Chapter 2 — Physical and Chemical Changes

Introduction to change in matter
Matter is anything that has mass and occupies space. When we talk about changes in matter, we refer to what happens to the arrangement, motion or bonding of its particles. Changes can involve only physical aspects such as shape, size or state, or they can involve chemical transformations where the substances themselves are altered and new materials form.

Classifying changes
There are two broad classes of change. Physical changes affect how matter looks or is arranged but do not alter the substance’s identity. Examples include melting, freezing, cutting and dissolving. Chemical changes, also called chemical reactions, produce new substances with different properties. For instance, burning wood yields ash, smoke and gases which are chemically different from the original wood.

How to observe changes
Careful observation is essential. Notice colour changes, smell, temperature change, the appearance of bubbles, formation of solids or changes in texture. These clues help decide whether a change might be chemical. Keep a notebook and record time, conditions (like warmth or stirring) and measurements where possible. Observations should be described precisely: what was done, what was seen, and when it occurred.

Why this matters
Understanding these basic ideas helps explain everyday events — why ice melts, why food sours, why metals rust — and prepares students for practical work. It helps make safe choices at home and school, for example when heating substances or disposing of chemicals. The ability to classify and explain changes builds logical thinking that is useful in science and daily life.

Safety and ethics
Always respect safety: wear goggles when needed, avoid tasting unknown substances, and follow teacher instructions. Be honest in recording results and thinking about environmental impacts such as pollution from chemical processes. These habits form good scientific practice.

What are physical changes?
Physical changes alter the physical form or appearance of a substance but do not change its chemical composition. The particles remain the same kind, though their arrangement or motion may change. Common physical changes include changes of state (melting, freezing, boiling, condensation and sublimation), changes in shape or size (cutting, bending), and mixing/dissolving.

Changes of state explained
Particles in solids are closely packed and vibrate about fixed positions. When heat is added, particles gain energy and can move more freely, leading to melting and forming a liquid. Further heating can cause particles to escape into the gas phase by evaporation or boiling. Cooling removes energy and causes gas to condense or liquids to freeze. These changes depend on temperature and pressure and are typically reversible under suitable conditions.

Dissolving and mixtures
Dissolving is a physical process in which a solute (like sugar or salt) disperses at the particle level in a solvent (like water). The resulting solution is a homogeneous mixture. The solute can often be recovered by physical methods such as evaporation or crystallisation. Mixing two solids or liquids without chemical reaction also produces a mixture that can be separated physically.

Mechanical changes
Cutting, grinding and stretching change the form or size of an object. For example, grinding a rock into powder changes texture and surface area but not chemical identity. Stretching an elastic band is reversible within its elastic limit because bonds are not permanently altered.

Applications and observations
Physical changes are used every day: freezing to preserve food, boiling to cook, filtering to separate solids, and dissolving to make solutions. Observational skills are important: note temperature, time taken, and whether the change can be reversed easily. Safety matters too — for instance, handle hot liquids with care and use protective equipment when needed.

Nature of chemical changes
Chemical changes happen when atoms are rearranged, bonds are broken and new bonds form, producing substances different from the starting materials. These reactions change chemical identity: reactants turn into products with new properties, such as different colours, melting points, solubilities and smells. Chemical changes often involve energy change — either released (exothermic) or absorbed (endothermic).

How to recognise chemical changes
Certain observations point to chemical change: persistent colour change, evolution of gas bubbles not caused by heating alone, formation of an insoluble solid (precipitate), change in temperature without external heating or cooling, and appearance of a new smell. While each sign is helpful, none alone proves a chemical change; combine several observations for stronger evidence.

Role of indicators
Indicators are substances that change colour when they undergo a chemical response to the acidity or basicity of a solution. Litmus, phenolphthalein and universal indicator are common examples used to reveal acidic or alkaline conditions. Indicators themselves undergo a reversible chemical change that provides a visible sign. Other chemical tests use reagents to form characteristic precipitates or colour reactions that confirm the presence of particular ions.

Everyday chemical changes
Common chemical changes include combustion (burning), rusting, cooking (proteins denature and new compounds form), fermentation (sugars converted to alcohol and carbon dioxide), and decomposition (food spoils). Many of these have important uses — for example, fermentation in making bread and yogurt — and also safety implications, such as the risk of fire or formation of harmful gases.

