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Metals. Chapter 5. Metals. Look at the periodic table How many elements are metals? Look around this room How many things are made out of metals or metal alloys? Why do we use metals so much? Can you think of any properties of metals?. Uses of Metals. Why do we use metals?
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Metals Chapter 5
Metals • Look at the periodic table • How many elements are metals? • Look around this room • How many things are made out of metals or metal alloys? • Why do we use metals so much? • Can you think of any properties of metals?
Uses of Metals • Why do we use metals? • The strength of metals is useful when building robust structures. • The lightness and strength of some metals are useful for boats and aircraft • The ability to form wire from metals is applied in many different objects from copper wiring to jewellery • Electricity is transmitted by metals in both the home or industry • Metals can be used to make diverse objects which can be moulded and shaped
Useful Properties of Metals • Elements we classify as metals have all or most of the following properties. They: • Are good conductors of electricity • Are good conductors of heat • Are malleable • Can be shaped by beating or rolling • Are ductile • Can be drawn into a wire • Exhibit a range of melting temperatures and relatively high boiling temperatures • Generally have high densities • Are lustrous or reflective (shiny) • Are often hard, with high tensile strength
Exceptions • Not all metals have all of these properties. • Mercury is a liquid at room temperature meaning it has an unusually low melting temperature. • Chromium is brittle rather than malleable, meaning it breaks rather easily. • Group one metals as you have seen are soft • These elements however exhibit most of the properties and are classified as metals
Your Turn • Page 80 of text. • Questions 2 and 3
Properties and Structure • Each of the properties of metals gives us some information about a metals structure. • Can you make any inferences about the structure of metals.
Clues from the Properties • It is reasonable to assume from those two tables that • The atoms in metals are not uncharged. • As solid metals are particular shapes and do not spontaneously change shape the atoms cannot move around. • They must be able to move however when beaten as metals are malleable and ductile. • They must have strong forces between particles. • The charged particles must be able to move among the atoms and pass on energy. • So how are these particles arranged?
Metals • Most metals have only one or two electrons in their outer shell. • It is relatively easy for another particle to draw the valence electrons away from a metal atom. • This is due to most metals having a low first ionization energy.
Lets consider uncharged sodium1s22s22p63s1 • The valence electron is held very weakly as it is further from the nucleus than the 10 inner shell electrons. • If sodium was by itself the electron would remain in place orbiting sodium. • But if a sodium atom was surrounded on all sides by other sodium atoms, the valence electron, which is on the surface of the atom, will experience some attraction from the positively charged nuclei of the surrounding nuclei.
Sodium1s22s22p63s1 • It makes sense then that the valence electron on each sodium atom present would be drawn in to the space between atoms by the combined attractive forces. • Once in between atoms it will be moving about constantly between all the positively charged sodium atoms so it will no longer remain in the locality of the atom it came from.
The metallic lattice • The electrons that are detached from their atoms are called delocalised electrons. • But what about the sodium atom? • We now have a large set of atoms that have lost an electron. They each have a net charge of 1+, since they still have 11 protons but only 10 electrons. • They are now positively charged sodium ions or sodium cations.
The Metallic Bonding Model • Chemists believe that, in a solid sample of metal: • Positive ions are arranged in a closely packed structure. This structure is described as a three-dimensional lattice of fixed cations. • The smaller negatively charged valence electrons are free to move throughout the lattice. These electrons are called delocalised electrons because they belong to the lattice as a whole. • The inner shell electrons are localised and not free moving.
Electrostatic Forces • The ions are held in the lattice by the electrostatic force of attraction between them and the delocalised electrons. This attraction extends throughout the lattice and is called metallic bonding. • Electrostatic forces are the forces between particles that are caused by their electric charges.
How do the forces affect the structure? • A lot of pushing and pulling in this mixture of positively charged cations and negatively charged delocalised electrons. Why? • This occurs until the sodium ions settle into a stable arrangement. • It is stable when the force of attraction between the sodium ions and delocalised electrons is maximised and the force of repulsion between the sodium ions is minimised.
