Chapter 6 — Chemical Reactions

Definition and everyday meaning.

A chemical reaction is a process in which one or more substances (reactants) change into new substances (products) with different physical and chemical properties. This change involves breaking of old chemical bonds and formation of new bonds between atoms. Chemical reactions are behind many familiar events: food cooking, wood burning, iron rusting, milk turning sour and the fizz produced when an antacid tablet dissolves in water.

How to observe a reaction.

Although we cannot see atoms directly in Class 8, we can notice the results. Typical signs that a chemical reaction has taken place include:

• Change in colour — for example, copper turning green or iron forming reddish rust.
• Evolution of a gas — bubbles form when vinegar reacts with baking soda.
• Formation of a precipitate — a solid appears when two clear solutions mix, such as when silver nitrate and sodium chloride are mixed.
• Temperature change — the mixture becomes warmer (exothermic) or colder (endothermic) without external heating or cooling.
• Change of smell or the appearance of light — burning wood gives light and new smells.

Difference from physical change.

Physical changes alter only the form or state of a substance (like melting or tearing) and are often reversible. Chemical changes produce new substances with new properties and are usually not reversible by simple physical means. For instance, burning paper cannot be undone to get the paper back because ash and gases are new products.

Recording observations and safety.

When you do experiments, write the reactants, what you see, and what you infer as products. Use proper safety: wear goggles, handle acids and reactive solids carefully, and follow your teacher’s instructions. Learning to link observable signs to the idea of bond-making and bond-breaking will help you understand reactions in later classes.

Why equations are useful.

Chemical equations give a compact way to show what happens in a reaction. They list reactants on the left, products on the right and use chemical formulas to represent substances. Equations can be written as word equations (using names) or symbol equations (using chemical formulas). Symbolic equations are precise and are the standard in science.

Steps to write and balance an equation.

• Write correct chemical formulas for all reactants and products.
• Place an arrow between reactants and products to show direction.
• Count the number of atoms of each element on both sides.
• Adjust coefficients (whole numbers in front of formulas) to make the atom counts equal on both sides.
• Do not change subscripts — changing a subscript changes the substance itself.
• Simplify coefficients to the smallest whole-number ratio if possible.

State symbols and conditions.

Add (s), (l), (g) and (aq) to show solid, liquid, gas and aqueous solution. You may also write conditions such as heat or a catalyst above or below the arrow. These details give more information about the reaction environment.

Example of balancing.

Hydrogen + oxygen → water is written H2 + O2 → H2O. Count atoms: left H = 2, O = 2; right H = 2, O = 1. Place 2 before H2O to get H2 + O2 → 2H2O, then adjust hydrogen by placing 2 before H2: 2H2 + O2 → 2H2O. Now atoms balance: H = 4 and O = 2 on both sides.

Checking work and safety.

Always re-count atoms after balancing and ensure mass is conserved. While writing equations is safe on paper, remember real reactions can release energy or harmful gases, so experimental work must be done with teacher supervision.

Definition and general pattern.

A combination or synthesis reaction is one in which two or more reactants combine to form a single product. The general pattern is A + B → AB. Many formation reactions in nature and industry follow this pattern, forming compounds from elements or simpler substances. Combination reactions are often exothermic and may release heat or light.

Types and examples.

Combination reactions occur between elements and between compounds and elements. For example, non-metal elements combine to form covalent compounds: hydrogen and chlorine combine to form hydrogen chloride: H2 + Cl2 → 2HCl. Metals and non-metals may form ionic compounds: sodium reacts with chlorine to form sodium chloride: 2Na + Cl2 → 2NaCl. Metal oxides form when metals burn in oxygen, e.g., 2Mg + O2 → 2MgO. In many of these reactions energy must be supplied to initiate the process, such as a spark to start combustion or heat to allow atoms to collide with enough energy.

Industrial importance.

Combination reactions are central to many manufacturing processes. The Haber process for ammonia is a combination of nitrogen and hydrogen to make ammonia (N2 + 3H2 → 2NH3) under controlled conditions with catalysts and pressure. The production of many salts and oxides also relies on combination reactions under engineered conditions.

Observations and safety in the lab.