Laboratory practice
When investigating chemical changes, carry out tests carefully: use small amounts of reagents, wear protection, and avoid inhaling gases. Record observations, use controls where possible, and repeat tests to confirm results. Understanding both the signs and methods of testing prepares students to explore chemistry further.

Recognising chemical change
To decide whether a change is chemical, observe several features. Colour change may indicate a new substance. Gas production is shown by bubbling or fizziness; if the gas has a characteristic reaction (for example, turning limewater milky), it helps identify the gas. Formation of a precipitate means an insoluble solid has formed in a solution. Temperature change without external heating signals an energy change within the system. A new odour or taste (do not taste in the lab) may also point to chemical transformation. Use more than one observation to make a strong conclusion.

Common laboratory tests
Some simple, reliable tests are used at school. Limewater (calcium hydroxide solution) turns milky on passing carbon dioxide, showing CO2 formation. Silver nitrate solution gives a white precipitate with chloride ions; this test can detect soluble chloride salts. Bubbles of hydrogen gas may be tested by bringing a burning splint near the gas — a faint ‘pop’ indicates hydrogen. Universal indicator or litmus paper shows acidity or alkalinity by clear colour changes. These tests should be performed with care and small quantities.

Temperature as evidence
Exothermic reactions release heat and feel warm; examples are many combustion reactions and mixing strong acids with water. Endothermic reactions absorb heat and feel cold, such as dissolving ammonium nitrate in water. Always use a thermometer for accurate readings and gloves when handling hot or cold containers.

Interpreting mixed signals
Some physical processes can show gas bubbles (boiling) or colour changes (dissolving dyes) without chemical change. Therefore, confirm chemical change by combining tests: for example, if bubbling occurs and limewater turns milky, this together confirms CO2 production by a chemical reaction rather than simple boiling.

Safe practice
Always wear eye protection, work under teacher supervision, and dispose of reagents properly. Use small amounts, keep flammable substances away from flames, and label your samples. Accurate observations and repeat tests make conclusions reliable.

Understanding reversibility
Reversible changes are those that can be returned to the original state by simple physical means. Many physical changes are reversible because the underlying particles and bonds are not permanently altered. For example, water frozen into ice can be melted back to water; salt dissolved in water can be recovered by evaporation. Reversibility often depends on conditions such as temperature, pressure and time.

Irreversible changes explained
Irreversible changes cannot easily be reversed to obtain the original substances. Most chemical changes are irreversible without further chemical processes. Examples include burning paper, baking a cake and rusting of iron. In these cases, new substances with different properties form, and reversing would require additional chemical reactions or complex industrial processes.

Borderline cases
Some changes can be reversible under special conditions but appear irreversible in everyday life. For example, a stretched elastic band may regain its shape if not overstretched, but if broken it cannot. A crushed metal can sometimes be remelted and reformed (a physical process), but if the metal oxidises badly, the oxide layer may make simple recovery difficult. These examples show that reversibility depends on the nature of change and the available methods for reversal.

How to decide in experiments
When performing an experiment, ask: can I get the original substances back by physical means (cooling, removing solvent, filtration)? If yes, the change is reversible. If only by chemical treatment or not at all, the change is likely irreversible. Record observations and, if possible, test recovery methods such as evaporation to obtain dissolved solids back.

Practical importance
Knowing whether a change is reversible helps in everyday decisions: preserving food by freezing (reversible), disposing burnt waste (irreversible), choosing repair methods for materials, and designing recycling processes. It also helps students plan experiments and predict outcomes.

Why separate mixtures?
Many mixtures occur in daily life and experiments: mud in water, tea leaves in brewed tea, salt in seawater. Separating mixtures recovers useful substances, removes impurities and helps study components. Filtration, sedimentation and decantation are simple, commonly used physical methods that rely on differences in particle size, density and solubility.

Sedimentation and decantation
Sedimentation uses gravity to let heavier particles settle at the bottom of a container over time. For muddy water, if left undisturbed, solids gather as sediment. Decantation follows: the clear liquid on top is carefully poured away without disturbing the settled solids. This method is quick and easy but not perfect: fine particles may stay suspended or be poured off accidentally. It is useful as a first step before finer separation.

