Hello, and welcome to today's chemistry lesson. We are going to explore Chapter Two: Chemical Changes and Reactions. By the end of this lesson, you will understand what makes a chemical change different from a physical change, the various ways chemical reactions occur, the four main types of chemical reactions, and how energy changes during these reactions.
Let us begin with the fundamental question: what exactly is a chemical change?
A chemical change is a permanent and irreversible change in which the chemical composition of a substance is altered. One or more new substances are formed, possessing different chemical compositions and different properties from the original substance. Unlike physical changes, chemical changes cannot be easily reversed.
There are four key characteristics you should remember. First, a chemical change is permanent and irreversible. Second, the composition of the substance undergoing change is altered. Third, a chemical change is either exothermic, where heat is evolved, or endothermic, where heat is absorbed. Fourth, and crucially, the total mass of the reactants and products remains the same — this is the Law of Conservation of Mass.
Every chemical change involves a chemical reaction. This is the process where chemical bonds in the reactants are broken, and new bonds are formed to create products. A chemical bond, remember, is simply the force that holds atoms together in a molecule or compound.
Now, what conditions are necessary for a chemical change to occur? Chemical reactions happen when particles collide, and this requires either close contact or energy supply.
Let us examine six important conditions.
First, mixing or close contact. Some reactions occur simply when two substances are mixed in their solid states. For example, when white lead nitrate and white potassium iodide are mixed, they react to form yellow lead iodide and potassium nitrate. The equation is: Pb(NO₃)₂(s) + 2KI(s) → 2KNO₃(s) + PbI₂(s).
Second, solution state. Many reactions occur when substances are dissolved in water or molten state. For instance, sodium chloride solution reacts with silver nitrate solution to form a white precipitate of silver chloride and sodium nitrate. The equation is: NaCl(aq) + AgNO₃(aq) → AgCl↓ + NaNO₃(aq).
Third, heat. Some reactions only occur when heated. Copper carbonate decomposes on heating into copper oxide and carbon dioxide: CuCO₃(s) → CuO(s) + CO₂(g). The triangle symbol above the arrow indicates heating.
Fourth, light. These are called photochemical reactions or photolysis. Molecules absorb light energy, become activated, and react rapidly. Hydrogen and chlorine combine in sunlight to form hydrogen chloride: H₂ + Cl₂ → 2HCl. Silver nitrate and hydrogen peroxide are stored in brown bottles because they decompose in light.
Fifth, electricity. These are electrochemical reactions. When electric current passes through acidulated water, it decomposes into hydrogen and oxygen: 2H₂O → 2H₂ + O₂.
Sixth, pressure. Some reactions need high pressure. Nitrogen and hydrogen combine under high pressure to form ammonia: N₂ + 3H₂ → 2NH₃.
Finally, catalysts. A catalyst is a substance that speeds up a reaction without being consumed. Potassium chlorate decomposes at 700 degrees Celsius, but with manganese dioxide as catalyst, it decomposes at just 300 degrees Celsius: 2KClO₃ → 2KCl + 3O₂. A positive catalyst accelerates reactions, while a negative catalyst or inhibitor slows them down.
How do we recognise that a chemical reaction has occurred? There are four characteristic changes to observe.
First, evolution of gas. When zinc reacts with dilute sulphuric acid, hydrogen gas evolves with effervescence — the formation of gas bubbles in a liquid. The equation is: Zn + H₂SO₄ → ZnSO₄ + H₂.
Second, change of colour. When iron is added to blue copper sulphate solution, the colour fades to light green due to formation of ferrous sulphate: Fe + CuSO₄ → FeSO₄ + Cu.
Third, formation of precipitates. When silver nitrate solution mixes with sodium chloride solution, an insoluble white solid called silver chloride forms: AgNO₃ + NaCl → AgCl↓ + NaNO₃.
Fourth, change of state. Ammonia gas and hydrogen chloride gas combine to form solid ammonium chloride: NH₃(g) + HCl(g) → NH₄Cl(s).
Now we come to the heart of this chapter: the four main types of chemical reactions.
Type one: Direct Combination or Synthesis. This is when two or more substances combine to form a single new substance. The general form is A plus B gives AB.
Two elements may combine, like iron and chlorine forming iron three chloride: 2Fe + 3Cl₂ → 2FeCl₃. An element and compound may combine, like carbon monoxide burning in oxygen to form carbon dioxide: 2CO + O₂ → 2CO₂. Or two compounds may combine, like ammonia and hydrogen chloride forming ammonium chloride: NH₃ + HCl → NH₄Cl.