When performing combination reactions in the laboratory, watch for bright flames (e.g., magnesium burning gives a dazzling white light), heat release, and formation of solid products. Some combinations are vigorous or explosive; proper protective equipment and teacher supervision are essential. Note the state symbols and conditions when writing these reactions, and always balance the resulting equations according to the conservation of mass.

Definition and general pattern.

Decomposition reactions are those in which a single compound breaks down into two or more simpler substances. The general pattern is AB → A + B. Decomposition often requires an input of energy — heat, light or electricity — to break the bonds in the original compound. These reactions are essential in chemical manufacturing and laboratory work to obtain simpler substances or elements from compounds.

Types of decomposition.

There are several common kinds. Thermal decomposition uses heat to break down a compound, such as calcium carbonate on heating: CaCO3 → CaO + CO2. Electrolytic decomposition uses electric current to separate elements from compounds, as in electrolysis of molten sodium chloride to obtain sodium metal and chlorine gas. Photochemical decomposition is driven by light; some compounds break down when exposed to sunlight. Catalytic decomposition uses a catalyst to lower energy needs, for example the use of manganese dioxide to speed up hydrogen peroxide decomposition.

Observations and uses.

Decomposition reactions often produce gases, colour changes, or solid residues. For example, heating copper carbonate gives a black copper(II) oxide and releases carbon dioxide gas. Electrolysis of water yields hydrogen and oxygen gases that can be collected separately. Decomposition is used industrially for producing lime from limestone, obtaining oxygen from hydrogen peroxide, and extracting metals via electrolysis.

Laboratory and safety notes.

Many decompositions involve high temperature or the release of gases; carry them out with ventilation and protective gear. When gas evolves, set up apparatus to collect it so mass appears conserved in a closed system. Remember to balance decomposition equations and to indicate conditions such as heat or electricity above the reaction arrow when writing symbolic equations.

Single displacement reactions.

In a single displacement reaction, one element displaces another element from a compound. The general form is A + BC → AC + B when A is more reactive than B. The reactivity of elements determines whether displacement will occur; more reactive metals displace less reactive metals from their salts. For example, zinc metal displaces hydrogen from dilute hydrochloric acid: Zn + 2HCl → ZnCl2 + H2(g). Iron placed in copper sulfate solution gives iron sulfate and copper metal: Fe + CuSO4 → FeSO4 + Cu. The displaced element often appears as a free metal.

Double displacement reactions.

Double displacement (metathesis) reactions involve exchange of ions between two compounds: AB + CD → AD + CB. These reactions commonly occur in aqueous solutions and often produce a precipitate, a gas, or a weak electrolyte. For example, mixing silver nitrate and sodium chloride solutions produces a white precipitate of silver chloride: AgNO3(aq) + NaCl(aq) → AgCl(s) + NaNO3(aq). Another example is formation of carbon dioxide when an acid reacts with a carbonate.

Predicting and testing reactions.

To predict single displacement, use the reactivity series of metals: a metal will displace a metal below it in the series. In double displacement, use solubility rules to predict whether a precipitate forms. Laboratory tests include observing colour changes, precipitate formation, gas evolution and measuring mass changes. Displacement reactions are used in metallurgy for extracting metals and in analytical chemistry to test for specific ions.

Safety and writing equations.

Carry out displacement experiments with care; some produce toxic gases or heat. Write balanced symbol equations and include state symbols such as (aq) and (s). For single displacements, remember the reactivity order to know whether the reaction will proceed spontaneously.

Definition and essential features.

Combustion is a rapid chemical reaction in which a substance reacts with oxygen and releases heat, and often light. Fuels are materials that burn to release energy; common fuels include wood, coal, kerosene, petrol (gasoline) and natural gas. The products of combustion depend on the fuel and how much oxygen is available. Complete combustion of hydrocarbons produces carbon dioxide and water, while incomplete combustion produces carbon monoxide, soot (carbon) and other harmful substances.

Complete versus incomplete combustion.

Complete combustion of a hydrocarbon like methane is represented by CH4 + 2O2 → CO2 + 2H2O, producing a clean blue flame and maximum heat. Incomplete combustion occurs when oxygen supply is limited; for example, 2CH4 + 3O2 → 2CO + 4H2O or C (soot) may form. Incomplete combustion gives less energy and produces carbon monoxide (CO), which is a colourless, odourless and toxic gas.