Filtration explained
Filtration forces a mixture through a porous medium that traps solid particles but allows liquid to pass. Filter paper, cloth or sand filters are common. In a laboratory funnel with filter paper, the residue is the solid left on the paper; the filtrate is the liquid collected below. Filtration separates insoluble solids from liquids effectively and is used for preparing clear solutions, collecting precipitates and making drinks like tea and coffee.

Choosing methods and combined approaches
Combine methods when needed: allow coarse particles to settle (sedimentation), then decant the liquid and pass it through filter paper to catch finer particles. For very fine suspensions, centrifugation (not part of this class) speeds up sedimentation. Always consider particle size and whether the solid is soluble. Filtration cannot separate dissolved substances; evaporation or distillation is required for that purpose.

Practical tips and safety
Pour slowly during decantation, tilt the container and use a clean receiving vessel. Wet the filter paper before filtering to help it stick to the funnel. Label samples, clean apparatus after use, and dispose residue according to instructions. Wear gloves and goggles when handling harmful substances, and perform separations under teacher supervision.

Evaporation basics
Evaporation is the process by which a liquid turns into vapour at temperatures below its boiling point, often from the surface. In separation, evaporation removes a solvent (usually water) and leaves dissolved solids behind. This simple method is commonly used to obtain salt from seawater or to concentrate solutions. Evaporation rate depends on temperature, surface area and air movement.

Crystallisation as a purification technique
Crystallisation recovers a solute from its solution in a purer form. A saturated or nearly saturated solution is prepared and then cooled slowly or allowed to evaporate gently. As the solubility decreases, solute particles arrange themselves into a regular lattice to form crystals. Slow evaporation or cooling favours larger, purer crystals because impurities remain dissolved in the mother liquor.

Steps in crystallisation
First dissolve the solid completely in the minimum amount of hot solvent to make a concentrated solution. Filter the hot solution to remove insoluble impurities. Allow the filtered solution to cool slowly; crystals form as temperature drops. Separate crystals by filtration and wash them with a small amount of cold solvent to remove adhering impurities. Dry the crystals carefully.

Practical examples and uses
Making rock sugar at home uses slow crystallisation from a sugar syrup. Chemical industries use crystallisation to purify substances like salts and organic compounds. In the lab, recrystallisation is a standard method for purifying solids. Evaporation is also used in cooking (reducing sauces) and in obtaining salt pans from seawater.

Precautions
Avoid overheating which can decompose solutes. Use gentle heat and supervise hot liquids. Label containers, work on stable surfaces, and keep flammable materials away from open flames. Collect and dispose of residues responsibly. Record temperatures and times to reproduce results reliably.

Why distil?
Distillation separates liquids based on differing boiling points. When a mixture is heated, the more volatile component vaporises first. That vapour is then cooled in a condenser to form liquid (distillate) which is collected separately. Distillation is essential when one component is to be recovered in pure form, such as obtaining drinking water from impure sources or separating alcohol from a fermented mixture.

Simple distillation apparatus and operation
A typical setup includes a round-bottom flask for the liquid mixture, a heat source, a thermometer to monitor vapour temperature, a condenser where vapour cools to liquid, and a receiver to collect the distillate. Heat gently so that vaporisation is controlled. Cooling water enters the condenser from the lower end and leaves from the upper end to maintain effective condensation. Collect fractions when the temperature rises to the next component’s boiling point.

Types of distillation
Simple distillation suits mixtures with large boiling point differences or when one component is non-volatile (e.g., saltwater). Fractional distillation, using a fractionating column, separates liquids with closer boiling points by providing repeated vaporisation-condensation cycles, improving purity. Steam distillation is used for heat-sensitive substances like essential oils.

Applications and limitations
Industries use distillation for petroleum refining, alcohol production and water purification. However, distillation cannot separate mixtures of substances with identical boiling points and requires energy. For very small boiling point differences, fractional distillation or other methods may be needed to achieve good separation.

Safety and good practice
Never heat a closed system; always allow vapour to escape through the condenser. Secure apparatus with clamps and stands to avoid spills. Ensure good ventilation and keep flammable liquids away from open flames. Monitor temperature and cooling water flow, and handle hot glassware with care. Dispose of residues properly and follow teacher instructions.

