Hello, and welcome to today's chemistry lesson. In this session, we will explore the fascinating world of chemical reactions — what they are, how we recognise them, and the different types that occur all around us. We will learn about combination and decomposition reactions, displacement and double displacement reactions, and how energy changes during these processes. We will also discover the nature of oxides and why some metals are more reactive than others.
Let us begin with the fundamental question: what exactly is a chemical reaction? A chemical reaction is any chemical change in matter that involves the transformation of one or more substances into entirely new substances with different properties. The substances that undergo change are called reactants, and the new substances formed are called products.
Consider what happens when methane burns in oxygen. The reactants are methane, CH₄, and oxygen, O₂. After the reaction, we obtain carbon dioxide, CO₂, water vapour, H₂O, and energy in the form of heat and light. Similarly, when sodium hydroxide reacts with hydrochloric acid, we get sodium chloride and water.
At the molecular level, a chemical reaction involves breaking chemical bonds between atoms and rearranging those atoms to form new bonds. A chemical bond is simply the attractive force that holds atoms or ions together in a molecule.
How can we tell when a chemical reaction has occurred? There are several characteristic changes we can observe.
First, evolution of gas. When zinc reacts with dilute sulphuric acid, hydrogen gas is produced with visible effervescence — bubbles forming in the liquid. The equation is: zinc plus sulphuric acid gives zinc sulphate plus hydrogen. Or, when sodium sulphite reacts with dilute hydrochloric acid, sulphur dioxide gas with a suffocating odour is liberated.
Second, change of colour. Drop some iron pieces into blue copper sulphate solution, and watch the solution gradually turn light green as iron displaces copper. The reddish-brown copper deposits on the iron. Alternatively, when blue copper sulphate reacts with hydrogen sulphide gas, a black solid of copper sulphide forms.
Third, formation of precipitates. When silver nitrate solution mixes with sodium chloride solution, a white insoluble solid called silver chloride settles at the bottom. Similarly, adding ferrous sulphate to sodium hydroxide produces a dirty green precipitate of ferrous hydroxide.
Fourth, change of state. Hydrogen sulphide gas and chlorine gas react to form solid sulphur and hydrogen chloride gas. Ammonia gas and hydrogen chloride gas combine to form solid ammonium chloride. We represent these states using (s), (l), (g), and (aq) for aqueous solutions.
Fifth, change in energy. Reactions either release energy, usually as heat and light, or absorb energy from the surroundings.
For chemical reactions to occur, certain conditions are often necessary.
Close contact between reactants is essential. Sodium metal explodes on contact with cold water, forming sodium hydroxide and hydrogen.
Many reactions proceed faster when reactants are in solution form, as this brings molecules into contact more readily. Adding dilute sulphuric acid to barium chloride solution immediately produces a white precipitate of barium sulphate.
Heat energy drives many reactions. Iron and sulphur powders mixed at room temperature do not react, but when heated, they combine to form iron sulphide.
Light energy enables photochemical reactions. Photosynthesis is the most important example: carbon dioxide and water, in the presence of chlorophyll and sunlight, produce glucose and oxygen.
Electricity powers electrochemical reactions. When electric current passes through acidulated water, it decomposes into hydrogen and oxygen gases. Note that pure water conducts electricity poorly, so we add a small amount of acid, alkali, or salt.
High pressure facilitates certain reactions. Nitrogen and hydrogen combine at 200 atmospheres pressure to form ammonia.
Finally, catalysts alter reaction rates without being consumed themselves. A positive catalyst speeds up reactions. Manganese dioxide lowers the decomposition temperature of potassium chlorate from 700 degrees Celsius to about 300 degrees Celsius. Finely divided iron catalyses ammonia production from nitrogen and hydrogen. Promoters like molybdenum enhance catalyst efficiency but cannot work alone. A negative catalyst or inhibitor slows reactions down. Phosphoric acid, for instance, slows the decomposition of hydrogen peroxide.
In living organisms, enzymes — complex protein molecules — act as biological catalysts. Without enzymes, digestion would take decades; with enzymes, it completes in mere hours.
Now let us examine the four main types of chemical reactions.
First, combination or synthesis reactions, where two or more substances unite to form a single product. Two elements may combine: magnesium burns in oxygen with a dazzling white flame to form magnesium oxide. An element and a compound may combine: carbon monoxide burns in oxygen to form carbon dioxide. Two compounds may combine: ammonia and hydrogen chloride gases form solid ammonium chloride.
Second, decomposition reactions, where a compound breaks down into simpler substances, usually on heating. Mercuric oxide decomposes to mercury and oxygen. Water electrolyses into hydrogen and oxygen. Potassium nitrate yields potassium nitrite and oxygen. Calcium carbonate, or limestone, decomposes to calcium oxide and carbon dioxide. When you heat lead nitrate, you observe reddish-brown nitrogen dioxide gas, colourless oxygen that reignites a glowing splint, and pale yellow lead monoxide remaining as residue.