Energy and environmental aspects.

Fuels are chosen for high calorific value (energy released per kilogram). Combustion in engines and power plants converts chemical energy into mechanical or electrical energy. However, combustion also produces pollutants such as CO2 (a greenhouse gas), NOx and particulates. These pollutants affect air quality and contribute to climate change, so efficient combustion and pollution control are important.

Practical safety and prevention of hazards.

Ensure good ventilation and correct burners to avoid incomplete combustion and buildup of CO. Use flame arrestors, maintain stoves and engines, and keep flammable liquids in safe containers. In experiments, use proper apparatus and keep flammable materials away from open flames. Understanding combustion helps you choose safer and cleaner fuels and to recognise the signs of incomplete combustion such as yellow smoky flames and smell.

Basic idea.

Oxidation and reduction are paired processes that describe the transfer of oxygen, hydrogen or electrons in chemical reactions. For Class 8 we use simple rules: oxidation is gain of oxygen or loss of hydrogen; reduction is loss of oxygen or gain of hydrogen. Since electrons are transferred too, oxidation and reduction always happen together — when one substance is oxidised another is reduced.

Everyday examples and simple tests.

Rusting of iron is an oxidation process where iron combines with oxygen in the presence of water to form rust. Burning carbon to form carbon dioxide is another example of oxidation: C + O2 → CO2. In contrast, when hydrogen removes oxygen from copper(II) oxide it reduces it to metallic copper: CuO + H2 → Cu + H2O. Here oxygen moves from copper to hydrogen; copper is reduced and hydrogen is oxidised.

Oxidising and reducing agents.

The oxidising agent causes oxidation by accepting electrons or supplying oxygen; it itself becomes reduced. Oxygen is a common oxidising agent. The reducing agent causes reduction by donating electrons or hydrogen; it itself becomes oxidised. Hydrogen or carbon often act as reducing agents in extraction and laboratory reactions.

Importance and simple identification.

Understanding oxidation and reduction helps explain corrosion, bleaching, combustion and metal extraction. In simple classroom tests, note that a substance that removes oxygen from another (or removes hydrogen) is a reducing agent. Always write the two half-reactions if needed: one showing oxidation and the other showing reduction, then combine them to see the overall change. This introduces you to the idea of electron movement, which you will study more formally in higher classes.

What are acids and bases?

Acids are substances that taste sour, react with metals to produce hydrogen and turn blue litmus red. Bases or alkalis taste bitter, feel slippery and turn red litmus blue. In water, acids produce hydrogen ions (H+) and bases produce hydroxide ions (OH-). Examples include hydrochloric acid (HCl) and sodium hydroxide (NaOH). Many household substances are acidic (lemon juice) or basic (soap solution).

Neutralisation reaction.

Neutralisation is the reaction of an acid with a base to form a salt and water. The general form is Acid + Base → Salt + Water. For example, HCl + NaOH → NaCl + H2O. Neutralisation often releases heat (exothermic) and is used to treat acid spills, adjust soil acidity and prepare salts in the laboratory.

Preparing salts by neutralisation and other methods.

Salts can be prepared by neutralising an acid with a base, by reacting a metal with an acid, or by reacting an acid with a metal oxide. For example, copper oxide reacts with sulfuric acid to form copper sulfate and water: CuO + H2SO4 → CuSO4 + H2O. When acids react with carbonates, carbon dioxide gas is released: 2HCl + Na2CO3 → 2NaCl + CO2 + H2O. In classroom preparations, indicators and careful addition of reagents help reach the neutral point before crystallising the salt.

Indicators and pH.

Indicators such as litmus show whether a solution is acidic or basic. Universal indicator gives a range of colours for pH values from 0 to 14: acids below 7, neutral around 7, bases above 7. For more accurate measurement, a pH meter is used. Understanding pH is important for agriculture, digestion, water treatment and many chemical processes.

Energy changes in chemical reactions.

Chemical reactions involve changes in energy. If a reaction releases heat to the surroundings it is called exothermic; if it absorbs heat from the surroundings it is endothermic. These energy changes can be felt as temperature increases or decreases during the reaction and are important for identifying the type of reaction.

