What this chapter is about
This chapter explores the two broad categories into which most elements are classified: metals and non-metals. A Class 10 student learns to distinguish these categories based on their physical and chemical properties, and understands why they behave differently in reactions.
The chapter covers how metals react with oxygen, water, acids and other metal salts. It introduces the reactivity series, which ranks metals by how readily they lose electrons and react. Students also learn about extraction of metals from their ores, including methods like roasting, calcination, reduction and electrolytic refining.
Finally, the chapter explains corrosion (why iron rusts, why silver tarnishes) and how to prevent it. By the end, students should be able to predict reaction outcomes, write balanced equations for metal reactions, and explain everyday phenomena like why copper vessels turn green or why aluminium is used for food packaging.
Key ideas
- Metals are generally lustrous, malleable, ductile, good conductors of heat and electricity, and have high melting points. Non-metals are generally brittle, dull, poor conductors, and have low melting points.
- Metals form positive ions (cations) by losing electrons; non-metals form negative ions (anions) by gaining electrons.
- The reactivity series arranges metals in order of decreasing reactivity: potassium, sodium, calcium, magnesium, aluminium, zinc, iron, lead, copper, mercury, silver, gold.
- Ionic compounds are formed when metals transfer electrons to non-metals. These compounds have high melting points, conduct electricity when molten or dissolved, and are usually soluble in water.
- Extraction method depends on reactivity: highly reactive metals (sodium, aluminium) need electrolysis; moderately reactive metals (zinc, iron) are reduced using carbon; least reactive metals (copper, silver) occur free or need simple heating.
- Corrosion is the slow chemical attack on metal surfaces by air, water or chemicals. Rusting of iron requires both oxygen and moisture.
- Alloys are homogeneous mixtures of metals (or metal with non-metal) made to improve properties. Examples: brass (copper + zinc), bronze (copper + tin), stainless steel (iron + chromium + nickel + carbon).
Formulas and facts to remember
- Metal + Oxygen → Metal oxide (basic oxide). Example: 2Mg + O₂ → 2MgO
- Metal + Water → Metal hydroxide + Hydrogen. Example: 2Na + 2H₂O → 2NaOH + H₂
- Metal + Dilute acid → Salt + Hydrogen. Example: Zn + 2HCl → ZnCl₂ + H₂
- More reactive metal displaces less reactive metal from its salt solution. Example: Fe + CuSO₄ → FeSO₄ + Cu
- Metal oxide + Carbon → Metal + Carbon monoxide (reduction). Example: ZnO + C → Zn + CO
- Copper, silver, gold and platinum are least reactive and do not react with dilute acids.
- Amphoteric oxides react with both acids and bases. Examples: aluminium oxide, zinc oxide.
- Rusting of iron: 4Fe + 3O₂ + 2xH₂O → 2Fe₂O₃·xH₂O (hydrated iron(III) oxide)
Worked examples
Example 1: Predicting a displacement reaction
A student dips an iron nail into copper sulphate solution. Will a reaction occur? Write the equation.
Solution: Step 1: Check the reactivity series. Iron is above copper, so iron is more reactive. Step 2: A more reactive metal displaces a less reactive metal from its solution. Step 3: Reaction occurs. Iron displaces copper. Balanced equation: Fe + CuSO₄ → FeSO₄ + Cu Observation: The blue colour of copper sulphate fades, and a reddish-brown deposit of copper forms on the nail.
Example 2: Choosing an extraction method
Why is aluminium extracted by electrolysis while iron is extracted by reduction with carbon?
Solution: Step 1: Aluminium is very high in the reactivity series (above carbon). Step 2: Carbon cannot reduce aluminium oxide because aluminium holds oxygen more strongly than carbon does. Step 3: Electrolysis supplies enough energy to break the strong Al–O bonds. Step 4: Iron is below carbon in reactivity, so carbon can reduce iron oxide. Therefore, iron ore is heated with coke (carbon): Fe₂O₃ + 3C → 2Fe + 3CO.
Example 3: Preventing corrosion
A bicycle manufacturer wants to prevent the iron frame from rusting. Suggest two methods and explain how each works.
Solution: Method 1: Painting or oiling the frame. This creates a barrier that keeps oxygen and moisture away from iron. Method 2: Galvanisation (coating with zinc). Zinc is more reactive than iron, so even if the coating is scratched, zinc corrodes first, protecting the iron beneath. This is called sacrificial protection.
Common mistakes
Thinking all metals react with dilute acids → Copper, silver and gold are below hydrogen in reactivity and do not displace hydrogen from acids.
Writing hydrogen as H instead of H₂ in equations → Hydrogen exists as diatomic molecules; always write H₂.
Believing aluminium does not corrode because it looks shiny → Aluminium corrodes quickly, but its oxide layer is protective and sticks to the surface, preventing further damage.
Confusing calcination and roasting → Calcination heats ore in absence of air (for carbonates); roasting heats ore in presence of air (for sulphides).
Assuming alloys are compounds → Alloys are mixtures, not compounds; their composition can vary.
Quick revision
- Metals lose electrons to form cations; non-metals gain electrons to form anions.
- Reactivity series: K, Na, Ca, Mg, Al, Zn, Fe, Pb, H, Cu, Hg, Ag, Au (decreasing order).
- More reactive metal displaces less reactive metal from its salt solution.
- Highly reactive metals: electrolysis; moderately reactive: carbon reduction; least reactive: simple heating or occur free.
- Rusting needs both oxygen and water; prevent by coating, galvanising or alloying.
- Ionic compounds have high melting points and conduct electricity only when molten or in solution.