Third, displacement reactions, where a more reactive element replaces a less reactive one in a compound. A more reactive metal displaces a less reactive metal from its salt solution. Zinc displaces copper from copper sulphate, turning the blue solution colourless while depositing reddish-brown copper. Iron similarly displaces copper, producing green iron sulphate solution. Metals above hydrogen in the reactivity series displace hydrogen from acids. Zinc reacts with dilute hydrochloric or sulphuric acid to produce hydrogen gas. Highly reactive sodium and potassium displace hydrogen from water violently, forming hydroxides. Among non-metals, chlorine displaces bromine from potassium bromide, or iodine from potassium iodide.
Fourth, double displacement reactions, where two compounds exchange ions in aqueous solution. Precipitation reactions produce an insoluble solid. Barium chloride and sodium sulphate yield white barium sulphate precipitate plus sodium chloride. Neutralisation reactions occur when acids react with bases to produce salt and water only. Sodium hydroxide neutralises hydrochloric acid to form sodium chloride and water. Zinc oxide, a base, reacts with nitric acid to give zinc nitrate and water.
Indicators help us identify acids and bases. Red litmus turns blue in alkaline solutions; blue litmus turns red in acidic solutions. Methyl orange is red or pink in acid, yellow in base. Phenolphthalein is colourless in acid, pink in base.
Neutralisation has practical importance. Antacids like milk of magnesia relieve indigestion by neutralising excess stomach acid. Toothpastes contain bases to neutralise mouth acids. Bee stings, being acidic, are treated with basic calamine or baking soda; wasp stings, being alkaline, are treated with acidic vinegar or lemon juice. Acidic soils are treated with lime to restore neutrality for healthy plant growth.
The metal activity series arranges metals in decreasing order of chemical reactivity. Potassium sits at the top as most reactive; platinum anchors the bottom as least reactive. This series helps predict displacement reactions: more reactive metals displace less reactive ones from their salts. Metals above hydrogen displace hydrogen from water or acids; those below cannot.
Energy changes classify reactions as exothermic or endothermic.
Exothermic reactions release heat, raising the temperature of surroundings. Burning carbon in oxygen produces carbon dioxide and heat. Adding water to quicklime generates enough heat to boil the water, forming slaked lime. Respiration, rusting, and burning fuels are all exothermic. All neutralisation reactions release heat — you can feel the beaker warming when acid and base react.
Endothermic reactions absorb heat, causing temperature to fall. Nitrogen and oxygen combine to form nitric oxide only at about 3000 degrees Celsius. Calcium carbonate decomposes at roughly 1000 degrees Celsius. Dissolving ammonium chloride in water makes the beaker feel cold — heat is absorbed from the surroundings.
Finally, let us understand oxides — compounds containing oxygen combined with metals or non-metals.
Metallic oxides form when metals react with oxygen. Sodium forms sodium oxide; calcium forms calcium oxide; copper forms copper oxide. Heating metal carbonates, nitrates, or sulphates also produces metallic oxides.
Most metallic oxides are basic oxides. Some dissolve in water forming alkalis: sodium oxide gives sodium hydroxide; potassium oxide gives potassium hydroxide. Basic oxides react with acids to produce salt and water. Calcium oxide reacts with hydrochloric acid to give calcium chloride and water.
Amphoteric oxides behave as both acids and bases. Zinc oxide reacts with hydrochloric acid to form zinc chloride and water, showing basic character. It also reacts with sodium hydroxide to form sodium zincate and water, showing acidic character. Lead monoxide and aluminium oxide similarly exhibit dual behaviour.
Non-metallic oxides form when non-metals combine with oxygen. Carbon burns to carbon dioxide; sulphur burns to sulphur dioxide.
Most non-metallic oxides are acidic oxides. They dissolve in water forming acids: carbon dioxide gives carbonic acid; sulphur dioxide gives sulphurous acid; sulphur trioxide gives sulphuric acid. Acidic oxides react with bases to produce salt and water.
Neutral oxides like carbon monoxide, nitric oxide, and water neither change indicator colours nor react with acids or bases to form salts.
Let us recap the key points from today's lesson.
First, a chemical reaction involves breaking and making of chemical bonds to form new substances with different properties.
Second, chemical reactions are recognised by evolution of gas, colour change, precipitate formation, state change, and energy change.
Third, the four main reaction types are: combination, where substances unite; decomposition, where compounds break down; displacement, where more reactive elements replace less reactive ones; and double displacement, where compounds exchange ions.
Fourth, exothermic reactions release heat while endothermic reactions absorb heat.
Fifth, catalysts alter reaction rates without being consumed; enzymes are biological catalysts.
Sixth, the metal activity series predicts which metals can displace others from compounds.
Seventh, metallic oxides are generally basic, non-metallic oxides are generally acidic, some metallic oxides are amphoteric, and certain oxides are neutral.
That brings us to the end of our lesson on chemical reactions. I hope you now feel confident in understanding how substances transform around us, from the burning of fuels to the digestion of food. Keep observing these reactions in your daily life, and remember — chemistry is happening everywhere, all the time. Until next time, stay curious and keep exploring the wonderful world of science.