Characteristics and examples.

Exothermic reactions give out heat and sometimes light. Combustion of fuels and many neutralisation reactions are exothermic. For example, the burning of a candle and the neutralisation of HCl by NaOH both raise the temperature of the surroundings. Endothermic reactions require continuous input of heat to proceed; thermal decomposition reactions like CaCO3 → CaO + CO2 are endothermic because heat must be supplied to break chemical bonds.

Energy profile diagrams.

On an energy profile, reactants start at one level and products at another. For exothermic reactions products lie at a lower energy level than reactants and the difference is released as heat. For endothermic reactions products lie at a higher energy level and energy must be absorbed. Both types require an initial activation energy to start the reaction, seen as a hump on the diagram; catalysts lower this hump.

Practical measurement and examples.

In the lab, measure temperature before and after mixing reagents to see whether heat is released or absorbed. Dissolving ammonium nitrate in water is endothermic and used in cold packs; dissolving sodium hydroxide in water is exothermic and releases heat. Understanding these energy changes is important in industrial process design, cooking, and safe handling of reactive chemicals.

Understanding reaction rate.

The rate of a chemical reaction describes how fast reactants are converted to products. For Class 8 we focus on four main factors that change the rate: concentration, surface area (particle size), temperature and catalysts. Each factor works by affecting how often particles collide and how much energy they have when they collide.

Concentration and pressure.

Higher concentration of reactants in a solution means more particles in the same volume, so collisions happen more frequently and the reaction speeds up. For gases, increasing pressure effectively increases concentration and raises the reaction rate. For example, more concentrated acid reacts faster with a metal than dilute acid.

Surface area and particle size.

When a solid reactant is made into smaller pieces or powder, its surface area increases and more particles are exposed for reaction. This increases the frequency of collisions with a liquid or gas and speeds the reaction. For instance, powdered zinc reacts with acid faster than a large lump of zinc.

Temperature and kinetic energy.

Raising temperature gives particles more kinetic energy, so they move faster and collide more often and more energetically. A larger fraction of collisions have enough energy to overcome the activation energy barrier, so rate increases. A rough rule of thumb is that many reaction rates double for a 10°C rise, though actual change depends on the reaction.

Catalysts and inhibitors.

Catalysts increase the rate by providing an alternative pathway with lower activation energy and are not consumed in the reaction. Enzymes are biological catalysts that work inside living cells at moderate temperatures. Inhibitors slow down reactions and catalyst poisons deactivate catalysts. Industrial processes rely heavily on catalysts to make reactions faster and more economical.

Practical observations.

In experiments, measure how long it takes for a change (like gas evolution or colour change) to occur under different conditions to compare rates. Plotting amount of product versus time shows how steeply the curve rises for faster reactions. Always control variables carefully so only the intended factor changes.

The law and its meaning.

The law of conservation of mass states that mass is neither created nor destroyed in a chemical reaction. In a closed system the total mass of reactants equals the total mass of products. This principle underlies the need to balance chemical equations: if atoms are conserved then mass must be conserved too.

Why balancing matters.

Balancing equations ensures the same number of each type of atom appears on both sides. For example, in the reaction 2H2 + O2 → 2H2O there are four hydrogen atoms and two oxygen atoms on both sides. If an equation is not balanced it would imply atoms disappeared or appeared, contradicting conservation of mass.

Simple experimental demonstrations.

You can demonstrate conservation of mass by performing a reaction in a sealed container where no gas escapes. For instance, place sodium bicarbonate and vinegar in a sealed flask with a delivery tube leading to a closed bottle; mass of the sealed system remains constant though gases form inside. If you allow gas to escape, the mass of the open system appears to decrease; this shows the importance of a closed system when testing the law.

Practical tips for balancing.

When balancing, change coefficients but never subscripts. Treat polyatomic ions that appear unchanged on both sides as single units to simplify balancing. Recount each element after adding coefficients. Conservation of mass also helps in laboratory calculations and in real-life chemical manufacturing where input and output masses must be tracked for yield and safety.

Role of catalysts.

Catalysts are substances that increase the rate of a chemical reaction without being consumed in the process. They do this by providing an alternative reaction pathway with a lower activation energy, so a higher proportion of collisions between particles are effective. Catalysts are invaluable in industry to speed up reactions, reduce energy costs and improve yields.

