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5c. Explaining the properties of metals using the metallic bonding model

Doc Brown's Chemistry: Chemical Bonding and structure GCSE level, IGCSE, O, IB, AS, A level US grade 9-12 level Revision Notes

 Explaining the physical properties of metals

(c) doc b The giant metallic lattice of metal ions

  • All metals are lustrous and, compared to non-metals, most metals are quite dense, hard (tough, high tensile strength), with high melting/boiling points, though there notable exceptions e.g.

    • mercury is a liquid at room temperature, group 1 alkali metals like sodium and potassium are less dense than water ('float') and have low melting points <100oC).

  • The strong metallic bonding generally results in dense, strong materials with high melting and boiling points.

    • Usually a relatively large amount of energy is needed to melt or boil metals.

    • The stronger the attraction between the atoms/ions in the giant metallic lattice, more kinetic energy of the particles (metal atoms) is needed to weaken the force between them sufficiently to break the giant lattice down in melting and eventually sufficiently great enough to overcome the attractive forces to boil the metal.

    • Energy changes for the physical changes of state of melting and boiling for a range of differently bonded substances are compared in a section of the Energetics Notes.

    • The strong bonding in metals gives them a high tensile strength, so alloys like steel are used in building construction, car bodies etc.

  • Metals are good conductors of electricity

    • Why are metals good conductors of electricity?

    • Metals are good at conducting electricity because these 'free' delocalised electrons carry the charge of an electric current when a potential difference (voltage) is applied across a piece of metal

      • e.g. copper wire is excellent to use as an electrical conductor in household wiring or any electrical appliances.

  • Metals are also good conductors of heat.

    • Why are metals good conductors of heat?

    • The fact that metals are good at conducting heat is also due to the free moving electrons.

    • Non–metallic solids conduct heat energy by hotter more strongly vibrating atoms, knocking against cooler less strongly vibrating atoms to pass the particle kinetic energy on.

    • BUT in metals, as well as this effect, the 'hot' high kinetic energy electrons move around freely to transfer the particle kinetic energy more efficiently to 'cooler' atoms.

    • This is a faster process than the transferring heat by the kinetic energy of atom vibration.

      • So, where a material needs to be a good heat conductor, metals quite naturally are used to make everything from radiators, cooking pans etc.

      • Its also hand that they are both strong and high melting when used as a saucepan!

  • Typical metals also have a silvery surface (lustrous) but remember this may be easily tarnished by corrosive oxidation in air and water.

    • Although many metals will corrode (oxidise) in the presence of air (oxygen) and water, the strong bonding prevents them dissolving in water or any other laboratory solvent. When metals like sodium 'dissolve in water, they do so via a chemical reaction forming a soluble compound (sodium hydroxide), and do NOT give a solution of sodium metal.

  • Unlike ionic and non-metallic element solids, metals are very malleable - easy to bend or hammer into shape

    • Why are metals very malleable and easily bent or pressed shaped?

    • Metals are be readily bent, pressed or hammered into shape because the strong bonding is retained even when the metal is stressed (at least up to a point!).

      • The layers of atoms can slide over each other without fracturing the structure.

      • The reason for this is the mobility of the electrons involved in the metallic bonding.

      • When planes of metal atoms are 'bent' or slide the electrons can run in between the atoms and maintain a strong bonding situation. This can't happen in ionic solids which tend to be brittle and the layers of immobile ions fracture easily.

      • Unfortunately, sometimes a pure metal is too malleable i.e. to weak for a given purpose, but this problem can be overcome by alloying the metal with other elements, which are usually metals too. The resulting alloy can be stronger and tailored to suite a particular application.

  • Metals usually have a high density

    • The strong bonding, particularly when several electrons are delocalised per atom, draws the metal ions in the metallic lattice close together, increasing the mass per unit volume i.e. increasing the density compared to most non-metallic elements.

      • e.g. densities in g/cm3 (multiply by 1000 to convert to kg/m3)

      • non-metals: sulfur 2.0;  bromine 3.1;  carbon 2.25 (graphite) and 3.51 (diamond)

      • metals: aluminium 2.7 (unusually low);  iron 7.9;  gold 19.3

  • For more on the properties and uses of metals see Transition Metals and Extra Industrial Chemistry pages and the note and diagram below.






What next?

Recommend next: Alloys - improved design, problems with fatigue and corrosion


Sub-index: Part 5 Metallic Bonding – structure and properties of metals


Index for ALL chemical bonding and structure notes


Perhaps of interest?

How can metals be made more useful? (useful links)


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