What is rusting?
Rusting is a specific kind of corrosion that affects iron and its alloys. It is a chemical reaction between iron, oxygen and water that produces hydrated iron(III) oxide—commonly called rust. Rusting changes both appearance and strength: a rusty metal becomes flaky, weak and less useful.

How rust forms
Rusting is an electrochemical process. Iron atoms lose electrons (oxidation) and oxygen gains electrons (reduction) in the presence of water or moisture. The process is accelerated by acids and salts because they increase the conductivity of the water film on metal surfaces, allowing electrons to move more easily and reactions to proceed faster. Salt spray in coastal regions makes rusting an important practical problem.

Ways to prevent rust
Prevention aims either to stop contact with air and water or to protect the metal surface. Painting or coating with oil forms a barrier. Galvanising coats iron with zinc; zinc corrodes in preference to iron and protects it sacrificially. Stainless steel contains chromium which forms a thin, protective oxide layer that prevents further corrosion. Regular maintenance, drying and using inhibitors also reduce rust.

Other forms of corrosion
Corrosion affects many metals differently. Copper develops a green patina (basic copper carbonate) which can protect the metal. Silver tarnishes to form black silver sulphide when exposed to sulphur compounds. Aluminium forms a thin oxide layer that prevents further corrosion in many conditions. Understanding each metal’s behaviour helps choose materials for construction, utensils and outdoor use.

Importance and safety
Rust weakens structures, increases maintenance costs, and can cause failures in bridges, vehicles and tools. Preventive treatments and appropriate material choice are essential in engineering and daily life. Handling rusty objects can cause tetanus risk if a wound occurs, so wear gloves and keep tetanus immunisations up to date when dealing with corroded metal.

Definition and basics
Combustion is a chemical reaction between a fuel and oxygen that releases heat and often light. It is an oxidation process where the fuel (which may be a hydrocarbon, wood, coal or gas) reacts with atmospheric oxygen. The reaction can be rapid, producing a visible flame, or slower without flame (smouldering). Combustion is central to cooking, heating and powering engines.

Complete vs incomplete combustion
Complete combustion happens when oxygen supply is ample; fuel burns to carbon dioxide and water, releasing maximum energy and usually producing a blue flame. Incomplete combustion occurs when oxygen is limited; it produces carbon monoxide (a poisonous gas), carbon (soot) and other partially oxidised products with less energy and a yellow, smoky flame. Ensuring adequate oxygen and clean-burning fuels reduces pollution and danger.

Signs and tests
Combustion gives out heat and light, produces gases and may leave solid residues (ash). Observing flame colour and amount of smoke helps assess completeness. Carbon monoxide cannot be seen or smelled, so good ventilation and detectors are essential where combustion occurs indoors. Practically, a blue flame on a gas stove shows efficient burning; a smoky yellow flame indicates incomplete combustion and poor efficiency.

Uses and hazards
Combustion powers vehicles, generates electricity in power stations and is used in industry for heating processes. Hazards include fire risk, burns and emission of pollutants. Carbon monoxide poisoning is dangerous because CO binds to haemoglobin and prevents oxygen transport. Use proper ventilation, maintain appliances, keep flammable materials away from flames, and have fire extinguishers ready.

Environmental considerations
Combustion of fossil fuels emits greenhouse gases (mainly CO2) contributing to climate change. Reducing unnecessary burning, improving efficiency and switching to cleaner fuels lower environmental impact. Learning about combustion helps students understand energy use, safety and environmental protection.

The law stated simply
The law of conservation of mass says that mass is neither created nor destroyed in a chemical reaction. For class 7, this means that if a reaction is carried out in a closed container so that nothing can escape, the total mass before and after the reaction remains the same. This idea helps explain why matter seems to disappear sometimes — often gases are released and escape from the system.

Common classroom demonstrations
A popular experiment mixes baking soda and vinegar in a sealed plastic bag or closed flask. When the reaction occurs, carbon dioxide gas is produced, but because the system is closed, the mass measured before and after stays constant. If the container is open, gas escapes and the measured mass falls, which could be misinterpreted unless one realises matter has left the system as gas.