Types and examples.

There are heterogeneous catalysts (in a different phase than reactants) such as solid platinum used in catalytic converters, and homogeneous catalysts (in the same phase) such as acids used in esterification. Enzymes are biological catalysts that allow reactions in living organisms to proceed at body temperature and are highly specific for their substrates. For example, manganese dioxide speeds up decomposition of hydrogen peroxide in the laboratory without being used up; the catalyst appears above the reaction arrow in an equation.

Inhibitors and catalyst poisons.

Inhibitors slow down reactions and can be useful where a slower rate is desired, like in preservatives for food. Catalyst poisons are substances that deactivate catalysts by binding to their active sites; for example sulfur compounds can poison metal catalysts. Understanding poisons is important for maintaining industrial processes and ensuring catalysts remain effective.

Practical considerations and safety.

Catalysts are not written among reactants because they are recovered unchanged; typically they are written above or below the arrow in a chemical equation. Handling catalysts requires care because some are toxic or expensive, and many must be recovered and recycled. In biology, enzymes operate under narrow temperature and pH ranges and can be denatured if conditions are not kept suitable.

What are indicators?

Indicators are substances that change colour to show whether a solution is acidic or basic. Simple indicators include litmus paper: blue litmus turns red in acid and red litmus turns blue in base. Natural indicators, such as red cabbage extract or turmeric, change colour because their molecules alter structure with changing pH. Universal indicator is a mixture that gives a spectrum of colours across the pH scale and helps estimate pH values roughly.

Understanding the pH scale.

pH is a scale from 0 to 14 that indicates acidity or alkalinity of an aqueous solution. pH 7 is neutral (pure water), pH less than 7 is acidic and pH greater than 7 is basic. pH is related to hydrogen ion concentration: lower pH means higher concentration of H+ ions. While indicators give qualitative or approximate pH, pH meters give more accurate numerical values and are used where precision matters.

Uses and demonstrations.

Indicators are used in titrations to find the equivalence point where acid and base neutralise each other. In classrooms, you can test household substances: lemon juice appears acidic, soap solution appears basic. Universal indicator will show red for strong acids, green for neutral solutions and violet for strong bases. Take care when using coloured solutions, as the colour of the solution may affect indicator visibility.

Limitations and safety.

Natural indicators may work over a narrow pH range and give ambiguous results with coloured samples. Indicators do not give exact pH values. For handling acids and bases, use gloves and goggles. Understanding pH is useful in agriculture (soil pH), swimming pool maintenance, medicine (stomach acidity) and environmental monitoring of water bodies.

Corrosion: what and why.

Corrosion is the slow chemical destruction of metals by reaction with substances in their environment, most commonly oxygen and water. Rusting of iron is a familiar example where iron reacts with oxygen and water to form hydrated iron(III) oxide (rust). Corrosion weakens structures, damages tools and causes large economic loss. The rate of corrosion increases in the presence of salts, acids and pollutants.

Prevention methods.

Several methods reduce or prevent corrosion. Painting or coating metal surfaces blocks contact with air and moisture. Oiling or greasing movable parts prevents moisture contact. Galvanising — coating iron or steel with zinc — protects because zinc is more reactive and corrodes in preference to iron (sacrificial protection). Electroplating covers the metal with a protective layer such as chromium. Cathodic protection uses a more reactive metal electrically connected to the iron structure to act as a sacrificial anode; this is used for pipelines and ship hulls.

Rancidity of fats and oils.

Rancidity is the oxidation of fats and oils in foods leading to unpleasant tastes and smells. Light, heat and air speed up rancidity. Preventive measures include storing food in airtight, dark containers, refrigeration, using antioxidants (like vitamin E) and vacuum packaging. Proper storage prolongs shelf life and keeps food safe and palatable.

Environmental and practical notes.

Preventing corrosion and rancidity saves resources and reduces waste. Choose suitable protective methods depending on cost and application: paint and coatings are economical for many structures, while galvanising is chosen where long-term protection is needed. Regular maintenance, prompt repair of coatings and controlling exposure to salt (for example near coasts) help reduce damage. For food, proper packaging and storage are essential to maintain quality.