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METAL BONDING

CHEMICAL BONDING Part 5 Metallic Bonding, Structure and Properties of Metals

Doc Brown's Science–Chemistry Chemical Bonding GCSE/IGCSE/O Level/AS/A2 Level Revision Notes

DIAGRAMS of METAL STRUCTURES and their PROPERTIES EXPLAINED – Metallic bonding is described and the properties of pure metals and alloys are described and explained using the giant metal lattice structure model which is used to explain the physical properties of metals. The structure of alloys is explained and why alloy metals are more useful than pure metals. These notes on bonding in metals and explaining the structure and properties of metallic structures are designed to meet the highest standards of knowledge and understanding required for students/pupils doing GCSE chemistry, IGCSE chemistry, O Level chemistry, KS4 science courses and a basic primer for AS/A Level chemistry courses.

Part 1 Introduction – why do atoms bond together? (I suggest you read 1st)

Part 2 Ionic Bonding – compounds and properties

Part 3 Covalent Bonding – small simple molecules and properties

Part 4 Covalent Bonding – macromolecules and giant covalent structures

Part 5 Metallic Bonding – structure and properties of metals (this page)

Part 6 More advanced concepts for advanced level chemistry (in preparation, BUT a lot on intermolecular forces (intermolecular bonding) in Equilibria Part 8)

See also How can metals be made more useful? (alloys of Al, Fe, steel etc.)

and The physical and chemical properties of transition metals


topPart 5.  METALLIC BONDING – structure and properties of metals

metal bonding model of element & alloys * physical properties of metals

Its a good idea to have some idea of where the metallic elements are in the periodic table

The black zig–zag line 'roughly' divides the metals on the left from the non–metals on the right of the elements of the Periodic Table.

Pd metals Part of the modern Periodic Table

Pd = period, Gp = group

metals => non–metals
Gp1 Gp2 Gp3 Gp4 Gp5 Gp6 Gp7 Gp0
1

1H  Note that H does not readily fit into any group

2He
2 3Li 4Be atomic number Chemical Symbol eg 4Be 5B 6C 7N 8O 9F 10Ne
3 11Na 12Mg 13Al 14Si 15P 16S 17Cl 18Ar
4 19K 20Ca 21Sc 22Ti 23V 24Cr 25Mn 26Fe 27Co 28Ni 29Cu 30Zn 31Ga 32Ge 33As 34Se 35Br 36Kr
5 37Rb 38Sr 39Y 40Zr 41Nb 42Mo 43Tc 44Ru 45Rh 46Pd 47Ag 48Cd 49In 50Sn 51Sb 52Te 53I 54Xe
6 55Cs 56Ba Transition Metals 81Tl 82Pb 83Bi 84Po 85At 86Rn
Gp 1 Alkali Metals  Gp 2 Alkaline Earth Metals  Gp 7 Halogens  Gp 0 Noble Gases

Chemical bonding comments about the selected elements highlighted in white

e.g. the 'white' highlighted elements are typical metals you are likely to have come across, either as a pure metal or in an alloy mixture of metals – all the atoms are held together by what is called 'metallic bonding' – details of the bonding model below

 


(c) doc b BONDING IN METALS

  • (c) doc bMETAL STRUCTURE
  • All metals have similar properties BUT, there can be wide variations in melting point, boiling point, density, electrical conductivity and physical strength.
  • To explain the physical properties of metals like iron or sodium we need a more sophisticated picture than a simple particle model of atoms all lined up in close packed rows and layers, though this picture is correctly described as another example of a giant lattice held together by metallic bonding.
  • A giant metallic lattice – the crystal lattice of metals consists of ions (NOT atoms) surrounded by a 'sea of electrons' that form the giant lattice.
  • The outer electrons (–) from the original metal atoms are free to move around between the positive metal ions formed (+).
  • These 'free' or 'delocalised' electrons from the outer shell of the metal atoms are the 'electronic glue' holding the particles together.
  • There is a strong electrical force of attraction between these free electrons (mobile electrons or 'sea' of delocalised electrons) (–) and the 'immobile' positive metal ions (+) that form the giant lattice and this is the metallic bond.
  • Metallic bonding is not directional like covalent bonding, it is like ionic bonding in the sense that the force of attraction between the positive metal ions and the mobile electrons acts in every direction about the fixed (immobile) metal ions of the metal crystal lattice, but in ionic lattices none of the ions are mobile. a big difference between a metal bond and an ionic bond.
  • Metals can become weakened when repeatedly stressed and strained.

    • This can lead to faults developing in the metal structure called 'metal fatigue' or 'stress fractures'.

    • If the metal fatigue is significant it can lead to the collapse of a metal structure.