When quicklime, calcium oxide, combines with water, it forms slaked lime with vigorous heat release: CaO + H₂O → Ca(OH)₂. This is used for whitewashing walls.
Type two: Decomposition. This is the opposite of combination. A compound breaks down into two or more simpler substances — elements or compounds. The general form is AB gives C plus D plus more products.
Decomposition may occur by heat, light, or electricity. Mercuric oxide decomposes on heating into mercury and oxygen: 2HgO → 2Hg + O₂. Silver chloride decomposes in sunlight: 2AgCl → 2Ag + Cl₂. Water decomposes by electric current: 2H₂O → 2H₂ + O₂.
Metal compounds decompose differently based on reactivity. Carbonates of potassium and sodium are stable to heat. Carbonates of calcium, magnesium, zinc, iron, lead, and copper decompose to metal oxides and carbon dioxide. Carbonates of mercury and silver decompose to metal, oxygen, and carbon dioxide.
Some decompositions are reversible. When heated, ammonium chloride splits into ammonia and hydrogen chloride gases, which recombine on cooling: NH₄Cl ⇌ NH₃ + HCl. This reversible decomposition by heat alone is called thermal dissociation.
Type three: Displacement. In this reaction, a more active element displaces a less active element from its compound. The general form is AB plus C gives AC plus B.
Activity series determines what displaces what. Potassium, sodium, and calcium displace hydrogen from cold water. Magnesium and zinc need heat to do so. Metals above hydrogen in the series displace hydrogen from acids.
Zinc displaces copper from copper sulphate solution: CuSO₄ + Zn → ZnSO₄ + Cu. The blue solution becomes colourless as reddish copper deposits form. Chlorine displaces iodine from potassium iodide: 2KI + Cl₂ → 2KCl + I₂.
Type four: Double Decomposition or Double Displacement. Two compounds exchange their radicals to form two new compounds. The general form is AB plus CD gives AD plus CB.
There are two subtypes. First, precipitation reactions, where an insoluble solid forms. Barium chloride and sodium sulphate form white barium sulphate precipitate: BaCl₂ + Na₂SO₄ → BaSO₄↓ + 2NaCl.
Second, neutralisation, where acid and base react to form salt and water only. The hydrogen ion from acid combines with hydroxide ion from base to form water. Sodium hydroxide and hydrochloric acid react: NaOH + HCl → NaCl + H₂O. Neutralisation has practical uses: treating bee stings, relieving acidity, neutralising acid spills, and treating acidic soil.
Finally, let us discuss energy changes in chemical reactions.
Every chemical reaction involves energy change because breaking bonds requires energy while forming bonds releases energy. These two energies differ, resulting in net absorption or release.
Exothermic reactions release heat, causing temperature rise. The products have less energy than reactants. Examples include: carbon burning in oxygen, combustion of methane, slaking of lime, neutralisation reactions, and respiration in our bodies where glucose burns to release energy: C₆H₁₂O₆ + 6O₂ → 6CO₂ + 6H₂O + Energy.
Endothermic reactions absorb heat, causing temperature fall. The products have more energy than reactants, requiring external energy supply. Examples include: formation of carbon disulphide from carbon and sulphur, formation of nitric oxide from nitrogen and oxygen at 3000 degrees Celsius, and decomposition of calcium carbonate: CaCO₃ → CaO + CO₂.
Photochemical reactions absorb light energy. Photosynthesis is the most vital example: 6CO₂ + 12H₂O → C₆H₁₂O₆ + 6H₂O + 6O₂. Electrochemical reactions absorb electrical energy, like electrolysis of water or molten sodium chloride.
Let me recap the key takeaways from this lesson.
First, a chemical change is permanent and irreversible, forming new substances with different properties, while mass is conserved.
Second, chemical reactions occur through mixing, solution, heat, light, electricity, pressure, or catalysts, and are recognised by gas evolution, colour change, precipitate formation, or state change.
Third, the four types of reactions are: combination or synthesis, decomposition, displacement, and double decomposition — which includes precipitation and neutralisation.
Fourth, exothermic reactions release heat and energy, while endothermic reactions absorb heat and require energy input.
Fifth, photochemical reactions need light, and electrochemical reactions need electricity.
Finally, catalysts speed up reactions without being consumed, and the activity series determines which elements can displace others.
You have now covered all the essential concepts of chemical changes and reactions. Understanding these fundamentals will help you predict how substances behave and transform in the world around you. Keep practising with equations, and remember to observe the conditions and characteristics we discussed. Until next time, stay curious and keep exploring the fascinating world of chemistry.