Lattice? • In solid crystals of sodium there are billions of sodium ions. • For each sodium ion there is how many delocalised electrons? • The positive sodium ions are in a highly organised geometric arrangement with the negatively charged delocalised electrons moving rapidly between the ions. • This structure is termed a metallic lattice. • The delocalised electrons in a metallic lattice are often referred to as a ‘sea of valence electrons’
Explaining the Properties of Metals • We can now explain why metals exhibit some of the properties we have discussed. • As they are highly mobile particles, the delocalised electrons can very rapidly transfer energy from one end of the metal to the other. • Since they are charged they can also transfer electrical energy. • The electrons that are present in the surface of the metal are excellent reflectors of light, so they are responsible for the lustre shown by metals.
Your Turn • Page 82. • Question 5
Limitations of the Metallic Bonding Model • Some things cannot be explained by the metal model. • The range of melting temperatures and densities of different metals. • The differences in electrical conductivity between metals. • The magnetic nature of metals such as cobalt, iron, and nickel. • They model cannot be drawn to scale. • Metal ions are not solid balls, any more than atoms are. Nor are they coloured. Infact ions are pretty much empty space. • Delocalised electrons move throughout the structure extremely rapidly, we cannot show this in a model. • The electrostatic forces cannot be seen.
Your Turn • Page 82 • Question 6
Modifying Metals • Few metals used are in their pure form. • Most metals used need to be modified in order to produce the desired properties for particular uses. • Pure iron is not hard enough, so it is converted to steel by mixing it with 2% carbon. • The properties of metals are altered by adding small amounts of another substance. The substances are melted together, mixed and allowed to cool. • These mixtures are called alloys. • There are two main types of alloys.
Substitutional Alloys • Substitutional alloys are made from elements that have fairly similar chemical properties and atoms of similar size. • Our coins are made from an alloy that is 75% copper and 25% nickel. • The nickel atoms take the place of some copper atoms in the lattice. • Both the nickel and copper ions are attracted to the sea of electrons so the lattice is held strongly in place. • Due to the slight difference in the size of the atoms there is a restriction when the layers within the lattice move relative to one another. This makes the alloy harder and less malleable.
Interstitial Alloys • A small proportion of an element with significantly smaller atoms is added to a metal. • Carbon can be added to iron to increase its hardness. The resulting product is called steel. • In steel, the smaller carbon atoms fit randomly in the hollows between the packed metal ions. • Like substitutional alloys the presence of different atoms makes it difficult for layers of the lattice to slide past one another. Making it less malleable than pure iron.
Your Turn • Page 89 • Question 7
Work Hardening and Heat Treatment • The way a metal has been prepared will also affect how it behaves. • Many metals are prepared in a liquid state and then cooled. • The rate at which a metal is cooled has significant effect on the properties of the solid.
Metal Crystals • A crystal is a region in a solid in which the particles are arranged in a regular way. • A sample of solid metal, however consists of a large number of small crystals. • At the point where one crystal meets another, the regular lattice is disrupted.
Galvanised Iron • Figure 5.17(b)This piece of steel (an alloy of iron, chromium and molybdenum) has individual crystals that formed as the molten metal mixture cooled down.
So who cares about crystals • The way a metal behaves, its malleability and its brittleness, will depend to some extent on the size of these crystals and the way they are arranged. • The smaller the crystals the harder the metal, as there will be less free movement of layers of ions over each other. • Smaller crystals, however, also mean more areas of disruption between them and this usually means the metal will be more brittle.
Solid Molten • Figure 5.18 In these representations of a metal, each circle represents an ion. (a) A solid crystalline metal; (b) A molten metal with a random arrangement of ions.
Work Hardening • Working (hammering) cold metals causes a rearrangement of crystal grains and a hardening of the metal. • Try bending a paper clip. Bend it once and it remains fairly pliable, keep bending it and it snaps. • Bending or working metal causes the crystal grains to be rearranged making the metal harder but more brittle.
Your Turn • Question 8 on Page 89. • The question regarding metals from your resource manual.