    • So it is important develop alloys which are well designed, well tested and will last the expected lifetime of the structure whether it be part of an aircraft (eg titanium aircraft frame) or a part of a bridge (eg steel suspension cables).

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 Explaining the physical properties of metals

  • 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 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, the more energy is needed to weaken the force between them sufficiently to break the giant lattice down in melting and completely 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' 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 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 upto a point!)

      • The layers of atoms can slide over each other without fracturing the structure (see below).

      • 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.

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

  • Potential problems with metal structures?

    • Although the metals used in construction are strong, in some situations they may become dangerously weak e.g.

      • If iron or steel becomes badly corroded, there is no strength in rust!, and, the thicker the rust layer, the thinner and weaker the supporting iron or steel layer, hence the possibility of structural failure. Therefore, most iron and steel structures exposed to the outside weather are maintained with a good coating of paint.

      • Also, if metal structures e.g. in aircraft or bridges, are continually strained under stress, the crystal structure of the metal can change so it becomes brittle. This effect is called metal fatigue (stress fractures) and may lead to a very dangerous situation of mechanical failure of the structure.

      • So it is important develop alloys which are well designed, well tested and will last the expected lifetime of the structure whether it be part of an aircraft (eg titanium aircraft frame) or a part of a bridge (eg steel suspension cables).

      • See notes on Corrosion of Metals and Rust Prevention

 Note on Alloy Structure  via a very simplified diagram

(c) doc b

An alloy is a mixture of a metal with other elements (metals or non-metals). Metals can be mixed together to make alloys to improve the metal's properties to better suit a particular purpose. An alloy mixture often has superior desired properties compared to the pure metal or metals i.e. the alloy has its own unique properties and a more useful metal. 

  1. Shows the regular arrangement of the atoms in a pure metal crystal and the white spaces show where the free electrons are (yellow circles actually positive metal ions).
  2. Shows what happens when the metal is stressed by a strong force. The layers of atoms can slide over each other and the bonding is maintained as the mobile electrons keep in contact with ions of the giant lattice, so the metal object remains intact BUT the metal is physically a different shape.
  3. Shows an alloy mixture. Alloys are NOT compounds but a physical mixing of a metal plus at least one other material (shown by red circle), it can be another metal e.g. nickel or manganese added to iron in steel, or a non–metal e.g. carbon, and it can be bigger or smaller than the iron atoms. Many alloys are produced like this to give a stronger metal. The presence of the other atoms (smaller or bigger) disrupts the symmetry of the layers and reduces the 'slip ability' of one layer next to another. The result is a stronger harder less malleable metal.
  4. The main point about using alloys is that you can make up, and try out, all sorts of different compositions until you find the one that best suits the required purpose in terms of tensile/compression strength, malleability, electrical conductivity or corrosion resistance etc.
    • The are hundreds of alloys of steel made by alloying iron with other metals to increase the strength or anti-corrosion properties of the metal.
      • Steel is used in building and bridge construction, car bodies, railway lines and countless other objects that need to have a high tensile strength.
    • Pure metals can be either too soft (e.g. like copper or tin) or too brittle (e.g. like zinc) to be used directly and are therefore often alloyed to make superior metals like brass or bronze.
    • The properties of metals are readily matched to a particular use e.g.
      • Aluminium alloys are strong and light (relatively low density for a metal), they do not corrode easily and so are used in aircraft construction, greenhouse frames and not as expensive as titanium alloys.
      • Cooking pans made of stainless steel are good conductors of heat, strong with good anti-corrosion properties and steel has a high melting point.
      • Copper is malleable and ductile, easily drawn out into wire, an excellent conductor of electricity, and so is widely used in electrical circuitry.
    • Steel alloys of varying strength and anti-corrosion properties are used in thousands of products and constructions e.g. reinforcing rods in concrete buildings, bridge girders, car engines, domestic appliances from washing machines to electric kettles, saucepans, tools like chisels, ship hulls and superstructure, very hard drill bits,

    • For more on specific metals and alloys see my notes on Transition Metals

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Revision notes information to help revise KS4 Science Additional Science Triple Award Separate Sciences GCSE/IGCSE/O level Chemistry Revision–Information Study Notes for revising for AQA GCSE Science, Edexcel GCSE Science/IGCSE Chemistry & OCR 21st Century Science, OCR Gateway Science WJEC/CBAC GCSE science–chemistry CCEA/CEA GCSE science–chemistry (and courses equal to US grades 8, 9, 10) basic aid notes for GCE Advanced Subsidiary Level AS Advanced Level A2 IB Revise AQA OCR Edexcel Salters CIE, CCEA/CEA & WJEC advanced level courses for pre–university students (equal to US grade 11 and grade 12 and Honours/honors level courses)


WHAT NEXT? – other pages to do with metals

 



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