Careful procedure
To demonstrate conservation, weigh the closed system accurately before mixing. Mix the reactants inside without opening the system, allow the reaction to finish, and weigh again. Use a balance with appropriate precision. For more visible demonstrations, burn a small candle under a sealed jar placed on a balance; the total mass of jar plus candle will remain nearly constant though the candle alone loses mass because combustion products are contained within the closed jar.

Interpretation and errors
Small differences in measurements arise from experimental error: heat loss, sticking of residues to apparatus, or balance inaccuracies. Emphasise that conservation applies to the closed system; if products escape (as gases or splashes), measured mass of the remaining items will change but total mass including escaped materials still obeys the law.

Importance for chemistry
This law underpins the balancing of chemical equations taught later. It ensures that atoms and mass are accounted for in reactions and helps chemists design processes that conserve materials and reduce waste. Understanding the law reinforces careful experimental technique and clear thinking about systems and boundaries.

Basic definitions
Acids are substances that release hydrogen ions (H+) in water and usually taste sour (do not taste chemicals in the lab). Bases release hydroxide ions (OH-) in water and feel slippery and taste bitter (again, do not taste in the lab). These chemical properties cause acids and bases to react differently with metals, carbonates and other materials.

Indicators and testing
Indicators are substances that change colour depending on the acidity or alkalinity of a solution. Litmus paper is a common classroom indicator: it turns red in acidic solutions and blue in basic solutions. Universal indicator provides a range of colours across the pH scale, from red for strongly acidic to purple for strongly alkaline. Natural indicators, such as red cabbage extract or turmeric, also show colour changes and provide simple, safe tests.

Chemical changes involving acids and bases
Acids react with metals to produce hydrogen gas and a salt — for example, dilute hydrochloric acid reacts with zinc to give zinc chloride and hydrogen gas. Acids react with carbonates to produce carbon dioxide gas, water and a salt; bubbling gas and limewater tests can confirm CO2 formation. Neutralisation is a chemical reaction where an acid and a base form a salt and water; this reaction releases or absorbs energy and changes the chemical identity of the reactants.

Everyday importance and safety
Acids and bases are common at home: lemon juice and vinegar are acids; baking soda is a base. They are used in cleaning, food preparation and medicine. Handle acids and bases with care: use small quantities, wear eye protection when needed, and neutralise spills before disposal when instructed. Indicators help test water quality and soil acidity in gardening and agriculture.

Practical classroom activities
Simple experiments include testing household liquids with litmus or universal indicator, observing reactions of acids with carbonates (carefully), and demonstrating neutralisation with antacid tablets. Record colour changes and gas evolution, and relate observations to the chemical changes taking place. These activities build understanding of how chemical tests reveal underlying reactions.

Everyday examples of changes
Physical and chemical changes are all around us. Cooking food involves chemical changes: proteins denature, starches gelatinise and new flavours form. Freezing, drying and refrigeration are physical changes used to preserve food. Cleaning with soap and detergents involves chemical action that removes dirt and oils. Rust prevention, painting, and metal coating are practical efforts to control chemical changes that damage materials.

Industrial and environmental applications
Chemical changes are used to make medicines, plastics, fertilizers and fuels. Physical separation techniques like filtration, evaporation and distillation purify products and recover resources. Understanding these changes helps industries reduce waste, recycle materials and limit pollution. For example, wastewater treatment uses settling, filtration and chemical treatment to remove contaminants.

Safety rules for school and home
Follow simple safety rules: always work under supervision for experiments, wear goggles and aprons when required, do not taste chemicals, and wash hands after handling substances. Use small quantities, label containers clearly, and keep flammable substances away from open flames. Know the location of safety equipment like fire extinguishers, first aid kits and eyewash stations, and learn emergency procedures.

Responsible disposal and environmental care
Dispose chemical wastes as instructed; do not pour harmful chemicals down the drain. Recycle materials like paper, glass and metals to reduce waste. Prevent pollution by cleaning spills immediately and using safe alternatives when possible. Teaching and practising safe disposal helps protect water, soil and living organisms.

Connecting classroom to life
Relate experiments to household tasks: filtration to make tea, evaporation to dry clothes, neutralisation to treat acidity. Encourage thoughtful questions: How can I make a process safer? Can this material be reused? Understanding changes and safety builds good habits, helps make informed choices and prepares students for higher study and everyday decision making.