Metals: Metallic Bonding and Properties of Metals
FoundationA metallic bond is the attraction between metal cations and a sea of delocalised electrons. Free electrons make metals conduct electricity and heat; sliding layers give ductility and malleability; alloys, with atoms of different sizes, resist layer-slipping and are harder than pure metals.
- Metallic bond = electrostatic attraction between metal cations and a sea of delocalised electrons
- Mobile delocalised electrons make metals conduct both electricity and heat
- Slipping cation layers without bond-breaking give metals ductility and malleability
文體小知識：科學知識說明文——以離域電子海模型解釋金屬導電、導熱、有延展性和熔點高，並說明合金為何比純金屬硬。
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1The outermost electrons of metal atoms detach easily. Once detached, they belong to no single atom but flow freely through the whole piece of metal — the delocalised electrons. The remaining atoms become metal cations, packed in orderly layers. The electrostatic attraction between these cations and the delocalised electrons is the metallic bond. This 'sea of electrons' model explains the conductivity, heat conduction, malleability and high melting points of metals at one stroke.
2Connect a metal to a power supply and the delocalised electrons drift one way, forming a current — metals conduct electricity. Heat also travels swiftly via the free electrons and the vibrating cations knocking one another, so metals conduct heat well. Copper is made into wires and rice-cooker bowls, aluminium into cookware and heat sinks; nearly every metal object around us relies on these two properties.
3When a force is applied, layers of cations can slip past one another, delocalised electrons immediately filling the new positions, and the metallic bond does not break. Metals can therefore be rolled into sheets or drawn into wires without shattering — malleability and ductility. Metallic bonds are generally strong, so metals melt at high temperatures: tungsten melts at about 3410°C and serves as lamp filaments, while mercury, with weaker bonds, is the only metal that is liquid at room temperature.
4Mixing atoms of a different size into a metal produces an alloy. Because the atoms differ in size, the cation layers can no longer slip neatly, so alloys are usually harder than pure metals. Steel, an alloy of iron and carbon, is hard and used in buildings and bridges; brass, of copper and zinc, goes into instruments and locks; stainless steel adds chromium and nickel for rust resistance. Most metal materials of daily life are alloys, not pure metals.
1Summary 重點摘要
A metallic bond is the attraction between metal cations and a sea of delocalised electrons. Free electrons make metals conduct electricity and heat; sliding layers give ductility and malleability; alloys, with atoms of different sizes, resist layer-slipping and are harder than pure metals.
2Vocabulary cards Vocabulary
The electrostatic attraction between metal cations and delocalised electrons.
Outer-shell electrons that have left their atoms and flow freely through the metal; also called free electrons.
Metal cations in layered lattices immersed in a sea of delocalised electrons.
The ability to be drawn into wires without breaking, e.g. copper wire.
The ability to be rolled into sheets without shattering, e.g. aluminium foil.
The ability to carry a current; metals conduct via moving delocalised electrons.
The ability to transfer heat; metals conduct via free electrons and vibrating cations.
A mixture of a metal with other metals or non-metals; usually harder than the pure metal.
An alloy of copper and zinc, harder than pure copper, used in instruments and locks.
Steel with chromium and nickel added, highly rust-resistant, used in cutlery and medical tools.
3Step by step Steps
- Step 1: The electron-sea model of the metallic bondThe outer-shell electrons of metal atoms are loosely held and detach easily. Take copper: each atom gives up its outer electrons and becomes a Cu cation; the cations line up in layers, while the released electrons belong to no single ion and roam the whole metal — the electron sea. The attraction between cations and sea is the metallic bond, binding the metal into one piece. In answers, first state that metal cations sit in a sea of delocalised electrons, then add the mutual attraction, and the model is complete.
- Step 2: Conduction — electrons on the moveWithout a supply, delocalised electrons jostle in all directions; once connected to a source, they drift together in one direction — a current. For example, copper wire fills our buildings because copper has many easily moving delocalised electrons, conducts almost as well as silver but far more cheaply. Compare solid sodium chloride: its ions are locked in the lattice and cannot move, so solid salt does not conduct. 'Mobile charges present' is the common prerequisite for conduction.
- Step 3: Heat conduction — a relay of energyMetals conduct heat through two relays: free electrons, which speed up on heating and carry energy swiftly around the metal, and cations, whose stronger vibrations pass energy by collision with neighbours. For example, an aluminium pot on the stove heats its handle quickly through this relay; a plastic handle, lacking free electrons, conducts slowly and insulates instead. Heat conduction and electrical conduction share the same origin — delocalised electrons — and often appear together in exam questions.
