehydrated, and turns blue again when water is added.
Substances without water of crystallisation are called anhydrous. Examples include common salt, NaCl, and potassium nitrate, KNO₃. These can be obtained from hydrated salts by direct heating, heating in dry hot air, heating under vacuum, or using dehydrating agents such as warm concentrated sulphuric acid.
Hydrated substances exhibit three important properties related to moisture.
Efflorescence is the phenomenon where a compound loses its water of crystallisation on exposure to dry air, resulting in loss of crystalline shape and crumbling into powder. The crystals crumble to powder. Washing soda and Glauber's salt are efflorescent. This occurs when vapour pressure in the hydrated crystals exceeds atmospheric vapour pressure, and is minimized during humid conditions.
Deliquescence is when certain water-soluble substances absorb moisture from air, become moist, lose crystalline form, and ultimately dissolve in the absorbed water to form a saturated solution at ordinary temperatures. Calcium chloride, CaCl₂, zinc chloride, ZnCl₂, and caustic soda, NaOH, are deliquescent. This happens when vapour pressure inside the crystals is very low compared to atmospheric vapour pressure, and is minimized during dry conditions.
Hygroscopy is when substances absorb moisture from air but do not form solutions. Concentrated sulphuric acid, phosphorus pentoxide, P₂O₅, quicklime, CaO, and silica gel are hygroscopic.
We must distinguish between drying agents and dehydrating agents.
Drying agents or desiccants remove moisture from substances without chemically reacting with them. Anhydrous calcium chloride, concentrated sulphuric acid, and phosphorus pentoxide are common drying agents. They represent physical changes.
Dehydrating agents remove chemically combined water from compounds — hydrogen and oxygen in the two to one ratio. Concentrated sulphuric acid can dehydrate blue vitriol to white anhydrous copper sulphate, and can even extract water from cane sugar, leaving behind black sugar charcoal. This represents chemical change with new substances forming.
Finally, let us discuss soft water and hard water.
Soft water readily forms lather with soap. Hard water does not. Hardness is caused by dissolved bicarbonates, chlorides, or sulphates of calcium and magnesium, which dissolve when water flows over limestone, dolomite, or gypsum beds.
Temporary hardness is caused by calcium and magnesium bicarbonates. It can be removed by boiling, which converts bicarbonates to insoluble carbonates that precipitate out. Clark's process uses slaked lime, Ca(OH)₂, to remove temporary hardness, forming insoluble carbonates that settle out.
Permanent hardness is caused by calcium and magnesium sulphates and chlorides. It cannot be removed by boiling. Washing soda, also called soda ash, Na₂CO₃, precipitates these as insoluble carbonates. For large-scale softening, the permutit process uses artificial zeolite, chemically sodium aluminium orthosilicate with formula Na₂Al₂Si₂O₈·xH₂O, simplified as Na₂P. Calcium and magnesium ions exchange with sodium ions, softening the water. The permutit is regenerated by running concentrated brine through it.
Hard water has advantages — it tastes better than soft water free from dissolved salts, provides essential calcium and magnesium for bones and teeth, and prevents lead poisoning from pipes by forming insoluble lead sulphate layer. However, it causes boiler scale formation that reduces efficiency and can cause boiler bursts, wastes soap by forming insoluble curd or scum, and leaves mineral deposits on appliances and skin. Synthetic detergents overcome the soap problem as they form soluble salts with calcium and magnesium ions, and are unaffected by water hardness.
Let us recap the key points.
First, water is a universal solvent due to its polar nature, dissolving solids, liquids, and gases to form aqueous solutions.
Second, solubility of solids generally increases with temperature, while solubility of gases decreases with temperature but increases with pressure.
Third, water of crystallisation is the fixed amount of water loosely held in chemical combination with some compounds as an integral part of their crystalline structure; hydrated substances contain it, anhydrous substances do not.
Fourth, efflorescence is loss of water to air, deliquescence is absorption of water until dissolution, and hygroscopy is absorption without dissolution.
Fifth, drying agents remove surface moisture physically, while dehydrating agents remove chemically combined water.
Sixth, temporary hardness is removed by boiling, Clark's process using lime, or adding washing soda; permanent hardness requires washing soda or the permutit process for large-scale treatment.
Water is truly remarkable — its properties sustain life, shape our environment, and enable countless chemical processes. Understanding these principles will help you appreciate this vital resource and excel in your chemistry studies. Keep exploring, keep questioning, and see you in the next lesson.