Metals and Non-metals: Unit Review

Introduction

This unit links what metals are like to what they do. Conductivity and malleability arise from metallic structure, while reaction patterns reflect electron transfer and the reactivity series. Metal–acid and displacement equations then lead naturally to extraction and corrosion. A strong review answer chooses the correct model for the question and states the conditions under which its prediction applies.

Core explanation

Many metals conduct heat and electricity, can be shaped and have a shiny freshly cut surface. Metallic bonding provides a useful model: positive metal ions or atomic cores occupy a lattice while delocalised electrons move through it. Their movement supports electrical conduction, and the non-directional attraction allows layers to shift without the immediate bond failure expected from a simple rigid model. These are tendencies, not rules without exceptions. Non-metals more commonly lack the same metallic conductivity and malleability, but graphite conducts and some elements sit near the metal/non-metal boundary.

Metals often lose electrons to form positive ions, while many non-metals gain electrons to form anions in ionic compounds or share electrons in covalent compounds. For example, magnesium and oxygen form MgO with Mg²⁺ and O²⁻ in a lattice; carbon and oxygen can form molecular CO₂ with shared electrons. Bonding type affects melting behaviour, solubility and conduction, but properties should be explained from the particular structure rather than from a label alone.

The reactivity series orders common metals for school-level comparisons. A useful segment runs Mg, Al, Zn, Fe, Pb, H, Cu, Ag, Au from higher to lower metal reactivity, with hydrogen used as a reference marker. Metals suitably above H can often release H₂ from dilute non-oxidising acids: Zn + 2HCl → ZnCl₂ + H₂. Copper below H does not ordinarily produce H₂ from dilute HCl. This is not universal acid resistance; oxidising acids can take different pathways. Rate also differs from direction: a protective oxide can delay a reaction even when the metal is high in the series.

In a salt-solution displacement, compare the solid metal with the metal ion in solution. Zinc above copper gives Zn + CuSO₄ → ZnSO₄ + Cu; copper below zinc does not give the reverse simple reaction. The net ionic equation Zn + Cu²⁺ → Zn²⁺ + Cu shows the electron transfer. Zinc loses two electrons and is oxidised, making it the reducing agent. Cu²⁺ gains them and is reduced, making it the oxidising agent. A balanced equation must conserve atoms and charge, but balance by itself does not prove that the direction occurs.

Reactivity also helps explain occurrence and extraction. Gold can persist native, while more reactive metals are usually found in compounds. Suitable less reactive metal oxides can be reduced by carbon-based methods, whereas aluminium needs electrolysis of its compounds in an industrial molten system. This broad pattern does not replace knowledge of the actual ore, energy supply and processing stages. Alloying is another way to change metal performance: steel includes iron and carbon, brass is mainly copper with zinc, and traditional bronze is copper with tin. Added atoms and processing can hinder deformation and alter hardness.

Corrosion is environmental metal degradation; rusting is the narrower iron case. Ordinary iron rusting needs both water and oxygen and can produce hydrated iron(III)-oxide-rich material, often represented Fe₂O₃·xH₂O. Salt can accelerate it but is not required. Paint, oil, grease and plastic limit reactant access while intact. Galvanising adds zinc: it is a barrier and can oxidise sacrificially near a small wet scratch when connected to iron. Aluminium's compact oxide often protects; copper patina and silver tarnish have different compositions from iron rust. These comparisons show why a series ranking alone cannot predict everyday durability.

Step-by-step reasoning

1. Identify whether the question concerns structure, acid reaction, salt displacement, extraction or corrosion. 2. For reactivity, compare the correct partners and specify the acid or salt conditions. 3. Write valid product formulas and balance atoms; for ionic equations, also balance charge. 4. Interpret observations through electron transfer or product-layer behaviour, and note limits of a simplified prediction.

Visual explanation

Draw a centre box labelled “metal atoms can lose electrons.” Branch to metallic bonding and properties, to the reactivity ladder and displacement, to ore extraction, and to corrosion. Put a separate branch for non-metal electron gain or sharing so structure and reaction patterns remain connected but distinct.

Real-world analogy

A map has roads, distances and landmarks; one symbol cannot answer every travel question. This unit's models likewise answer different questions. Metallic bonding explains many physical properties, while the series predicts relative reaction directions and surface chemistry explains corrosion resistance.

Real-world example

An outdoor steel fence uses an iron-based alloy for strength but needs protection from water and oxygen. A zinc coating can offer both a barrier and sacrificial action. The same series that explains zinc's ability to displace copper also helps explain why zinc is consumed before connected iron in a wet protective couple.

Why?

Why are properties and reactions taught together? Material choice depends on both. A conductor must carry current, yet its surface must survive the intended environment. Structure explains the first demand; redox and corrosion chemistry explain much of the second.

Common misconception

“A metal high in the series must corrode visibly faster outdoors.” Aluminium is highly reactive but forms an adherent oxide that often slows further attack. Iron forms rust that can expose fresh metal. The reactivity series and protective-layer behaviour answer different parts of the problem.

Worked example

An iron nail is placed in aqueous CuSO₄, and a second iron nail is kept in wet air. Predict both changes. Iron lies above Cu, so Fe + CuSO₄ → FeSO₄ + Cu can deposit copper on the first nail; Fe is oxidised to Fe²⁺ and Cu²⁺ reduced. The second nail can rust because water and oxygen are present, making hydrated iron(III)-oxide-rich products over time. The two processes involve iron oxidation, but the electron acceptors and products differ.

Quick check

1. Which two conditions are required for ordinary iron rusting, and which metal can displace copper from CuSO₄: zinc or silver? Answer: Water and oxygen are required for rusting; zinc can displace copper because zinc is above copper.

Exam focus

State the comparison that justifies each prediction. For a metal–acid question, identify the hydrogen reference and acid conditions. For displacement, name the solid metal and aqueous ion. For corrosion, distinguish rust from other metal surface products.

Advanced insight

The reactivity series is related to electrode potentials under specified conditions, whereas observed speed also depends on activation barriers, transport and passivating layers. Material engineering adds composition, microstructure and local environment. Using these layers of explanation prevents a single qualitative rule from being overextended.

Summary

Metallic structure explains common conductivity and malleability. The reactivity series guides metal–acid and salt-displacement predictions, while redox tracks electrons. Ore extraction, alloys and corrosion show how that chemistry affects real materials. Always state the conditions and limits of the model used.

Practice questions

1. Why does a metal typically conduct electricity? Answer: Mobile delocalised electrons can move through its metallic structure. 2. Predict Cu(s) in dilute non-oxidising HCl as a source of H₂. Answer: No significant H₂ displacement is expected because copper is below hydrogen. 3. Write the net ionic equation for Zn(s) in aqueous CuSO₄. Answer: Zn(s) + Cu²⁺(aq) → Zn²⁺(aq) + Cu(s). 4. Name the two roles of zinc in galvanised steel. Answer: It acts as a barrier and can oxidise sacrificially to protect nearby connected iron in wet conditions. 5. Why can aluminium resist further corrosion after initial surface reaction? Answer: Its thin adherent Al₂O₃ layer can limit further access of environmental reactants.