- Step 4: Malleability — layers that can slipMetal cations sit in neat layers. Under a hammer blow or a pull, one layer can slide to a new position, delocalised electrons immediately surround the new arrangement, and the metallic bond reconnects — the structure does not fracture. For example, gold is beaten into gilding foil and aluminium rolled into kitchen foil, both thanks to layer-slipping. Contrast an ionic crystal: a slipped layer brings like charges face to face, mutual repulsion shatters the crystal — salt grains crack under a blow while iron can be rolled into plates.
- Step 5: High melting points — tungsten versus mercuryThe metallic bond rests on strong attraction between cations and the electron sea; melting must supply energy for ions to break away, so metals generally melt high. Tungsten melts at about 3410°C, glowing white-hot without melting, hence lamp filaments are tungsten. Mercury's metallic bond is weaker: it melts at about −39°C and is liquid at room temperature — the only liquid metal, once used in thermometers. The stronger the metallic bond, the higher the melting point — a rule often tested in comparison questions.
- Step 6: Alloys — layers that cannot slipMixing atoms of a different size into a pure metal makes an alloy. For example, brass is copper with zinc; zinc atoms are slightly larger and sit between the copper layers, making the planes uneven, so they can no longer slip neatly under a force — the alloy is harder. Steel is iron with small carbon atoms wedged between iron layers, with the same effect; stainless steel adds chromium and nickel for hardness plus rust resistance. Alloys are mixtures whose proportions can be tuned, designing materials of chosen hardness and durability — far more useful industrially than pure metals.
- Step 7: Choosing materials from propertiesEvery use of a metal maps onto its structure and properties: for conduction choose copper or aluminium cables, rich in delocalised electrons and affordable; for heat resistance choose tungsten filaments, with their extreme melting point; for hardness and looks choose brass locks and instruments; for rust resistance choose stainless-steel cutlery. For example, overhead power lines use aluminium — light and conductive enough — to spare the pylons. In data questions, argue structure (electron sea, layers), then properties (conduction, malleability, high melting point, hardness), then the use — a complete chain.
4Diagrams Diagrams
5Worked example Example
Explanation question: a workshop chooses copper for cables, tungsten for filaments and brass for locks. Using metallic bonding and the giant metallic structure, explain: (1) why copper conducts electricity; (2) why metals are generally ductile and malleable; (3) why tungsten suits lamp filaments; (4) why brass is harder than pure copper.
- Describe the model: copper consists of layers of Cu cations and delocalised electrons; the cations sit in an electron sea with mutual attraction — the metallic bond
- Conduction: when connected, delocalised electrons drift in one direction, forming a current — copper conducts
- Ductility and malleability: under force the cation layers slip, delocalised electrons fill the new positions, bonds do not break — sheets and wires form
- Tungsten filament: strong metallic bond gives a melting point of about 3410°C, surviving white-hot temperatures
- Brass harder: zinc and copper atoms differ in size, wedged between layers so the planes cannot slip neatly — harder than pure copper
- Answer: Copper consists of Cu cations and delocalised electrons: when a voltage is applied, the delocalised electrons drift in one direction and form a current, so copper conducts electricity. When a force acts on a metal, the cation layers can slip while delocalised electrons immediately fill the new positions, so the metallic bond does not break — metals are generally malleable and ductile. Tungsten's metallic bond is very strong, giving a melting point of about 3410°C, so a filament survives white heat. Brass is a copper-zinc alloy: the two atom sizes make the layers uneven and unable to slip neatly, so brass is harder than pure copper.
6Common mistakes Mistakes
7Quick check Quick check
Q1What is the nature of the metallic bond?
Q2Why do metals conduct electricity while solid salt does not?
Q3Why can metals be drawn into wires while salt crystals shatter under a blow?
Q4Which metals make up brass, and why is it harder than pure copper?
Q5Why are lamp filaments made of tungsten?
8Recap Recap
- Metallic bond = electrostatic attraction between metal cations and a sea of delocalised electrons
- Mobile delocalised electrons make metals conduct both electricity and heat
- Slipping cation layers without bond-breaking give metals ductility and malleability
- Strong metallic bonds mean high melting points: tungsten (about 3410°C) for filaments; mercury is the only liquid metal
- Alloys are mixtures: mismatched atom sizes block layer-slipping, so they are harder than pure metals (steel, brass, stainless steel)